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Chemistry: Comparative Table for Grade 11 & 12

 

Grades 11-12 Theory SLOs                                                                                                                                    2

Grades 11-12 Experimentation SLOs                                                                                                                 286

 

 

Grades 11-12 Theory SLOs

 

 

 

2006 National Curiculum

CAIE A Levels Curriculum 2025-2027

IB DP Curriculum 2016

NCC 2023

Guidance on SLOs

(Eleboration on the extent and depth of study and assessment expectations)

Essential Ideas

Rationale

Interdisciplinary Connection

How are Broad Topics Conceptualised

1 Stoichiometry

2 Atomic Structure

3 Theories of Covalent Bonding and Shapes of molecules

4 States of Matter I: Gases

5 States of Matter II: Liquids

6 States of Matter Ill: Solids

7 Chemical Equilibrium

8 Acids, Bases and Salts

9 Chemical Kinetics

11 Thermochemistry

12 Electrochemistry

13 s and p Block Elements

14 d and f Block Elements: Transition Elements

15 Organic Compounds

16 Hydrocarbons

17 Alkyl Halides and Amines

18 Alcohols, Phenols and Ethers

19 Carbonyl Compounds 1: Aldehydes and Ketones

20 Carbonyl Compounds 2: Carboxylic Acids and Functional Derivatives

21 Biochemistry

22 Industrial Chemistry

23 Environmental Chemistry

24 Analytical Chemistry

Chapter 1 Stoichiometry

Introduction

1.1 Mole and Avogadro's Number

1.2 Mole Calculations

1.3 Percentage Composition

1.4 Excess and Limiting Reagents

1.5 Theoretical Yield and Actual Yield as percentage

Chapter 2 Atomic Structure

Introduction

2.1 Discharge Tube Experiments

2.2 Application of Bohr's Model

2.2.1 Derivation of Radius, Energy, Frequency, Wave Length, Wave Number

2.2.2 Spectrum of Hydrogen Atom

2.2.3 Defects of Bohr's Theory

2.3 Planck's Quantum Theory_

2.3.1 Postulates With Derivation of E =hcv

2.4 X-Rays

2.4.1 Production, Properties and Uses

2.4.2 Types

2.4.3 X-rays and Atomic Number

2.4.4 Moseley's Experiment

2.4.5 Moseley's Law

2.5 Quantum Numbers and Orbitals

2.5.1 Principle Quantum Number

2.5.2 Azimuthal Quantum Number

2.5.3 Magnetic Quantum Number

2.5.4 Spin Quantum Number

2.5.5 Shapes of s, p and d Orbitals

2.6 Electronic Configuration

2.6.1 Aufbau Principle

2.6.2 Pauli's Exclusion Principle

2.6.3 Hund's Rule

2.6.4 Electronic Configurations

Chapter 3 Theories of Covalent Bonding and Shapes of molecules

Introduction

3.1 Shapes of molecules

3.1.1 VSEPR

3.1.2 Resonance

3.2 Theories of covalent bonding

3.2.1 VBT and hybridization

3.2.2. MOT

3.3 Bond Characteristics

3.3.1 Bond Energy

3.3.2 Bond Length

3.3.3 Ionic Character

3.3.4 Dipole Moment

3.4 Effect of Bonding on Physical and Chemical Properties

3.4.1 Solubility of Ionic and Covalent Compounds

3.4.2 Reactions of Ionic and Covalent Compounds

3.4.3 Directional and Non Directional Nature of Ionic and Covalent Bonds

Chapter 4 States of Matter I: Gases

Introduction

4.1 Kinetic Molecular Theory of Gases

4.1.1 Postulates of Kinetic Molecular Theory

4.1.2 Pressure and Its Units

4.2 Absolute Temperature Scale on the Basis of Charles Law

4.2.1 Brief recall of Boyle's and Charles' Law

4.2.2 Graphical Explanation of Absolute Zero

4.3 Avogadro's Law

4.4 Ideal Gas Equation

4.4.1 Derivation

4.4.2 Gas Constant and its Units

4.5 Deviation From Ideal Gas Behavior

4.5.1 Graphical Explanation

4.5.2 Causes for Deviation

4.6 Van der Waals Equation

4.6.1 Volume Correction

4.6.2 Pressure Correction

4.7 Dalton's Law of Partial Pressure

4.8 Graham's Law of Diffusion and Effusion

4.9 Liquefaction of Gases

4.9.1 Joule-Thomson Effect

4.9.2 Linde's Method of Liquefaction of Gases

4.10 Fourth State of Matter: Plasma

Chapter 5 States of Matter II: Liquids

Introduction

5.1 Kinetic Molecular Interpretation of Liquids

5.1.1 Simple properties of Liquids Describing Diffusion, Compression,

Expansion, Motion of Molecules, Kinetic Energy

5.2 Intermolecular Forces (Vander Waals Forces)

5.2.1 Dipole-Dipole interaction

5.2.2 Hydrogen Bonding

5.2.3 London Forces

5.3 Energetics of Phase Changes

5.3.1 Molar Heat of Fusion, Molar Heat of Vaporization, Molar Heat of Sublimation

5.3.2 Energy Changes and Intermolecular Attractions

5.3.3 Change of State and Dynamic Equilibrium

5.4 Liquid Crystals

5.4.1 Brief Description

5.4.2 Uses from Daily Life

Chapter 6 States of Matter Ill: Solids

Introduction

6.1 Kinetic Molecular Interpretation of Solids

6.1 .1 Simple Properties of Solids Describing Vibration of Molecules, Intermolecular Forces, Kinetic Energy

6.2 Types of Solids

6.2.1 Amorphous

6.2.2 Crystalline

6.3 Properties of Crystalline Solids

6.3.1 Symmetry

6.3.2 Geometrical Shape

6.3.3 Melting Point

6.3.4 Cleavage Plane

6.3.5 Habit of Crystal

6.3.6 Crystal Growth

6.3. 7 Anisotropy

6.3.8 Isomorphism

6.3.9 Polymorphism

6.3.10 Allotropy

6.3.11 Transition Temperature

6.4 Crystal Lattice

6.4.1 Unit Cell

6.4.2 NaCl Crystal

6.4.3 Lattice Energy

6.5 Types of Crystalline Solids

6.5.1 Ionic Solids

6.5.2 Covalent Solids

6.5.3 Metallic Solids

6.5.4 Molecular Solids

Chapter 7 Chemical Equilibrium

Introduction

7.1 Reversible Reactions and Dynamic Equilibrium

7 .1 .1 Concept and Explanation

7 .1.2 Law of Mass Action and Expression for Equilibrium Constant

7 .1.3 Relationship between Kc, Kp, Kx, Kn

7 .1.4 Importance of K and Reaction Quotient

7.2 Factors Affecting Equilibrium ( Le-Chatelier's Principle)

7.2.1 Effect of Change in Concentration

7.2.2 Effect of Change in Pressure or Volume

7.2.3 Effect of Change in Temperature

7.3 Industrial Application of Le-Chatelier's Principle (Haber's Process)

7.4 Solubility Product and Precipitation Reactions

7 .5 Common Ion Effect

Chapter 8 Acids, Bases and Salts

Introduction

8.1 Acidic, Basic and Amphoteric Substances

8.2 Bronsted-Lowery Definitions of Acids and Bases

8.2.1 Proton Donors and Acceptors

8.2.2 Relative Strength of Acids and Bases

8.3 Conjugate Acid-Base Pairs

8.4 Expressing the Strength of Acids and Bases

8.4.1 Ionization Equation of Water

8.4.2 pH, pOH and pKw

8.4.3 Acid Ionization Constant, Ka and pKa

8.4.4 Leveling Effect

8.4.5 Base Ionization Constant, Kb and pKb

8.4.6 Relationship of Ka and Kb

8.5 Lewis Definitions of Acids and Bases

8.6 Buffer Solutions and their Applications

8.7 Salt Hydrolysis

Chapter 9 Chemical Kinetics

Introduction

9.1 Chemical Kinetics

9.2 Rates of Reactions

9.2.1 Rate law or Rate Expression

9.2.2 Elementary and overall Rate Constant and Units

9.2.3 Order of Reaction and its Determination

9.2.4 Factors Affecting Rate of Reaction

9.3 Collision Theory, Transition State and Activation Energy

9.4 Catalysis

9.4.1 Characteristics of Catalysts

9.4.2 Homogeneous Catalysis

9.4.3 Heterogeneous Catalysis

9.4.4 Enzyme Catalysis

Chapter 10 Solutions and Colloids

Introduction

10.1 General Properties of Solutions

10.1.1 Solution, Suspension and Colloids

10.1.2 Hydrophilic and Hydrophobic Molecules

10.1.3 The Nature of Solutions in Liquid Phase

10.1.4 The Effect of Temperature and Pressure on Solubility

10.2 Concentration Units

10.2.1 Percent

10.2.2 Molarity

10.2.3 Molality

10.2.4 Mole fraction

10.2.5 Parts per million, billion, and trillion

10.3 Raoult's Law

10.3.1 Non-Volatile Non- -Electrolyte Solutes in Volatile solvents

10.3.2 When both Components are Volatile

10.4 Colligative Properties of dilute Solutions

10.4.1 Vapour Pressure Lowering

10.4.2 Boiling Point Elevation and Freezing Point Depression

10.4.3 Molar Mass Determination by Vapor Pressure Lowering, Boiling Point Elevation and Freezing Point Depression

10.4.4 Osmotic Pressure and Reverse Osmosis

10.5 Colloids

10.5.1 Properties of Colloids

10.5.2 Types of Colloids

Chapter 11 Thermochemistry

Introduction

11.1 Energy in Chemical Reactions

11.2 Thermodynamics

11.3 Internal Energy

11.4 First Law of Thermodynamics

11.5 Standard State and Standard Enthalpy Changes

11.6 Heat Capacity

11.7 Calorimeter

11.8 Hess's Law: Enthalpy Change Calculations

11.9 Born Haber Cycle

Chapter 12 Electrochemistry

Introduction

12.1 Oxidation-Reduction Concepts

12.1.1 Oxidation and Reduction

12.1.2 Oxidation Numbers

12.1.3 Recognizing Oxidation Reduction Reactions

12.1.4 Balancing Oxidation Reduction Equations by Oxidation Number Method

12.1.5 Balancing Oxidation Reduction Equations by the Half Reaction Method

12.1.6 Chemistry of Some Important Oxidizing and Reducing Agents

12.2 Electrode, Electrode Potential and Electrochemical Series

12.3 Types of Electrochemical Cells

12.3.1 Electrolytic Cells

12.3.2 Electrolysis of Aqueous NaCl

12.3.3 Voltaic Cells

12.3.3.1 Standard State Cell Potential for Voltaic Cell

12.3.3.2 Standard State Reduction Half Cell Potential

12.3.3.3 Standard State Cell Potentials and Spontaneous Reaction

12.3.4 Batteries

12.3.4.1 Primary Batteries

12.3.4.2 Secondary Batteries

12.3.4.3 Fuel Cells

12.3.5 Corrosion and its Prevention

Chapter 13 s and p Block Elements

Introduction

13.1 Period 3 (Na to Ar)

13.1.1 Physical and Atomic Properties of the Elements

13.1.1.1 Electronic Structure

13.1.1.2 Trends in Atomic Radius

13.1.1.3 Trends in First Ionization Energy

13.1.1.4 Trends in Electronegativity

13.1.1.5 Trends in Electrical Conductivity

13.1.1.6 Trends in Melting and Boiling Points

13.1.2 Reactions of the Period 3 Elements with Water, Oxygen and Chlorine

13.1.3 Physical Properties of the Oxides

13.1.3.1 Structure

13.1.3.2 Melting and Boiling Points

13.1.3.3 Electrical Conductivity

13.1.4 Acid-Base Behavior of the Oxides

13.1.4.1 Trends in Acid Base Behavior

13.1.4.2 Reactions of Oxides with Water, Acids and Bases

13.1.5 Chlorides of the Period 3 Elements

13.1.5.1 Structure

13.1.5.2 Melting and Boiling Points

13.1.5.3 Electrical Conductivity

13.1.5.4 Solubility in Water

13.1.6 Hydroxides of the Period 3 Elements

13.1.6.1 Sodium and Magnesium Hydroxides

13.1.6.2 Aluminum Hydroxide

13.1.6.3 Other Hydroxides

13.2 Group 1-Elements

13.2.1 Atomic and Physical Properties

13.2.1.1 Trends in Atomic Radius

13.2.1.2 Trends in First Ionization Energy

13.2.1.3 Trends in Electronegativity

13.2.1.4 Trends in Melting and Boiling Points

13.2.1.5 Trends in Density

13.2.2 Trends in Reactivity with Water

13.2.3 Reactions with Oxygen

13.2.3.1 Reactions with Air or Oxygen and the formation of Normal Oxides,Peroxides, speroxides and their Stability

13.2.3.2 Reactions of Oxides with Water and Dilute Acids

13.2.4 Reactions with Chlorine

13.2.5 Effect of Heat on Nitrates, Carbonates and Hydrogen-Carbonates explaining the Trend in Terms of the Polarizing Ability of the Positive Ion

13.2.6 Flame Tests: Origin of Flame Colors

13.3 Group 2- Elements

13.3.1 Atomic and Physical Properties

13.3.1.1 Trends in Atomic Radius

13.3.1.2 Trends in First Ionization Energy

13.3.1.3 Trends in Electronegativity

13.3.1.4 Trends in Melting and Boiling Points

13.3.2 Trends in Reactivity with Water

13.3.3 Reactions with Oxygen and Nitrogen

13.3.3.1 Simple Oxides

13.3.3.2 Formation of Peroxides on Heating with Oxygen

13.3.3.3 Formation of Nitrides on Heating in Air

13.3.4 Trends in Solubility of the Hydroxides, Sulphates and Carbonates

13.3.5 Trends in Thermal Stability of the Nitrates and Carbonates

13.3.6 How Beryllium Differs from other Members of its Group?

13.3.6.1 Why is Beryllium Chloride Covalent and not Ionic?

13.3.6.2 Amphoteric Beryllium Hydroxide

13.4 Group 4 -Elements

13.4.1 Physical Properties : Melting and Boiling Points

13.4.2 The Trend from Non-Metal to Metal

13.4.3 Oxidation State

13.4.4 Possible Oxidation States

13.4.4.1 Inert Pair Effect in Formation of Ionic Bonds

13.4.4.2 Inert Pair Effect and the Formation of Covalent Bonds

13.4.5 Chlorides of Carbon, Silicon and Lead

13.4.5.1 Structures and Stability

13.4.5.2 Reactions with Water

13.4.6 Oxides

13.4.6.1 Structure of Carbon Dioxide and Silicon Dioxide

13.4.6.2 Acid Base Behavior of Group IV Oxides

13.5 Group 7-Elements:Halogens

13.5.1 Atomic and Physical Properties

13.5.1.1 Trends in Atomic Radius

13.5.1.2 Trends in Electronegativity

13.5.1.3 Trends in Electron Affinity

13.5.1.4 Trends in Melting and Boiling Points

13.5.1.5 Bond Enthalpies

13.5.1.5.1 Bond Enthalpies in Halogens

13.5.1.5.2 Bond Enthalpies in Hydrogen Halides

13.5.2 Strength of Halogens as Oxidizing Agents : F>Cl>Br>I

13.5.3 The Acidity of Hydrogen Halides

13.5.4 Halide Ions as Reducing Agents and Trends in Reducing Strength Ability of Halide Ions

Chapter 14 d and f Block Elements: Transition Elements

Introduction

14.1 General Features

14.1.1 General Features of Transition Elements

14.1.2 Electronic Structure

14.1.3 Binding Energy

14.1.4 Variable Oxidation States

14.1.5 Catalytic Activity

14.1.6 Magnetic Behaviour

14.1. 7 Alloy formation

14.2 Coordination Compounds

14.2.1 Complex Ion

14.2.2 Nomenclature of Coordination compounds

14.2.3 Shapes of Complex Ions with Coordination number 2, 4 and 6

14.2.4 Colour of Complexes

14.3 Chemistry of Some important Transition Elements

14.3.1 Vanadium

14.3.1.1 Oxidation States

14.3.1.2 As Catalyst in Contact Process

14.3.2 Chromium

14.3.2.1 Oxidation States

14.3.2.2 Chromate - Dichromate Equilibrium

14.3.2.3 Reduction of Chromate VI Ions with Zn and an Acid

14.3.2.4 Potassium Dichromate as an Oxidizing Agent in Organic Chemistry

14.3.2.5 Potassium Dichromate as an Oxidizing Agent in Titrations

14.3.3 Manganese

14.3.3.1 Oxidation States

14.3.3.2 Potassium Manganate VII as an Oxidizing Agent in Organic Chemistry

14.3.3.3 Potassium Manganate VII as an Oxidizing Agent in Titrations

14.3.4 Iron

14.3.4.1 Oxidation States

14.3.4.2 Iron as Catalyst in Haber's Process

14.3.4.3 Iron as Catalyst in Reaction between Persulphate and Iodide Ions

14.3.4.4 Reactions of Hexaaquairon(II) and Hexaaquairon(llI) with Water and Ammonia

14.3.4.5 Reactions of Iron (II) and (Ill) Ions with Carbonate, and Thiocyanate Ions

14.3.5 Copper

14.3.5.1 Oxidation States

14.3.5.2 The Reaction of Hexaaquacopper(II) Ions with Hydroxide Ions, Ammonia, and Carbonate Ions

Chapter 15 Organic Compounds

Introduction

15.1 Sources

15.1.1 Fossil Remains: Coal, Petroleum, Natural Gas

15.1.2 Plants and Natural Products Chemistry

15.1.3 Partial and Total Synthesis

15.1.4 Products of Biotechnology

15.2 Coal as a Source of Organic Compounds

15.2.1 Destructive Distillation of Coal

15.2.2 Conversion of Coal to Petroleum

15.3 Characteristics of Organic Compounds

15.4 Uses of Organic Compounds

15.5 New Allotrope of Carbon: Bucky Ball

15.6 Functional Groups and Homologous Series

15.7 Detection of Elements in Organic Compounds

Chapter 16 Hydrocarbons

Introduction

16.1 Types of Hydrocarbons

16.2 Alkanes and Cycloalkanes

16.2.1 Nomenclature

16.2.2 Physical Properties

16.2.3 Structure

16.2.4 Relative Stability

16.2.5 Reactivity

16.3 Radical Substitution Reactions

16.3.1 Overview

16.3.2 Reaction Mechanism

16.4 Oxidation of Organic compounds

16.5 Alkenes

16.5.1 Nomenclature

16.5.2 Relative Stability

16.5.3 Structure

16.5.4 Preparation of Alkenes

16.5.4.1 Dehydration of Alcohols

16.5.4.2 Dehydrohalogenation of Alkyl Halides

16.5.5 Reactivity

16.5.6 Reactions

16.5.6.1 Hydrogenation

16.5.6.2 Hydrohalogenation

16.5.6.3 Hydration

16.5.6.4 Halogenation

16.5.6.5 Halohydration

16.5.6.6 Epoxidation

16.5.6.7 Ozonolysis

16.5.6.8 Polymerization

16.5. 7 Conjugation

16.6 Isomerism

16.6.1 Chiral Centre

16.6.2 Carbon-Based Chiral Centers

16.6.3 Optical Activity

16.6.4 Optical Isomers

16.6.5 Stereoisomerism

16.6.6 Structural Isomerism

16.7 Alkynes

16. 7 .1 Nomenclature

16.7.2 Relative Stability

16.7.3 Structure

16.7.4 Physical Properties

16.7.5 Preparation of Alkynes by Elimination Reactions

16.7.6 Reactivity

16.7.7 Acidity of Terminal Alkynes

16.7.8 Addition Reactions of Alkynes

16. 7 .8.1 Hydrogenation

16.7.8.2 Dissolving Metal reduction

16. 7 .8.3 Hydrohalogenation

16. 7 .8.4 Hydration

16. 7 .8.5 Bromination

16.7.8.6 Ozonolysis

16.8 Benzene and Substituted Benzenes

16.8.1 Nomenclature

16.8.2 Physical Properties

16.8.3 Structure Molecular Orbital Aspects

16.8.4 Resonance, Resonance Energy and Stabilization

16.8.5 Reactivity And Reactions

16.8.5.1 Addition Reaction

16.8.5.2 Electrophilic Aromatic Substitution Reactions

16.8.5.2.1 General Introduction

16.8.5.2.2 Nitration

16.8.5.2.3 Sulfonation

16.8.5.2.4 Halogenation

16.8.5.2.5 Friedel-Crafts Alkylation

16.8.5.2.6 Friedel-Crafts Acylation

16.8.5.2.7 Substituent Effects - (Table of Substituent Effects)

16.8.5.2.8 Making Polysubstituted Benzenes

Chapter 17 Alkyl Halides and Amines

Introduction

17.1 Alkyl halides

17 .1 .1 Nomenclature

17 .1.2 Physical Properties

17 .1.3 Structure

17 .1.4 Preparations of Alkyl Halides

17.1.4.1 Reaction of Alcohols with Hydrogen Halides

17.1.4.2 Reaction Of Alcohols with other Halogenating Agents (SOCb, PX3)

17.1.4.3 Radical Halogenation of Alkanes

17 .1.5 Reactivity

17.1.6 Nucleophilic Substitution Reactions

17.1.6.1 General Introduction

17.1.6.2 Important Concepts

17.1.6.2.1 Carbocations and Their Stability

17.1.6.2.2 Nucleophile and Base

17.1.6.2.3 Substrate and Leaving Group

17.1.6.3 SN1 Mechanism

17.1.6.4 SN2 Mechanism

17 .1. 7 1, 2 Elimination Reactions

17.1.7.1 Overview

17.1.7.2 E1 Mechanism

17.1.7.3 E2 Mechanism

17.1.8 Substitution versus Elimination

17.2 Organometallic Compounds (Grignard's Reagents)

17.2.1 Preparation of Grignard's Reagents

17.2.2 Reactivity

17.2.3 Reactions of Grignard's Reagents

17.2.3.1 with Aldehydes and Ketones

17.2.3.2 with Esters

17.2.3.3 with CO2

17.3 Amines

17.3.1 Nomenclature

17.3.2 Physical Properties

17.3.3 Structure

17.3.4 Basicity

17.3.5 Preparation of Amines

17.3.5.1 Alkylation of Ammonia by Alkyl Halides

17.3.5.2 Reductions of Nitrogen Containing Functional Groups

17.3.5.2.1 Nitriles

17.3.5.2.2 Nitro

17.3.5.2.3 Amides

17.3.6 Reactivity

17.6.7 Reactions of Amines

17.6.7.1 Overview

17.6.7.2 Alkylation of Amines By Alkyl Halides

17.6.7.3 Reaction of Amines with Aldehydes and Ketones

17.6.7.4 Preparation of Amides

17.6.7.5 Preparation of Diazonium Salts

Chapter 18 Alcohols, Phenols and Ethers

Introduction

18.1 Alcohols

18.1.1 Nomenclature

18.1.2 Physical Properties

18.1.3 Structure

18.1.4 Acidity

18.1.5 Preparations of Alcohols

18.1.5.1 Hydration of Alkenes (review)

18.1.5.2 Hydrolysis of Alkyl Halides (review)

18.1.5.3 Reaction of RMgX With Aldehydes And Ketones

18.1.5.4 Reduction of Aldehydes and Ketones

18.1.5.5 Reaction of RMgX with Esters (review)

18.1.5.6 Reduction of Carboxylic Acids and Esters

18.1.6 Reactivity (review)

18.1. 7 Reactions of Alcohols

18.1. 7 .1 Reaction with HX to give Alkyl Halides (review)

18.1. 7 .2 Reaction with SOCl2, PX3 to give Alkyl Halides (review)

18.1. 7 .3 Acid Catalyzed Dehydration (review)

18.1. 7.4 Preparation of Esters

18.1. 7 .5 Oxidation

18.1. 7 .6 Cleavage of 1,2-diols

18.1.8 The Sulfur Analogues (Thiols, RSH ) Phenols

18.2.1 Nomenclature

18.2.2 Structure

18.2.3 Physical Properties

18.2.4 Acidity

18.2.5 Preparation of Phenols from

18.2.5.1 Benzene Sulfonic Acid

18.2.5.2 Chlorobenzene

18.2.5.3 Acidic Oxidation of Cumene

18.2.5.4 Hydrolysis of Diazonium Salts

18.2.6 Reactivity

18.2. 7 Reactions of Phenols

18.2. 7 .1 Electrophilic Aromatic Substitutions (review)

18.2. 7 .2 Reaction with Sodium Metal

18.2. 7 .3 Oxidation

18.2.8 Difference between Alcohol and Phenol Ethers

18.3.1 Nomenclature

18.3.2 Preparation

18.3.3 Physical properties

18.3.4 Chemical reactivity

Chapter 19 Carbonyl Compounds 1: Aldehydes and Ketones

Introduction

19.1 Nomenclature

19.2 Physical Properties

19.3 Structure

19.4 Preparations of Aldehydes and Ketones

19.4.1 Ozonolysis of Alkenes (review)

19.4.2 Hydration of Alkynes (review)

19.4.3 Oxidation of Alcohols (review)

19.4.4 Friedel-Crafts Acylation of Aromatics (review)

19.5 Reactivity

19.6 Reactions of Aldehydes and Ketones

19.6.1 Nucleophilic Addition Reactions (Acid and Base Catalyzed)

19.6.2 Relative Reactivity

19.6.3 Reduction of Aldehydes and Ketones

19.6.3.1 To Hydrocarbons

19.6.3.2 Using Hydrides to Give Alcohols

19.6.3.3 Using Carbon Nucleophiles

19.6.3.4 Using Nitrogen Nucleophiles

19.6.3.5 Using Oxygen Nucleophiles

19.6.4 Oxidation Reactions

Chapter 20 Carbonyl Compounds 2: Carboxylic Acids and Functional Derivatives

Introduction

20.1 Nomenclature

20.2 Physical Properties

20.3 Structure

20.4 Acidity

20.5 Preparations of Carboxylic Acids

20.5.1 Carbonation of Grignard's Reagent (review)

20.5.2 Hydrolysis of Nitriles

20.5.3 Oxidation of Primary Alcohols (review)

20.5.4 Oxidation of Aldehydes (review)

20.5.5 Oxidation of Alkyl benzenes (review)

20.6 Reactivity

20.7 Reactions of Carboxylic Acids

20.7.1 Conversion to Carboxylic Acid Derivatives

20.7.1.1 Acyl Halides

20.7.1.2 Acid Anhydrides

20.7.1.3 Esters

20.7.1.4 Amides

20.7.2 Summary of Reactions that lnterconvert Carboxylic Acids Derivatives

20.7.3 Reduction to Alcohols

20.7.4 Decarboxylation Reactions

20.7.5 Reactions of Carboxylic Acid Derivatives

20.7.5.1 Reactions of Acyl Halides, Friedel-Crafts Acylation (review)

20.7.5.2 Reactions of Acid Anhydrides, Hydrolysis

20.7.5.3 Reactions of Esters, Hydrolysis, Reduction, and with Grignard's Reagent

20.7.5.4 Reactions of Amides, Hydrolysis and Reduction

20.7.5.5 Reactions of Nitriles, Hydrolysis, Reduction, and reactions with Grignard's Reagent

Chapter 21 Biochemistry

Introduction

21.1 Carbohydrates

21.1.1 Classification

21.1.2 Functions

21.1.3 Nutritional Importance

21.2 Proteins

21.2.1 Classification

21.2.2 Structure

21.2.3 Properties

21.2.4 Importance of Proteins

21.3 Enzymes

21.3.1 Role of Enzyme as a Biocatalyst

21.3.2 Factors Affecting Enzyme activity

21.3.3 Industrial Application of Enzyme

21.4 Lipids

21.4.1 Classification

21.4.2 Structure

21.4.3 Properties of Lipids

21.4.4 Nutritional and Biological Importance of lipids

21.5 Nucleic Acids

21.5.1 Structural Components of DNA and RNA

21.5.2 Nucleic Acid Polymers

21.5.3 Storage of Genetic Information

21.6 Minerals of Biological Significance

21.6.1 Sources of Important Minerals

21.6.2 Biological Significance of Iron Calcium Phosphorous and Zinc

Chapter 22 Industrial Chemistry

Introduction

22.1 Introduction to the Chemical Process Industry and Raw Materials used

22.2 Safety Considerations in Process Industries

22.3 Dyes

22.4 Pesticides

22.5 Petrochemicals

22.6 Synthetic Polymers (PVC and Nylon)

22.7 Cosmetics: Lipsticks, Nail Varnish and Remover, hair Dyes

22.8 Adhesives

Chapter 23 Environmental Chemistry

Introduction

23.1 Chemistry of the Troposphere

23.1.1 Chemical Reactions in the Atmosphere

23.1.2 COx, NOx, VOCs, SOx , 0 3

23.1.3 Automobile, Pollutants and the Catalytic Converter

23.1.4 Industrial Smog

23.1.5 Photochemical Smog

23.1.6 Global Warming and Climate Change

23.1. 7 Acid Rain

23.2 Chemistry of the Stratosphere: Production and Destruction of Ozone

23.3 Water Pollution and Water Treatment

23.3.1 Types of Water Pollutants

23.3.1.1 Suspended Solids and Sediments

23.3.1.2 Dissolved Solids

23.3.1.3 Thermal Pollution

23.3.2 Waste water treatment

23.4 Green Chemistry

Chapter 24 Analytical Chemistry

Introduction

24.1 Classical Method of Analysis:

Combustion Analysis and determination of Molecular Formula

24.2 Modern Methods of Analysis

24.2.1 Spectroscopy

24.2.2 Spectroscopic Methods

24.2.2.1 Infra Red (IR)

24.2.2.2 Ultra-Violet/ Visible (UV-VIS)

24.2.2.3 Nuclear Magnetic Resonance (NMR)

24.2.2.4 Atomic Emission and Absorption

24.2.2.5 Mass Spectrometry MS)

AS

Physical chemistry

1 Atomic structure

2 Atoms, molecules and stoichiometry

3 Chemical bonding

4 States of matter

5 Chemical energetics

6 Electrochemistry

7 Equilibria

8 Reaction kinetics

Inorganic chemistry

9 The Periodic Table: chemical periodicity

10 Group 2

11 Group 17

12 Nitrogen and sulfur

Organic chemistry

13 An introduction to AS Level organic chemistry

14 Hydrocarbons

15 Halogen compounds

16 Hydroxy compounds

17 Carbonyl compounds

18 Carboxylic acids and derivatives

19 Nitrogen compounds

20 Polymerisation

21 Organic synthesis Analysis

22 Analytical techniques

A2

Physical chemistry

23 Chemical energetics

24 Electrochemistry

25 Equilibria

26 Reaction kinetics

Inorganic chemistry

27 Group 2

28 Chemistry of transition elements

Organic chemistry

29 An introduction to A Level organic chemistry

30 Hydrocarbons

31 Halogen compounds

32 Hydroxy compounds

33 Carboxylic acids and derivatives

34 Nitrogen compounds

35 Polymerisation

36 Organic synthesis

Analysis

37 Analytical techniques

AS Level subject content Physical chemistry

1 Atomic structure

1.1 Particles in the atom and atomic radius

1.2 Isotopes

1.3 Electrons, energy levels and atomic orbitals

1.4 Ionisation energy

2 Atoms, molecules and stoichiometry

2.1 Relative masses of atoms and molecules

2.2 The mole and the Avogadro constant

2.3 Formulae

2.4 Reacting masses and volumes (of solutions and gases)

3 Chemical bonding

3.1 Electronegativity and bonding

3.2 Ionic bonding

3.3 Metallic bonding

3.4 Covalent bonding and coordinate (dative covalent) bonding

3.5 Shapes of molecules

3.6 Intermolecular forces, electronegativity and bond properties

3.7 Dot-and-cross diagrams

4 States of matter

4.1 The gaseous state: ideal and real gases and pV = nRT

5 Chemical energetics

5.1 Enthalpy change, ΔH

5.2 Hess’s Law

6 Electrochemistry

6.1 Redox processes: electron transfer and changes in oxidation number (oxidation state)

7 Equilibria

7.1 Chemical equilibria: reversible reactions, dynamic equilibrium

7.2 Brønsted–Lowry theory of acids and bases

8 Reaction kinetics

8.1 Rate of reaction

8.2 Effect of temperature on reaction rates and the concept of activation energy

8.3 Homogeneous and heterogeneous catalysts

Inorganic chemistry

9 The Periodic Table: chemical periodicity

9.1 Periodicity of physical properties of the elements in Period 3

9.2 Periodicity of chemical properties of the elements in Period 3

9.3 Chemical periodicity of other elements

10 Group 2

10.1 Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their compounds

11 Group 17

11.1 Physical properties of the Group 17 elements

11.2 The chemical properties of the halogen elements and the hydrogen halides

11.3 Some reactions of the halide ions

12 Nitrogen and sulfur

12.1 Nitrogen and sulfur

Organic chemistry

13 An introduction to AS Level organic chemistry

13.1 Formulae, functional groups and the naming of organic compounds

13.2 Characteristic organic reactions

13.3 Shapes of organic molecules; σ and π bonds

13.4 Isomerism: structural and stereoisomerism

14 Hydrocarbons

14.1 Alkanes

14.2 Alkenes

15 Halogen compounds

15.1 Halogenoalkanes

15.1 Halogenoalkanes (continued)

16 Hydroxy compounds

16.1 Alcohols

17 Carbonyl compounds

17.1 Aldehydes and ketones

18 Carboxylic acids and derivatives

18.1 Carboxylic acids

18.2 Esters

19 Nitrogen compounds

19.1 Primary amines

19.2 Nitriles and hydroxynitriles

20 Polymerisation

20.1 Addition polymerisation

21 Organic synthesis

21.1 Organic synthesis

Analysis

22 Analytical techniques

22.1 Infrared spectroscopy

22.2 Mass spectrometry

A Level subject content

Physical chemistry

23 Chemical energetics

23.1 Lattice energy and Born-Haber cycles

23.2 Enthalpies of solution and hydration

23.3 Entropy change, ΔS

23.4 Gibbs free energy change, ΔG

24 Electrochemistry

24.1 Electrolysis

24.2 Standard electrode potentials E

; standard cell potentials E

and the Nernst equation cell

25 Equilibria

25.1 Acids and bases

25.2 Partition coefficients

26 Reaction kinetics

26.1 Simple rate equations, orders of reaction and rate constants

26.2 Homogeneous and heterogeneous catalysts

Inorganic chemistry

27 Group 2

27.1 Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their compounds

28 Chemistry of transition elements

28.1 General physical and chemical properties of the first row of transition elements, titanium to copper

28.2 General characteristic chemical properties of the first set of transition elements, titanium to copper

28.3 Colour of complexes

28.4 Stereoisomerism in transition element complexes

28.5 Stability constants, Kstab

Organic chemistry

29 An introduction to A Level organic chemistry

29.1 Formulae, functional groups and the naming of organic compounds

29.2 Characteristic organic reactions

29.3 Shapes of aromatic organic molecules; σ and π bonds

29.4 Isomerism: optical

30 Hydrocarbons

30.1 Arenes

31 Halogen compounds

31.1 Halogen compounds

32 Hydroxy compounds

32.1 Alcohols

32.2 Phenol

33 Carboxylic acids and derivatives

33.1 Carboxylic acids

33.2 Esters

33.3 Acyl chlorides

34 Nitrogen compounds

34.1 Primary and secondary amines

34.2 Phenylamine and azo compounds

34.3 Amides

34.4 Amino acids

35 Polymerisation

35.1 Condensation polymerisation

35.2 Predicting the type of polymerisation

35.3 Degradable polymers

36 Organic synthesis

36.1 Organic synthesis

37 Analytical techniques

37.1 Thin-layer chromatography

37.2 Gas / liquid chromatography

37.3 Carbon-13 NMR spectroscopy

37.4 Proton (1H) NMR spectroscopy

Core

1. Stoichiometric relationships

2. Atomic structure

3. Periodicity

4. Chemical bonding and structure

5. Energetics/thermochemistry

6. Chemical kinetics

7. Equilibrium

8. Acids and bases

9. Redox processes

10. Organic chemistry

11. Measurement and data processing

Additional higher level (AHL)

12. Atomic structure

13. The periodic table—the transition metals

14. Chemical bonding and structure

15. Energetics/thermochemistry

16. Chemical kinetics

17. Equilibrium

18. Acids and bases

19. Redox processes

20. Organic chemistry

21. Measurement and analysis

Option

A. Materials

B. Biochemistry

C. Energy

D. Medicinal chemistry

Core

Topic 1: Stoichiometric relationships 13.5

1.1 Introduction to the particulate nature of matter and chemical change

1.2 The mole concept

1.3 Reacting masses and volumes

Topic 2: Atomic structure 6

2.1 The nuclear atom

2.2 Electron configuration

Topic 3: Periodicity 6

3.1 Periodic table

3.2 Periodic trends

Topic 4: Chemical bonding and structure 13.5

4.1 Ionic bonding and structure

4.2 Covalent bonding

4.3 Covalent structures

4.4 Intermolecular forces

4.5 Metallic bonding

Topic 5: Energetics/thermochemistry 9

5.1 Measuring energy changes

5.2 Hess’s Law

5.3 Bond enthalpies

Topic 6: Chemical kinetics 7

6.1 Collision theory and rates of reaction

Topic 7: Equilibrium 4.5

7.1 Equilibrium

Topic 8: Acids and bases 6.5

8.1 Theories of acids and bases

8.2 Properties of acids and bases

8.3 The pH scale

8.4 Strong and weak acids and bases

8.5 Acid deposition

Topic 9: Redox processes 8

9.1 Oxidation and reduction

9.2 Electrochemical cells

Topic 10: Organic chemistry 11

10.1 Fundamentals of organic chemistry

10.2 Functional group chemistry

Topic 11: Measurement and data processing 10

11.1 Uncertainties and errors in measurement and results

11.2 Graphical techniques

11.3 Spectroscopic identification of organic compounds

Additional higher level (AHL)

Topic 12: Atomic structure 2

12.1 Electrons in atoms

Topic 13: The periodic table—the transition metals 4

13.1 First-row d-block elements

13.2 Coloured complexes

Topic 14: Chemical bonding and structure 7

14.1 Covalent bonding and electron domain and molecular geometries

14.2 Hybridization

Topic 15: Energetics/thermochemistry 7

15.1 Energy cycles

15.2 Entropy and spontaneity

Topic 16: Chemical kinetics 6

16.1 Rate expression and reaction mechanism

16.2 Activation energy

Topic 17: Equilibrium 4

17.1 The equilibrium law

Topic 18: Acids and bases 10

18.1 Lewis acids and bases

18.2 Calculations involving acids and bases

18.3 pH curves

Topic 19: Redox processes 6

19.1 Electrochemical cells

Topic 20: Organic chemistry 12

20.1 Types of organic reactions

20.2 Synthetic routes

20.3 Stereoisomerism

Topic 21: Measurement and analysis 2

21.1 Spectroscopic identification of organic compounds

Options

A: Materials

Core topics

A.1 Materials science introduction

A.2 Metals and inductively coupled plasma (ICP) spectroscopy

A.3 Catalysts

A.4 Liquid crystals

A.5 Polymers

A.6 Nanotechnology

A.7 Environmental impact—plastics

Additional higher level topics

A.8 Superconducting metals and X-ray crystallography (HL only)

A.9 Condensation polymers (HL only)

A.10 Environmental impact—heavy metals (HL only)

B: Biochemistry

Core topics

B.1 Introduction to biochemistry

B.2 Proteins and enzymes

B.3 Lipids

B.4 Carbohydrates

B.5 Vitamins

B.6 Biochemistry and the environment

Additional higher level topics

B.7 Proteins and enzymes (HL only)

B.8 Nucleic acids (HL only)

B.9 Biological pigments (HL only)

B.10 Stereochemistry in biomolecules (HL only)

C: Energy

Core topics

C.1 Energy sources

C.2 Fossil fuels

C.3 Nuclear fusion and fission

C.4 Solar energy

C.5 Environmental impact—global warming

Additional higher level topics

C.6 Electrochemistry, rechargeable batteries and fuel cells (HL only)

C.7 Nuclear fusion and nuclear fission (HL only)

C.8 Photovoltaic and dye-sensitized solar cells (HL only)

D: Medicinal chemistry

Core topics

D.1 Pharmaceutical products and drug action

D.2 Aspirin and penicillin

D.3 Opiates

D.4 pH regulation of the stomach

D.5 Anti-viral medications

D.6 Environmental impact of some medications

Additional higher level topics

D.7 Taxol—a chiral auxiliary case study (HL only)

D.8 Nuclear medicine (HL only)

D.9 Drug detection and analysis (HL only)

1 Chemical Foundation

Introduction to Chemistry

2 Nature of Science in Chemistry

History of Chemistry

TOK and Nature of Chemistry

Scientific Method

3 Physical Chemistry

Stoichiometry

Atomic Structure

Chemical Bonding

States and Phases of Matter

Energetics & Thermochemistry

Reaction Kinetics

Equilibria (Basic, Acid-Base, Ionic)

Electrochemistry and Redox

4 Inorganic Chemistry

Periodicity

Group 2

Group 17

Transition Metals

5 Envioronmental Chemistry

Air

Water

Green chemistry and Sustainability

6 Organic Chemistry

Introduction to Organic Chemistry (Nomenclature, Functional Group, Isomerism, Formulae)

Hydrocarbons (Alkanes, Alkenes, Alkynes, Benzene)

Halogenalkanes

Hydroxy Compounds (alcohols and phenols)

Carbonyl Compounds (Carboxylic Acids, Aldehydes, Ketones, Esters)

Nitrogen Compounds

Sulfur Compounds

Polymerization

Organic Synthesis

Biochemistry

7 Lab and Analysis Skills

Combustion Analysis

Mass spectrometry

Spectrocopy

NMR

Chromatography

8 Chemistry in Context

Materials

Energy

Medicine

Agriculture

Industry

 

 

 

 

Introduction to Chemistry
(Chemical Foundation)

 

 

 

1.1 Introduction to the particulate nature of matter and chemical change

Understandings:

• Atoms of different elements combine in fixed ratios to form compounds, which

have different properties from their component elements.

• Mixtures contain more than one element and/or compound that are not

chemically bonded together and so retain their individual properties.

• Mixtures are either homogeneous or heterogeneous.

Applications and skills:

• Deduction of chemical equations when reactants and products are specified.

• Application of the state symbols (s), (l), (g) and (aq) in equations.

• Explanation of observable changes in physical properties and temperature

during changes of state.

Guidance:

• Balancing of equations should include a variety of types of reactions.

• Names of the changes of state—melting, freezing, vaporization (evaporation

and boiling), condensation, sublimation and deposition—should be covered.

 

Understand the scientific method and how it applies to chemistry experimentation

Be able to use mathematical concepts such as units, conversion factors, and basic algebra to solve chemistry problems

Understand basic atomic and molecular structure, including the organization of the periodic table

Be able to write and balance chemical equations

Understand the concepts of chemical bonding, including ionic, covalent, and metallic bonding

Understand the properties and behavior of gases, liquids, and solids, including their intermolecular forces

Understand the concepts of acids and bases and be able to use pH calculations

Understand the concept of chemical equilibrium and its applications

Understand basic thermochemistry concepts, including enthalpy and entropy

Understand the basics of chemical kinetics and thermodynamics

Understand the basics of analytical chemistry, including titrations and spectroscopy

Understand the basics of organic chemistry, including functional groups and reactions.

In current curricula taught in Pakistan, students usually directly jump into studying one branch of chemistry after the other, without being provided a birds-eye overview of the field as a whole. These introductory SLOs provide this context.

These SLOs also help provide students with an understanding of what are the big questions that chemistry today is aiming to address, and how these endeavors are often interdisciplinary.

This also meets a key recommendation made by stakeholders in the InterProvincial Curriculum and Standards Workshops.

These introductory SLOs are not intended to be comprehensive, and are complemented by the Cross-Cutting Theme SLOs that are embedded into each of the subsequent Units. This will help students also understand what kind of career trajectories are possible by studying further chemistry, beyond simply becoming a scientist, engineer or a professor.

 

Measurement (Chemical Foundation)

 

 

Uncertainties and errors in measurement and results:

Qualitative data includes all non-numerical information obtained from observations not from measurement.

- Quantitative data are obtained from measurements, and are always associated with random errors/uncertainties, determined by the apparatus, and by human limitations such as reaction times.

- Propagation of random errors in data processing shows the impact of the uncertainties on the final result.

- Experimental design and procedure usually lead to systematic errors in measurement, which cause a deviation in a particular direction.

- Repeat trials and measurements will reduce random errors but not systematic errors

Graphical techniques:

Graphical techniques are an effective means of communicating the effect of an independent variable on a dependent variable, and can lead to determination of physical quantities.

- Sketched graphs have labelled but unscaled axes, and are used to show qualitative trends, such as variables that are proportional or inversely proportional.

- Drawn graphs have labelled and scaled axes, and are used in quantitative measurements

Spectroscopic identification of organic compounds:

The degree of unsaturation or index of hydrogen deficiency (IHD) can be used to determine from a molecular formula the number of rings or multiple bonds in a molecule.

- Mass spectrometry (MS), proton nuclear magnetic resonance spectroscopy (1H NMR) and infrared spectroscopy (IR) are techniques that can be used to help identify compounds and to determine their structure

Uncertainties and errors in measurement and results:

Qualitative data includes all non-numerical information obtained from observations not from measurement.

- Quantitative data are obtained from measurements, and are always associated with random errors/uncertainties, determined by the apparatus, and by human limitations such as reaction times.

- Propagation of random errors in data processing shows the impact of the uncertainties on the final result.

- Experimental design and procedure usually lead to systematic errors in measurement, which cause a deviation in a particular direction.

- Repeat trials and measurements will reduce random errors but not systematic errors

Graphical techniques:

Graphical techniques are an effective means of communicating the effect of an independent variable on a dependent variable, and can lead to determination of physical quantities.

- Sketched graphs have labelled but unscaled axes, and are used to show qualitative trends, such as variables that are proportional or inversely proportional.

- Drawn graphs have labelled and scaled axes, and are used in quantitative measurements

Spectroscopic identification of organic compounds:

The degree of unsaturation or index of hydrogen deficiency (IHD) can be used to determine from a molecular formula the number of rings or multiple bonds in a molecule.

- Mass spectrometry (MS), proton nuclear magnetic resonance spectroscopy (1H NMR) and infrared spectroscopy (IR) are techniques that can be used to help identify compounds and to determine their structure

The number of significant figures in a result is based on the figures given in the data. When adding or subtracting, the final answer should be given to the least number of decimal places. When multiplying or dividing the final answer is given to the least number of significant figures.

- Note that the data value must be recorded to the same precision as the random error.

- SI units should be used throughout the programme.

The electromagnetic spectrum (EMS) is given in the data booklet in section 3. The regions employed for each technique should be understood.

- The operating principles are not required for any of these methods. The data booklet contains characteristic ranges for IR absorptions (section 26), 1H NMR data (section 27) and specific MS fragments (section 28). For 1H NMR, only the ability to deduce the number of different hydrogen (proton) environments and the relative numbers of hydrogen atoms in each environment is required. Integration traces should be covered but splitting patterns are not required.

Usage:

Making quantitative measurements with replicates to ensure reliability—precision, accuracy, systematic, and random errors must be interpreted through replication

The idea of correlation—can be tested in experiments whose results can be displayed graphically

Improvements in instrumentation—mass spectrometry, proton nuclear magnetic resonance and infrared spectroscopy have made identification and structural determination of compounds routine. (1.8) Models are developed to explain certain phenomena that may not be observable—for example, spectra are based on the bond vibration model.

Crash of the Mars Climate Orbiter spacecraft.

- Original results from CERN regarding the speed of neutrinos were flawed.

Graphical representations of data are widely used in diverse areas such as population, finance and climate modelling. Interpretation of these statistical trends can often lead to predictions, and so underpins the setting of government policies in many areas such as health and education.

IR spectroscopy is used in heat sensors and remote sensing in physics.

- Protons in water molecules within human cells can be detected by magnetic resonance imaging (MRI), giving a three-dimensional view of organs in the human body.

All measurement has a limit of precision and accuracy, and this must be taken into account when evaluating experimental results.

Graphs are a visual representation of trends in data.

The inclusion of uncertainties and errors in measurement and results, graphical techniques, and spectroscopic identification of organic compounds in the Pakistan grade 9-12 syllabus is important as it helps students develop a strong foundation in scientific methodology, data analysis, and experimental techniques that are essential for pursuing careers in STEM fields.

 

Nature of Science in Chemistry

 

 

 

History of Chemistry

1. The ancient Egyptians, Greeks, and Chinese all made significant contributions to the field of chemistry.

2. The medieval Islamic world made significant advancements in alchemy, which laid the foundation for modern chemistry.

3. Robert Boyle is considered the ""father of modern chemistry"" for his work in the 17th century on the properties of gases.

4. Antoine Lavoisier is considered the ""father of modern chemistry"" for his work in the 18th century on the nature of matter and the law of conservation of mass.

5. Dmitri Mendeleev created the first periodic table of elements in 1869, which helped to organize the known elements and predict the properties of new ones.

6. Marie Curie was the first woman to win a Nobel Prize, and the first person to win multiple Nobel Prizes (in physics and chemistry) for her work on radioactivity.

7. The discovery of the structure of DNA in the 1950s by James Watson and Francis Crick revolutionized the field of biology and has had farreaching implications in medicine and genetics.

8. Chemistry plays a crucial role in many fields including medicine, agriculture, energy, and materials science.

TOK and Nature of Chemistry

1. Chemistry is an experimental science that combines academic study with the acquisition of practical and investigational skills

2. Chemistry is often called the central science as chemical principles underpin both the physical environment and all biological systems

3. Chemistry is a prerequisite for many other courses in higher education and serves as useful preparation for employment

4. Chemistry has its roots in the study of alchemy, the early days of alchemists who aimed to transmute common metals into gold

5. Observations remain essential at the core of chemistry and scientific processes carried out by the most eminent scientists in the past are the same ones followed by working chemists today and accessible to students in schools

6. The body of scientific knowledge has grown in size and complexity, and the tools and skills of theoretical and experimental chemistry have become specialized

7. Both theory and experiments should be undertaken by all students and should complement each other naturally

8. Allow students to develop traditional practical skills, mathematics skills, interpersonal skills, and digital technology skills.

Scientific Method

1. The scientific method is a process used to conduct scientific research and make discoveries.

2. The steps of the scientific method include:

3. Making observations and asking a question

4. Forming a hypothesis, or an educated guess, about the answer to the question

5. Designing and conducting experiments to test the hypothesis

6. Analyzing the data collected from the experiments

7. Drawing conclusions and determining whether the data supports or disproves the hypothesis

8. The scientific method is based on the principles of observation, experimentation, and replication.

Begin with a broad overview of the history of chemistry, highlighting key figures and contributions made by different civilizations and cultures.

Use visual aids such as timelines and diagrams to help students understand the progression of the field and how different discoveries have built upon one another.

Use real-world examples and case studies to illustrate the significance of chemistry in everyday life and its impact on various industries and fields.

Encourage students to think critically and make connections between historical discoveries and current scientific research.

Incorporate hands-on activities and experiments to demonstrate key concepts and principles.

Create opportunities for students to explore historical texts and primary sources to gain a deeper understanding of the scientific discoveries and the context in which they were made.

Use group work and discussions to help students make connections and discuss the implications of historical discoveries.

Encourage students to think about the contributions of different cultures and groups of people in the history of chemistry.

Show the ethical issues that arise from the history of chemistry, like the Marie Curie story for example.

Finally, use the history of chemistry as a way to inspire students to pursue careers in science, technology, engineering, and mathematics (STEM) fields.

The scientific method helps to minimize bias and error in the research process and ensures that findings can be independently verified.

Scientists use the scientific method to make new discoveries, improve existing technologies, and solve problems in a wide range of fields.

The scientific method is an iterative process, meaning that scientists may repeat steps, modify the hypothesis and make new observations.

The scientific method is not a fixed set of steps, but rather a general approach to conducting research.

It is important for students to understand that scientific research is an ongoing process, and that scientific knowledge is always subject to revision and improvement as new evidence is discovered.

"

 

The purpose of studying Chemistry at the introductory high school level is not only to prepare students for further study in the sciences. Most students will in fact not go on to study further science or STEM fields. The science that they learn in school may well remain their understanding of the subject for the rest of their lives. Hence an introductory physics curriculum must consider what citizens in a democractic society ought to know about the nature of science.

“Nature of Science” (NOS) means teaching about science’s underlying assumptions, and its methodologies. This involves some integrated study of the history of science, and some of the broad concepts from the philosophy of science.

It is important to study NOS because it helps students become critical thinkers about the scientific information the consume from the world around them.

Teaching NOS in the study of Physics, Biology and Chemistry is a cutting-edge international trend. For example:

- The United States has some NOS desired outcomes outlined in its Next Generation Science Standards, which have been co-created by multiple states to foster interdisciplinary science education

- New Zealand has since the last two decades incorporated an NOS module as part of its high school science curricula

- Brazil and Argentina have developed learning standards on NOS

- The IB curriculum substantially incorporates NOS in all its MYP and DP curricula

Teachers with science backgrounds can effectively teach introductory level modules on NOS with the suppport of teacher training, clear examples of assessment expectations and supportive online and textbook materials.The level of knowledege required up to Grade 12 on this topic is nicely elaborated on in the IB DP curriculum guidance documents and these can be adapated.

History of Chemistry:

The rationale behind including this is to provide a historical context for the development of chemistry and to highlight the contributions of different cultures and individuals to the field.

It also helps students understand the evolution of chemistry and how different discoveries and advancements have led to our current understanding of the subject.

TOK and Nature of Chemistry:

The rationale behind including this is to emphasize the importance of both theoretical knowledge and practical skills in the study of chemistry.

It highlights the interdisciplinary nature of chemistry and how it relates to other fields such as biology, physics, and engineering.

It also helps students understand the relevance of chemistry in higher education and in various career paths.

Scientific Method:

The rationale behind including this in is to teach students a systematic approach to conducting scientific research and problem-solving.

It emphasizes the importance of observation, hypothesis testing, and data analysis in the scientific process.

It also helps students understand how scientific knowledge is built through replication and consensus among the scientific community.

 

Stoichiometry
(Physical Chemistry)

- Interpret a balanced chemical equation in terms of interacting moles,

representative particles, masses and volumes of gases (at STP).

- Construct mole ratios from balanced equations for use as conversion factors in

stoichiometric problems.

- Perform stoichiometric calculations with balanced equations using moles,

representative particles, masses and volumes of gases (at STP).

- Identify the limiting reagent in a reaction.

- Knowing the limiting reagent in a reaction, calculate the maximum amount of

product(s) produced and the amount of any unreacted excess reagent.

(

- Given information from which any two of the following may be determined,

calculate the third: theoretical yield, actual yield, percentage yield.

- Calculate the theoretical yield and the percent yield when given the balanced

equation, the amounts of reactants and the actual yield.

- Use the volume ( 22.4 L ) of one mole of a gas at STP to work mole-volume problems

- Calculate the gram molecular mass of a gas from density measurements of gases at STP

- Use the mole to convert among measurements of mass, volume and number of particles(

- Fine out the limiting reactant in a chemical reaction and do the related calculations.

- Perform calculations based on moles, mass, volume and number of particles.

2 Atoms, molecules and stoichiometry

2.1 Relative masses of atoms and molecules

1 define the unified atomic mass unit as one twelfth of the mass of a carbon-12 atom

2 define relative atomic mass relative isotopic mass, relative molecular mass, and relative formula mass

in terms of the unified atomic mass unit

2.2 The mole and the Avogadro constant

1 define and use the term mole in terms of the Avogadro constant

2.3 Formulae

1 write formulae of ionic compounds from ionic charges and oxidation numbers (shown by a Roman numeral),

including:

(a) the prediction of ionic charge from the position of an element in the Periodic Table

(b) recall of the names and formulae for the following ions: zinc, silver, ammonium, nitrate, hydroxide, phosphate, carbonate, sulfate

2 (a) write and construct equations (which should be balanced), including ionic equations (which should not

include spectator ions)

(b) use appropriate state symbols in equations

3 define and use the terms empirical and molecular formula

4 understand and use the terms anhydrous, hydrated and water of crystallisation

5 calculate empirical and molecular formulae, using given data

2.4 Reacting masses and volumes (of solutions and gases)

1 perform calculations including use of the mole concept, involving:

(a) reacting masses (from formulae and equations) including percentage yield calculations

(b) volumes of gases (e.g. in the burning of hydrocarbons)

(c) volumes and concentrations of solutions

(d) limiting reagent and excess reagent

(When performing calculations, Students’ answers should reflect the number of significant figures given or

asked for in the question. When rounding up or down, Students should ensure that significant figures are

neither lost unnecessarily nor used beyond what is justified (see also Mathematical requirements section).)

(e) deduce stoichiometric relationships from calculations such as those in 2.4.1 (a)–(d)

1.2 The mole concept

Understandings:

• The mole is a fixed number of particles and refers to the amount, n, ofsubstance.

• Masses of atoms are compared on a scale relative to 12C and are expressed as relative atomic mass and relative formula/molecular mass

• Molar mass (M) has the units g mol

• The empirical formula and molecular formula of a compound give the simplest

ratio and the actual number of atoms present in a molecule respectively.

Applications and skills:

• Calculation of the molar masses of atoms, ions, molecules and formula units.

• Solution of problems involving the relationships between the number of

particles, the amount of substance in moles and the mass in grams.

• Interconversion of the percentage composition by mass and the empirical

formula.

• Determination of the molecular formula of a compound from its empirical

formula and molar mass.

• Obtaining and using experimental data for deriving empirical formulas from

reactions involving mass changes.

1.3 Reacting masses and volumes

Understandings:

• Reactants can be either limiting or excess.

• The experimental yield can be different from the theoretical yield.

• Avogadro’s law enables the mole ratio of reacting gases to be determined

from volumes of the gases.

• The molar volume of an ideal gas is a constant at specified temperature and

pressure.

• The molar concentration of a solution is determined by the amount of solute

and the volume of solution.

• A standard solution is one of known concentration.

Applications and skills:

• Solution of problems relating to reacting quantities, limiting and excess

reactants, theoretical, experimental and percentage yields.

• Calculation of reacting volumes of gases using Avogadro’s law.

• Solution of problems and analysis of graphs involving the relationship

between temperature, pressure and volume for a fixed mass of an ideal gas.

• Solution of problems relating to the ideal gas equation.

• Explanation of the deviation of real gases from ideal behaviour at low

temperature and high pressure.

• Obtaining and using experimental values to calculate the molar mass of a gasfrom the ideal gas equation.

• Solution of problems involving molar concentration, amount of solute and

volume of solution.

• Use of the experimental method of titration to calculate the concentration of a

solution by reference to a standard solution.

1. Understanding of balanced chemical equations in terms of moles, representative particles, masses, and volumes of gases (at STP).

2. Ability to calculate mole ratios from balanced equations for use in stoichiometric problems.

3. Ability to perform stoichiometric calculations using moles, representative particles, masses, and volumes of gases (at STP).

4. Understanding of limiting reagents and how to calculate the maximum amount of product and amount of any unreacted excess reagent.

5. Ability to calculate theoretical yield, actual yield, and percentage yield when given appropriate information.

6. Understanding of the volume of one mole of a gas at STP and how to use it in mole-volume problems.

7. Understanding of how to calculate the gram molecular mass of a gas from density measurements at STP.

8. Understanding of how to convert measurements of mass, volume, and number of particles using moles.

9. Understanding of the mole and Avogadro's constant and how to use it to define moles in terms of the Avogadro constant.

10. Understanding of how to write ionic compounds formula from ionic charges and oxidation numbers

11. Understanding of how to write balanced equations, including ionic equations, and use appropriate state symbols in equations.

12. Understanding of the terms empirical and molecular formula, anhydrous, hydrated, and water of crystallization.

13. Ability to calculate empirical and molecular formulae using given data.

14. Understanding of reacting masses and volumes of solutions and gases and ability to perform calculations involving reacting masses, volumes of gases, volumes and concentrations of solutions, limiting reagent and excess reagent, percentage yield calculations.

15. Understanding the mole concept, understanding the mole is a fixed number of particles, the relative atomic mass, relative isotopic mass, relative molecular mass, relative formula mass, molar mass, empirical and molecular formula, ability to calculate molar masses of atoms, ions, molecules, and formula units, ability to solve problems involving the relationships between the number of particles, the amount of substance in moles, and the mass in grams, ability to interconvert the percentage composition by mass and the empirical formula.

• The value of the Avogadro’s constant (L or NA) is given in the data booklet

• The generally used unit of molar mass (g mo

• Values for the molar volume of an ideal gas are given in the data booklet in

section 2.

• The ideal gas equation,

𝑃𝑃𝑃𝑃 = 𝑛𝑛𝑛𝑛𝑛𝑛

, and the value of the gas constant (R) are

given in the data booklet in sections 1 and 2.

and parts per million (ppm).

-3

, mol dm

-3

• Units of concentration to include: g dm

• The use of square brackets to denote molar concentration is required.

Usage:

Gas volume changes during chemical reactions are responsible for the

inflation of air bags in vehicles and are the basis of many other explosive

reactions, such as the decomposition of TNT (trinitrotoluene).

The concept of percentage yield is vital in monitoring the efficiency of

industrial processes.

Physical and chemical properties depend on the ways in which different atoms combine.

The mole makes it possible to correlate the number of particles with the mass that can be measured.

Mole ratios in chemical equations can be used to calculate reacting ratios by mass and gas volume.

These are already taught in all major curricula. The connection between experimental measurement, validation and theoretical calculations needs to be highlighted and that paradigm shift is important

 

Atomic Structure
(Physical Chemistry)

- Summarize Bohr's atomic theory

- Use Bohr's model for calculating radii of orbits.

- Use Bohr's atomic model for calculating energy of electron in a given

orbit of hydrogen atom.

- Relate energy equation (for electron) to frequency, wave length and

wave number of radiation emitted or absorbed by electron.

- Explain production, properties, types and uses of X-rays.

- Define photon as a unit of radiation energy.

- Describe the concept of orbitals.

- Explain the significance of quantized energies of electrons.

- Distinguish among principal energy levels, energy sub levels, and atomic orbitals.

- Describe the general shapes of s, p, and d orbitals.

- Relate the discrete-line spectrum of hydrogen to energy levels of electrons in the

hydrogen atom.

- Describe the hydrogen atom using the Quantum Theory.

- Use the Aufbau Principle, the Pauli Exclusion Principle, and Hund's Rule to write

the electronic configuration of the elements.

- Describe the orbitals of hydrogen atom in order of increasing energy.

- Explain the sequence of filling of electrons in many electron atoms.

- Write electron configuration of atoms.

- Calculate the frequency given the wavelength or wave number.

- Calculate the energy of a photon associated with a given wavelength or frequency

of radiation.

- Calculate energy differences between different energy levels of the hydrogen

atom.

1.1 Particles in the atom and atomic radius

1 understand that atoms are mostly empty space surrounding a very small, dense nucleus that contains

protons and neutrons; electrons are found in shells in the empty space around the nucleus

2 identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses

3 understand the terms atomic and proton number; mass and nucleon number

4 describe the distribution of mass and charge within an atom

5 describe the behaviour of beams of protons, neutrons and electrons moving at the same velocity in an

electric field

6 determine the numbers of protons, neutrons and electrons present in both atoms and ions given atomic or

proton number, mass or nucleon number and charge

7 state and explain qualitatively the variations in atomic radius and ionic radius across a period and down a

group

1.2 Isotopes

1 define the term isotope in terms of numbers of protons and neutrons

2 understand the notation for isotopes

3 state that and explain why isotopes of the same element have the same chemical properties

4 state that and explain why isotopes of the same element have different physical properties, limited to mass

and density

1.3 Electrons, energy levels and atomic orbitals

In 1.3 each atom or ion described will be in the ground state. Only the elements hydrogen to krypton will be

assessed.

1 understand the terms:

- shells, sub-shells and orbitals

- principal quantum number (n)

- ground state, limited to electronic configuration

2 describe the number of orbitals making up s, p and d sub-shells, and the number of electrons that can fill s, p

and d sub-shells

3 describe the order of increasing energy of the sub-shells within the first three shells and the 4s and 4p

sub-shells

4 describe the electronic configurations to include the number of electrons in each shell, sub-shell and orbital

5 explain the electronic configurations in terms of energy of the electrons and inter-electron repulsion

6 determine the electronic configuration of atoms and ions given the atomic or proton number and charge,

using either of the following conventions

7 understand and use the electrons in boxes notation

8 describe and sketch the shapes of s and p orbitals

9 describe a free radical as a species with one or more unpaired electrons

1.4 Ionisation energy

In 1.4 each atom or ion described will be in the ground state. Only the elements hydrogen to krypton will be

assessed.

1 define and use the term first ionisation energy, IE

2 construct equations to represent first, second and subsequent ionisation energies

3 identify and explain the trends in ionisation energies across a period and down a group of the Periodic Table

4 identify and explain the variation in successive ionisation energies of an element

5 understand that ionisation energies are due to the attraction between the nucleus and the outer electron

6 explain the factors influencing the ionisation energies of elements in terms of nuclear charge, atomic/ionic

radius, shielding by inner shells and sub-shells and spin-pair repulsion

7 deduce the electronic configurations of elements using successive ionisation energy data

8 deduce the position of an element in the Periodic Table using successive ionisation energy data

Understandings:

• Atoms contain a positively charged dense nucleus composed of protons and

neutrons (nucleons).

• Negatively charged electrons occupy the space outside the nucleus.

• The mass spectrometer is used to determine the relative atomic mass of an

element from its isotopic composition.

Applications and skills:

• Use of the nuclear symbol notation

𝑋𝑋

𝑍𝑍

𝐴𝐴

to deduce the number of protons,

neutrons and electrons in atoms and ions.

• Calculations involving non-integer relative atomic masses and abundance of

isotopes from given data, including mass spectra.

2.2 Electron configuration

Understandings:

• Emission spectra are produced when photons are emitted from atoms as

excited electrons return to a lower energy level.

• The line emission spectrum of hydrogen provides evidence for the existence

of electrons in discrete energy levels, which converge at higher energies.

• The main energy level or shell is given an integer number, n, and can hold a

maximum number of electrons, 2n

.

2

• A more detailed model of the atom describes the division of the main energy

level into s, p, d and f sub-levels of successively higher energies.

• Sub-levels contain a fixed number of orbitals, regions of space where there is

a high probability of finding an electron.

• Each orbital has a defined energy state for a given electronic configuration

and chemical environment and can hold two electrons of opposite spin.

Applications and skills:

• Description of the relationship between colour, wavelength, frequency and

energy across the electromagnetic spectrum.

• Distinction between a continuous spectrum and a line spectrum.

• Description of the emission spectrum of the hydrogen atom, including the

relationships between the lines and energy transitions to the first, second and

third energy levels.

• Recognition of the shape of an s atomic orbital and the p

x

, p

y

and p

z

atomic

orbitals.

• Application of the Aufbau principle, Hund’s rule and the Pauli exclusion

principle to write electron configurations for atoms and ions up to Z = 36.

• Details of the electromagnetic spectrum are given in the data booklet in

section 3.

• The names of the different series in the hydrogen line emission spectrum are

) and condensed electron

4

3p

2

3s

6

2p

2

2s

2

Guidance:

not required.

• Full electron configurations (eg 1s

4

3p

2

) should be covered.

Orbital diagrams should be used to represent the character and relative energy of

configurations (eg [Ne] 3s

orbitals. Orbital diagrams refer to arrow-in-box diagrams, such as the one given

• The electron configurations of Cr and Cu as exceptions should be covered.

below.

1. Describe the structure of atom as a central positively charged nucleus surrounded by negatively charged cloud of electrons due to electrostatic attraction

- understand that, unlike orbits, shells and subshells are energy levels of electrons and a bigger shell implies greater energy and average distance from nucleus

- electrons are quantum particles with probabilistic paths whose exact paths and locations cannot be mapped (with reference to the uncertainty principle)

- nucleus is made of protons and neutrons held together by strong force

- understand that atomic model is a model to aid understanding and if an atom were to be 'photographed' it will be a fuzzy cloud

2. Identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses

3. Understand the terms atomic and proton number; mass and nucleon number

4. Describe the distribution of mass and charge within an atom

5. Describe the behavior of beams of protons, neutrons and electrons moving at the same velocity in an electric field

6. Determine the numbers of protons, neutrons and electrons present in both atoms and ions given atomic or proton number, mass or nucleon number and charge

7. Explain qualitatively the variations in atomic radius and ionic radius across a period and down a group

8. Define the term isotope in terms of numbers of protons and neutrons

9. Understand the notation for isotopes

10. State that and explain why isotopes of the same element have the same chemical properties and different physical properties, limited to mass and density

11. Understand the terms: shells, sub-shells and orbitals, principal quantum number (n), ground state, limited to electronic configuration

12. Describe the number of orbitals making up s, p and d sub-shells, and the number of electrons that can fill s, p and d sub-shells

13. Describe the order of increasing energy of the sub-shells within the first three shells and the 4s and 4p sub-shells

14. Describe the electronic configurations to include the number of electrons in each shell, sub-shell and orbital

15. Explain the electronic configurations in terms of energy of the electrons and inter-electron repulsion

16. Determine the electronic configuration of atoms and ions given the atomic or proton number and charge, using either of the following conventions

17. Understand and use the electrons in boxes notation

18. Describe and sketch the shapes of s and p orbitals

19. Describe a free radical as a species with one or more unpaired electrons

20. Understand the concept of ionization energy and its trends across a period and down a group of the Periodic Table and the variation in successive ionization energies of an element

21. Understand that ionization energies are due to the attraction between the nucleus and the outer electron

22. Explain the factors influencing the ionization energies of elements in terms of nuclear charge, atomic/ionic radius, shielding by inner shells and sub-shells and spin-pair repulsion

23. Deduce the electronic configurations of elements using successive ionization energy data

24. Deduce the position of an element in the Periodic Table using successive ionization energy data

25. Use mass spectrometer to determine the relative atomic mass of an element from its isotopic composition.

26. Perform calculations involving non-integer relative atomic masses and abundance of isotopes from given data, including mass spectra.

27. Understand the concept of emission spectra and use it to deduce the electronic configuration of elements.

• Relative masses and charges of the subatomic particles should be known,

actual values are given in the data booklet. The mass of the

electron can be considered negligible.

• Specific examples of isotopes need not be learned.

• The operation of the mass spectrometer is not required.

students should be able to understand the structure of atoms, including the composition of the nucleus and the distribution of electrons in shells around the nucleus. They should also be able to describe the behavior of beams of protons, neutrons, and electrons in an electric field, and calculate the numbers of protons, neutrons, and electrons present in both atoms and ions based on given information.

students should be able to define isotopes and understand the notation used to describe them. They should also be able to explain why isotopes of the same element have the same chemical properties but different physical properties.

students should be able to understand terms such as shells, sub-shells, and orbitals, and describe the electronic configurations of atoms and ions. They should also be able to describe the order of increasing energy of the sub-shells within the first three shells and the 4s and 4p sub-shells.

students should be able to define and use the term "first ionization energy" (IE), and construct equations to represent first, second, and subsequent ionization energies. They should also be able to explain the factors influencing the ionization energies of elements in terms of nuclear charge, atomic/ionic radius, shielding by inner shells and sub-shells, and spin-pair repulsion.

Usage:

• Radioisotopes are used in nuclear medicine for diagnostics, treatment and

research, as tracers in biochemical and pharmaceutical research, and as

“chemical clocks” in geological and archaeological dating.

• PET (positron emission tomography) scanners give three-dimensional images

of tracer concentration in the body, and can be used to detect cancers.

Absorption and emission spectra are widely used in astronomy to analyse

light from stars.

Atomic absorption spectroscopy is a very sensitive means of determining the

presence and concentration of metallic elements.

Fireworks—emission spectra.

The mass of an atom is concentrated in its minute, positively charged nucleus.

: The electron configuration of an atom can be deduced from its atomic number.

The quantized nature of energy transitions is related to the energy states of electrons in atoms and molecules.

These SLOs are based on what is commonly taught in all major curricula. It has been noted that many false ideas perpetuate at the foundational level which usually make the subject challenging at higher levels. The idea that atomic structure is akin to solar system and shells are actual paths, which might be a deent introductory model, reaffirms the idea of a deterministic world. It is important to introduce the probabilistic nature of reality, and that has been added.

 

Chemical Bonding
(Physical Chemistry)

- Use VSEPR and VBT theories to describe the shapes of simple covalent molecules.

- Describe the features of sigma and pi bonds.

- Describe the shapes of simple molecules using orbital hybridization.

- Determine the shapes of some molecules from the number of bonded pairs and lone pairs of electrons around the central atom.

- Define bond energies and explain how they can be used to compare bond strengths of different chemical bonds.

- Predict the molecular polarity from the shapes of molecules.

- Describe how knowledge of molecular polarity can be used to explain some physical and chemical properties of molecules. (Analyzing

- Describe the change in bond lengths of hetero-nuclear molecules due to difference in Electronegativity values of bonded atoms.

- Describe the difference among molecular, network and metallic solids.

- Explain what is meant by the term ionic character of a covalent bond.

- Use ball and stick models to represent different molecular shapes.

- Guess the physical state of molecule form its structure

3 Chemical bonding

3.1 Electronegativity and bonding

1 define electronegativity as the power of an atom to attract electrons to itself

2 explain the factors influencing the electronegativities of the elements in terms of nuclear charge, atomic

radius and shielding by inner shells and sub-shells

3 state and explain the trends in electronegativity across a period and down a group of the Periodic Table

4 use the differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds

(the presence of covalent character in some ionic compounds will not be assessed) (Pauling electronegativity

values will be given where necessary)

3.2 Ionic bonding

1 define ionic bonding as the electrostatic attraction between oppositely charged ions (positively charged

cations and negatively charged anions)

2 describe ionic bonding including the examples of sodium chloride, magnesium oxide and calcium fluoride

3.3 Metallic bonding

1 define metallic bonding as the electrostatic attraction between positive metal ions and delocalised electrons

3.4 Covalent bonding and coordinate (dative covalent) bonding

1 define covalent bonding as electrostatic attraction between the nuclei of two atoms and a shared pair of

electrons

(a) describe covalent bonding in molecules including:

3.4 Covalent bonding and coordinate (dative covalent) bonding (continued)

(b) understand that elements in period 3 can expand their octet including in the compounds sulfur dioxide, phosphorus pentachloride, and sulfur hexafluoride

(c) describe coordinate (dative covalent) bonding, including in the reaction between ammonia and hydrogen

chloride gases to form the ammonium ion and in the aluminum trichloride molecule

2 (a) describe covalent bonds in terms of orbital overlap giving σ and π bonds:

- σ bonds are formed by direct overlap of orbitals between the bonding atoms

- π bonds are formed by the sideways overlap of adjacent p orbitals above and below the σ bond

(b) describe how the σ and π bonds form in molecules including H

(c) use the concept of hybridisation to describe sp, sp2 and sp3 orbitals

3 (a) define the terms:

- bond energy as the energy required to break one mole of a particular covalent bond in the gaseous

state

- bond length as the internuclear distance of two covalently bonded atoms

(b) use bond energy values and the concept of bond length to compare the reactivity of covalent molecules

3.5 Shapes of molecules

1 state and explain the shapes of, and bond angles in, molecules by using VSEPR theory, including simple examples:

2 predict the shapes of, and bond angles in, molecules and ions analogous to those specified in 3.5.1

3.6 Intermolecular forces, electronegativity and bond properties

1 (a) describe hydrogen bonding, limited to molecules containing N–H and O–H groups, including ammonia

and water as simple examples

(b) use the concept of hydrogen bonding to explain the anomalous properties of H

O (ice and water):

- its relatively high melting and boiling points

- its relatively high surface tension

- the density of the solid ice compared with the liquid water

2 use the concept of electronegativity to explain bond polarity and dipole moments of molecules

3 (a) describe van der Waals’ forces as the intermolecular forces between molecular entities other than

2

those due to bond formation, and use the term van der Waals’ forces as a generic term to describe all

intermolecular forces

(b) describe the types of van der Waals’ force:

- instantaneous dipole – induced dipole (id-id) force, also called London dispersion forces

- permanent dipole – permanent dipole (pd-pd) force, including hydrogen bonding

(c) describe hydrogen bonding and understand that hydrogen bonding is a special case of

permanent dipole – permanent dipole force between molecules where hydrogen is bonded to a highly

electronegative atom

4 state that, in general, ionic, covalent and metallic bonding are stronger than intermolecular forces

3.7 Dot-and-cross diagrams

1 use dot-and-cross diagrams to illustrate ionic, covalent and coordinate bonding including the representation

of any compounds stated in 3.4 and 3.5 (dot-and-cross diagrams may include species with atoms which

have an expanded octet or species with an odd number of electrons)

4.1 Ionic bonding and structure

Understandings:

• Positive ions (cations) form by metals losing valence electrons.

• Negative ions (anions) form by non-metals gaining electrons.

• The number of electrons lost or gained is determined by the electron

configuration of the atom.

• The ionic bond is due to electrostatic attraction between oppositely charged

ions.

• Under normal conditions, ionic compounds are usually solids with lattice

structures.

Applications and skills:

• Deduction of the formula and name of an ionic compound from its component

ions, including polyatomic ions.

• Explanation of the physical properties of ionic compounds (volatility, electrical

conductivity and solubility) in terms of their structure.

4.2. Covalent bonding

Understandings:

• A covalent bond is formed by the electrostatic attraction between a shared pair

of electrons and the positively charged nuclei.

• Single, double and triple covalent bonds involve one, two and three shared

pairs of electrons respectively.

• Bond length decreases and bond strength increases as the number of shared

electrons increases.

• Bond polarity results from the difference in electronegativities of the bonded

atoms.

Applications and skills:

• Deduction of the polar nature of a covalent bond from electronegativity values.

Guidance:

• Bond polarity can be shown either with partial charges, dipoles or vectors.

• Electronegativity values are given in the data booklet in section 8.

4.3 Covalent structures

Understandings:

• Lewis (electron dot) structures show all the valence electrons in a covalently

bonded species.

• The “octet rule” refers to the tendency of atoms to gain a valence shell with a

total of 8 electrons.

• Some atoms, like Be and B, might form stable compounds with incomplete

octets of electrons.

• Resonance structures occur when there is more than one possible position for

a double bond in a molecule.

• Shapes of species are determined by the repulsion of electron pairs according

to VSEPR theory.

• Carbon and silicon form giant covalent/network covalent structures.

Applications and skills:

• Deduction of Lewis (electron dot) structure of molecules and ions showing all

valence electrons for up to four electron pairs on each atom.

• The use of VSEPR theory to predict the electron domain geometry and the

molecular geometry for species with two, three and four electron domains.

• Prediction of bond angles from molecular geometry and presence of nonbondingpairs of electrons.

• Prediction of molecular polarity from bond polarity and molecular geometry.

• Deduction of resonance structures, examples include but are not limited to

• Explanation of the properties of giant covalent compounds in terms of their

structures.

4.4 Intermolecular forces

Understandings:

• Intermolecular forces include London (dispersion) forces, dipole-dipole forces

and hydrogen bonding.

• The relative strengths of these interactions are London (dispersion) forces <

dipoledipole

forces

< hydrogen

bonds.

Applications and skills

• Deduction of the types of intermolecular force present in substances, based on

their structure and chemical formula.

• Explanation of the physical properties of covalent compounds (volatility,

electrical conductivity and solubility) in terms of their structure and

intermolecular forces.

4.5 Metallic bonding

Understandings:

• A metallic bond is the electrostatic attraction between a lattice of positive ions

and delocalized electrons.

• The strength of a metallic bond depends on the charge of the ions and the

radius of the metal ion.

• Alloys usually contain more than one metal and have enhanced properties.

Applications and skills:

• Explanation of electrical conductivity and malleability in metals.

• Explanation of trends in melting points of metals.

• Explanation of the properties of alloys in terms of non-directional bonding.

1. Define electronegativity as the power of an atom to attract electrons to itself

2. Explain the factors influencing the electronegativities of the elements in terms of nuclear charge, atomic radius and shielding by inner shells and sub-shells

3. State and explain the trends in electronegativity across a period and down a group of the Periodic Table

4. Use the differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds

5. Define ionic bonding as the electrostatic attraction between oppositely charged ions (positively charged cations and negatively charged anions) and describe ionic bonding including the examples of sodium chloride, magnesium oxide and calcium fluoride

6. Define metallic bonding as the electrostatic attraction between positive metal ions and delocalized electrons

7. Define covalent bonding as electrostatic attraction between the nuclei of two atoms and a shared pair of electrons, describe covalent bonding in molecules, use the concept of hybridization to describe sp, sp2 and sp3 orbitals and use bond energy values and the concept of bond length to compare the reactivity of covalent molecules

8. State and explain the shapes of, and bond angles in, molecules by using VSEPR theory, predict the shapes of, and bond angles in, molecules and ions analogous to those specified

9. Describe the types of van der Waals’ force:

- instantaneous dipole – induced dipole (id-id) force, also called London dispersion forces

- permanent dipole – permanent dipole (pd-pd) force, including hydrogen bonding

- Hydrogen bonding as a special case of permanent dipole – permanent dipole force between molecules where hydrogen is bonded to a highly electronegative atom

10. Describe hydrogen bonding, limited to molecules containing N–H and O–H groups, including ammonia and water as simple examples

11. Use the concept of hydrogen bonding to explain the anomalous properties of H2O (ice and water)

12. Use the concept of electronegativity to explain bond polarity and dipole moments of molecules

13. Describe van der Waals’ forces as the intermolecular forces between molecular entities and explain the types of van der Waals’ force

14. State that, in general, ionic, covalent and metallic bonding are stronger than intermolecular forces

15. Use dot-and-cross and lewis dot diagrams to show the arrangement of electrons in covalent molecules and ions.

• The term “electron domain” should be used in place of “negative charge centre”.

• Electron pairs in a Lewis (electron dot) structure can be shown as dots,crosses, a dash or any combination.

• Allotropes of carbon (diamond, graphite, graphene, C60 buckminsterfullerene) and SiO2 should be covered.

• Coordinate covalent bonds should be covered.

• Bond polarity can be shown either with partial charges, dipoles or vectors.

• Electronegativity values are given in the data booklet in section 8.

Students should be familiar with the names of polyatomic ions

The term “London (dispersion) forces” refers to instantaneous induced dipoleinduced dipole forces that exist between any atoms or groups of atoms and should be used for non-polar entities. The term “van der Waals” is an inclusive term, which includes dipole–dipole, dipole-induced dipole and London (dispersion) forces.

• Trends should be limited to s- and p-block elements.

• Examples of various alloys should be covered.

students should be able to define electronegativity and bonding, use the differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds, describe ionic bonding, metallic bonding, and covalent bonding, and use the concept of hybridization to describe sp, sp2, and sp3 orbitals.

students should be able to use VSEPR theory to predict the shapes of molecules and bond angles.

For the topic of "3.6 Intermolecular forces, electronegativity and bond properties," students should be able to describe hydrogen bonding, van der Waals' forces, and use the concept of electronegativity to explain bond polarity and dipole moments of molecules.

Usage:

Ionic liquids are efficient solvents and electrolytes used in electric power

sources and green industrial processes.

Microwaves—cooking with polar molecules.

use of X-ray crystallography in structural determinations

Ionic compounds consist of ions held together in lattice structures by ionic bonds.

Covalent compounds form by the sharing of electrons.

Lewis (electron dot) structures show the electron domains in the valence shell and are used to predict molecular shape.

: The physical properties of molecular substances result from different types of forces between their molecules.

Metallic bonds involve a lattice of cations with delocalized electrons.

Larger structures and more in-depth explanations of bonding systems often require more sophisticated concepts and theories of bonding.

Hybridization results from the mixing of atomic orbitals to form the same number of new equivalent hybrid orbitals that can have the same mean energy as

the contributing atomic orbitals.

Currently bonds are shown as processes that occur at the atomic level to attain noble-gas configuration. While this allows students to easily predict the result of chemical bodning, it also gives rise to the idea that chemistry is replete with exceptions which must be memorized as is. By introducing the underlying reasons such as electronegativity as a result of atomic size and strucuture, bonds as electrostatic forces of attraction, intermolecular forces as dipole formation due to distortion of electron cloud, it is hoped that students will develop an intuition that is closer to how atomic model predicts it to be.

- using plasticizers

- controlling biodegradability

- melting points of cis-/trans- fats

- understanding of intermolecular forces to work with molecules in the body

- use of metals in nanotechnology

- water

States and Phases of Matter
(Physical Chemistry)

States of Matter I: Gases

- List the postulates of Kinetic Molecular Theory.

- Describe the motion of particles of a gas according to Kinetic Theory. State the values of standard temperature and pressure (STP).

- Relate temperature to the average kinetic energy of the particles in a substance.

- Use Kinetic Theory to explain gas pressure.

- Describe the effect of change in pressure on the volume of gas. Describe the effect of change in temperature on the volume of gas.

- Explain the significance of absolute zero, giving its value in degree Celsius and Kelvin.

- State and explain the significance of Avogadro's Law.

- Derive Ideal Gas Equation using Boyle's, Charles' and Avogadro's law.

- Explain the significance and different units of ideal gas constant. Distinguish between real and ideal gases.

- Explain why real gases deviate from the gas laws.

- Define and describe the properties of Plasma.

- Derive new form of Gas Equation with volume and pressure corrections for real gases.

- State and use Graham's Law of Diffusion.

- State and use Dalton's Law of Partial Pressures.

- Describe some of the implications of the Kinetic Molecular Theory, such as the velocity of molecules and Graham's Law.

- Explain Lind's method for the liquefaction of gases.

- Define pressure and give its various units.

- Define and explain plasma formation.

- Derive new form of Gas Equation with volume and pressure corrections for real gases.

- State and use Graham's Law of Diffusion.

- State and use Dalton's Law of Partial Pressures.

- Describe some of the implications of the Kinetic Molecular Theory, such as the velocity of molecules and Graham's Law.

- Explain Lind's method for the liquefaction of gases.

- Define pressure and give its various units.

- Define and explain plasma formation.

- lnterconvert pressure in pascals, kilopascals, atmospheres and bar.

- Calculate the partial pressure of a gas collected over water.

- Calculate the new volume of a gas when the pressure of the gas changes.

- Use the combined gas law in calculations.

- Determine the molar volume of the gas under various conditions. Apply the ideal gas laws to calculate the pressure or the volume of a gas.

States of Matter II: Liquids

- Describe simple properties of liquids e.g., diffusion, compression, expansion,

motion of molecules, spaces between them, intermolecular forces and kinetic

energy based on Kinetic Molecular Theory.

- Explain applications of dipole-dipole forces, hydrogen bonding and London forces.

- Explain physical properties of liquids such as evaporation, vapour pressure, boiling

point, viscosity and surface tension.

- Use the concept of Hydrogen bonding to explain the following properties of water:

high surface tension, high specific heat, low vapor pressure, high heat of vaporization, and high boiling point. And anomalous behaviour of water when its density shows maximum at 4 degree centigrade

- Define molar heat of fusion and molar heat of vaporization.

- Describe how heat of fusion and heat of vaporization affect the particles that make

up matter.

- Relate energy changes with changes in intermolecular forces.

- Define dynamic equilibrium between two physical states.

- Describe liquid crystals and give their uses in daily life.

- Differentiate liquid crystals from pure liquids and crystalline solids.

- Identify types of intermolecular attractions between the molecules of a liquid from

a given list of liquids based on its molecular structures.

- Compare and explain the volatility of different liquids at same temperature based

on intermolecular forces.

States of Matter Ill: Solids

- Describe simple properties of solids e.g., diffusion, compression, expansion,

motion of molecules, spaces between them, intermolecular forces and kinetic

energy based on kinetic molecular theory. (Undetanding)

- Differentiate between amorphous and crystalline solids.

- Describe properties of crystalline solids like geometrical shape, melting point,

cleavage planes, habit of a crystal, crystal growth, anisotropy, symmetry,

isomorphism, polymorphism, allotropy and transition temperature.

- Use oxygen and sulphur to define allotropes.

- Explain the significance of the unit cell to the shape of the crystal using NaCl as an

example.

- Name three types of packing arrangements and draw or construct models of them .

- Name three factors that affect the shape of an ionic crystal.

- Define lattice energy.

- Differentiate between ionic, covalent, molecular and metallic crystalline solids .

- Explain the low density and high heat of fusion of ice.

- Define and ex lain molecular and metallic solids.

- List some common amorphous solids encountered in daily life.

- Explain why a compound like CaCb will fluctuate in mass from day to day because

of humidity.

- Purif saline water by repeated freezing. (Appplying)

- List the characteristics of colloids and suspensions that distinguish them from solutions.

- Define hydrophilic and hydrophobic molecules.

- Explain the nature of solutions in liquid phase giving examples of completely miscible, partially miscible and immiscible liquid-liquid solutions.

- Explain the effect of temperature on solubility and interpret the solubility graph.

- Express solution concentration in terms of mass percent, molality, molarity, parts per million, billion and trillion and mole fraction.

- Define the terms colligative.

- Describe on a particle basis why a solution has a lower vapor pressure than the pure solvent.

- Explain on a particle basis how the addition of a solute to a pure solvent causes an elevation of the boiling point and depression of the freezing point of the resultant solution.

- Describe the role of solvation in the dissolving process.

- Define the term water of hydration.

- Explain concept of solubility and how it applies to solution saturation.

- Distinguish between the solvation of ionic species and molecular substances.

- List three factors that accelerate the dissolution process.

- Define heat of solution and apply this concept to the hydration of ammonium nitrate crystals.

- Explain how solute particles may alter the colligative properties.

- Explain osmotic pressure, reverse osmosis, and give their daily life applications.

- Describe types of colloids and their properties.

- List some colligative properties of liquids.

- Perform calculations involving percent (volume/volume) and percent (mass/volume) solutions.

- Calculate the molality of a solution.

- Calculate the freezing point depression and the boiling point elevation of aqueous solutions.

- Calculate molar mass of a substance using ebullioscopic and cryoscopic methods.

- Calculate the percent of water in a given hydrate.

- Explain the phenomenon freezing in a mixture of ice and salt.

4 States of matter

4.1 The gaseous state: ideal and real gases and pV = nRT

1 explain the origin of pressure in a gas in terms of collisions between gas molecules and the wall of the

container

2 understand that ideal gases have zero particle volume and no intermolecular forces of attraction

3 state and use the ideal gas equation pV = nRT in calculations, including in the determination of Mr

4.2 Bonding and structure

1 describe, in simple terms, the lattice structure of a crystalline solid which is:

(a) giant ionic, including sodium chloride and magnesium oxide

(b) simple molecular, including iodine, buckminsterfullerene C

and ice

(c) giant molecular, including silicon(IV) oxide, graphite and diamond

(d) giant metallic, including copper

2 describe, interpret and predict the effect of different types of structure and bonding on the physical

properties of substances, including melting point, boiling point, electrical conductivity and solubility

3 deduce the type of structure and bonding present in a substance from given information

 

1. Describe simple properties of liquids e.g., diffusion, compression, expansion, motion of molecules, spaces between them, intermolecular forces and kinetic energy based on Kinetic Molecular Theory.

2. Explain applications of dipole-dipole forces, hydrogen bonding and London forces.

3. Describe physical properties of liquids such as evaporation, vapor pressure, boiling point, viscosity and surface tension.

4. Use the concept of Hydrogen bonding to explain the following properties of water: high surface tension, high specific heat, low vapor pressure, high heat of vaporization, and high boiling point.

5. Define molar heat of fusion and molar heat of vaporization.

6. Describe how heat of fusion and heat of vaporization affect the particles that make up matter.

7. Relate energy changes with changes in intermolecular forces.

8. Define dynamic equilibrium between two physical states.

9. Describe liquid crystals and give their uses in daily life.

10. Differentiate liquid crystals from pure liquids and crystalline solids.

11. Describe simple properties of solids e.g., diffusion, compression, expansion, motion of molecules, spaces between them, intermolecular forces and kinetic energy based on kinetic molecular theory.

12. Differentiate between amorphous and crystalline solids.

13. Describe properties of crystalline solids like geometrical shape, melting point, cleavage planes, habit of a crystal, crystal growth.

Gases: Gases are composed of particles in constant motion and have properties such as pressure, volume, and temperature that can be explained by Kinetic Molecular Theory.

Liquids: Liquids have properties such as viscosity, surface tension, and boiling point that are affected by intermolecular forces such as dipole-dipole, hydrogen bonding, and London forces.

Solids: Solids are composed of particles that are closely packed and in fixed positions, and have properties such as melting point, crystal shape, and cleavage planes.

Plasma: Plasma is a state of matter that is composed of highly energized particles, and has unique properties such as high electrical conductivity and light emission.

Dynamic equilibrium: Dynamic equilibrium is the balance of opposing processes in a system, such as evaporation and condensation in a liquid, and can be described through changes in intermolecular forces and kinetic energy.

Usage:

 

The word phase has been introduced.

Some points about crystal structure have been added.

The hope is that students will appreciate that states and phases are more of a spectrum and less of distinct states limited in their abilities and properties.

 

Energetics & Thermochemistry
(Physical Chemistry)

- Define thermodynamics.

- Classify reactions as exothermic or endothermic.

- Define the terms system, surrounding, boundary, state function, heat, heat capacity, internal energy, work done and enthalpy of a substance.

- Name and define the units of thermal energy.

- Relate a change in enthalpy to the heat of reaction or heat of combustion of a reaction.

- Relate change in internal energy of a system with thermal energy at constant temperature and constant pressure.

- Define bond dissociation energy.

- Use the experimental data to calculate the heat of reaction using a calorimeter.

- Specify conditions for the standard heat of reaction.

- Apply Hess's Law to construct simple energy cycles.

- Describe how heat of combustion can be used to estimate the energy available from foods.

- Explain reaction pathway diagram in terms of enthalpy changes of the reaction. (Born Haber's Cycle)

- Use standard heats of formation to calculate the enthalpy change of a reaction.

- Determine the heat of a reaction which is experimentally inaccessible from the heats of a set of reaction which are experimentally measurable.

- Perform calculations involving energy cycles related to Hess's Law.

- Calculate lattice energy and enthalpy of formation of NaCI and MgO from given set of appropriate data.

5 Chemical energetics

5.1 Enthalpy change, ΔH

1 understand that chemical reactions are accompanied by enthalpy changes and these changes can be

exothermic (ΔH is negative) or endothermic (ΔH is positive)

2 construct and interpret a reaction pathway diagram, in terms of the enthalpy change of the reaction and of

the activation energy

3 define and use the terms:

(a) standard conditions (this syllabus assumes that these are 298 K and 101 kPa

(b) enthalpy change with particular reference to: reaction, formation, combustion, neutralisation

4 understand that energy transfers occur during chemical reactions because of the breaking and making of bonds

5 use bond energies (ΔH positive, i.e. bond breaking) to calculate enthalpy change of reaction, ΔH

6 understand that some bond energies are exact and some bond energies are averages

7 calculate enthalpy changes from appropriate experimental results, including the use of the relationships q = mcΔT and ΔH = –mcΔT/n

23.1 Lattice energy and Born-Haber cycles

1 define and use the terms:

(a) enthalpy change of atomisation, ΔH

(b) lattice energy, ΔH

(the change from gas phase ions to solid lattice)

2 (a) define and use the term first electron affinity, EA

(b) explain the factors affecting the electron affinities of elements

(c) describe and explain the trends in the electron affinities of the Group 16 and Group 17 elements

3 construct and use Born–Haber cycles for ionic solids

(limited to +1 and +2 cations, –1 and –2 anions)

4 carry out calculations involving Born–Haber cycles

5 explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a

lattice energy

23.2 Enthalpies of solution and hydration

1 define and use the term enthalpy change with reference to hydration,and solution

2 construct and use an energy cycle involving enthalpy change of solution, lattice energy and enthalpy change

of hydration

3 carry out calculations involving the energy cycles in 23.2.2

4 explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of an

enthalpy change of hydration

23.3 Entropy change, ΔS

1 define the term entropy, S, as the number of possible arrangements of the particles and their energy in a

given system

2 predict and explain the sign of the entropy changes that occur:

(a) during a change in state, e.g. melting, boiling and dissolving (and their reverse)

(b) during a temperature change

(c) during a reaction in which there is a change in the number of gaseous molecules

3 calculate the entropy change for a reaction, ΔS, given the standard entropies, S

, of the reactants and

products, ΔS

= ΣS

(products) – ΣS

(reactants)

(use of ΔS

= ΔSsurr

+ ΔSsys

is not required)

1 Measuring energy changes

Understandings:

Heat is a form of energy.

- Temperature is a measure of the average kinetic energy of the particles.

- Total energy is conserved in chemical reactions.

- Chemical reactions that involve transfer of heat between the system and the surroundings are described as endothermic or exothermic.

- The enthalpy change (∆H) for chemical reactions is indicated in kJ mol-1.

- ∆H values are usually expressed under standard conditions, given by ∆H°, including standard states

2 Hess's Law

Understandings:

The enthalpy change for a reaction that is carried out in a series of steps is equal to the sum of the enthalpy changes for the individual steps.

3 Bond enthalpies

Understandings:

Bond-forming releases energy and bond-breaking requires energy.

- Average bond enthalpy is the energy needed to break one mol of a bond in a gaseous molecule averaged over similar compound"

1. Understand that chemical reactions are accompanied by enthalpy changes and these changes can be exothermic (ΔH is negative) or endothermic (ΔH is positive)

2. Construct and interpret a reaction pathway diagram, in terms of the enthalpy change of the reaction and of the activation energy

3. Define and use terms such as standard conditions, enthalpy change, reaction, formation, combustion, neutralisation

4. Understand that energy transfers occur during chemical reactions because of the breaking and making of bonds

5. Use bond energies to calculate enthalpy change of reaction, ΔH

6. Understand that some bond energies are exact and some bond energies are averages

7. Calculate enthalpy changes from appropriate experimental results, including the use of the relationships q = mcΔT and ΔH = –mcΔT/n

8. Define and use terms such as enthalpy change of atomisation, ΔH, lattice energy, ΔH, first electron affinity, EA

9. Explain the factors affecting the electron affinities of elements

10. Describe and explain the trends in the electron affinities of the Group 16 and Group 17 elements

11. Construct and use Born–Haber cycles for ionic solids

12. Carry out calculations involving Born–Haber cycles

13. Explain the effect of ionic charge and ionic radius on the numerical magnitude of a lattice energy

14. Define and use the term enthalpy change with reference to hydration, and solution

15. Construct and use an energy cycle involving enthalpy change of solution, lattice energy and enthalpy change of hydration

16. Carry out calculations involving the energy cycles

17. Explain the effect of ionic charge and ionic radius on the numerical magnitude of an enthalpy change of hydration

18. Define the term entropy, S, as the number of possible arrangements of the particles and their energy in a given system

19. Predict and explain the sign of the entropy changes that occur during a change in state, temperature change and a reaction in which there is a change in the number of gaseous molecules

20. Calculate the entropy change for a reaction, ΔS, given the standard entropies, S, of the reactants and products

21. Understand the concept of heat as a form of energy

22. Understand the relationship between temperature and kinetic energy of particles

23. Understand that total energy is conserved in chemical reactions

24. Understand the concept of endothermic and exothermic reactions

25. Understand the concept of standard conditions and standard states in measuring energy changes

26. Understand the concept of Hess's Law and how to apply it to calculate enthalpy changes in a reaction carried out in multiple steps.

27. Understand the relationship between bond formation and energy, and bond breaking and energy

28. Understand the concept of average bond enthalpy.

• Enthalpy changes of combustion (∆Hc°) and formation (∆Hf°)should be covered.

- Consider reactions in aqueous solution and combustion reactions.

• Enthalpy of formation data can be found in the data booklet in section 12.

- An application of Hess's Law is ∆

𝐻𝐻

reaction

• Bond enthalpy values are given in the data booklet in section 11.

- Average bond enthalpies are only valid for gases and calculations involving bond enthalpies may be inaccurate because they do not take into account intermolecular forces.

It is hoped that ideas of bonds being a storage of energy will be dispelled. Similarly, students will come to appreciate how bond formationa and breaking involves change in energy, and that matter and energy cannot be interchanged in chemical reaction. For example, carbohydrates do not turn into energy when digested, or we lose fat by exhaling it as CO2 and H2O, not by converting it to energy.

What is the relationship between enthalpy changes and temperature in chemical reactions?

How does the first law of thermodynamics apply to chemical transformations?

How can bond breaking and bond forming affect the overall energy change in a chemical reaction?

How can the concept of energy change in a single step reaction be applied to changes involving ionic compounds?

How does the second law of thermodynamics relate to the direction of spontaneous change in chemical reactions?

What is the relationship between the overall transformation and the total entropy in a spontaneous chemical reaction?

How the energy available to do useful work is related to the total entropy of the universe in a spontaneous chemical reaction?

Terms for different enthalpy changes have been added that will help understand Born-Haber and Hess cycle better.

 

Reaction Kinetics
(Physical Chemistry)

- Define chemical kinetics.

- Explain and use the terms rate of reaction, rate equation, order of reaction, rate constant and rate determining step.

- Explain qualitatively factors affecting rate of reaction.

- Given the order with respect to each reactant, write the rate law for the reaction.

- Explain what is meant by the terms activation energy and activated complex.

- Relate the ideas of activation energy and the activated complex to the rate of a

reaction.

- Use the collision theory to explain how the rate of a chemical reaction is influenced by the temperature, concentration, size of molecules and.

- Given a potential energy diagram for a reaction, discuss the reaction mechanism for the reaction.

- Explain effects of concentration, temperature and surface area on reaction rates.

- Explain the significance of the rate-determining step on the overall rate of a multi step reaction.

- Describe the role of the rate constant in the theoretical determination of reaction rate.

- Describe that increase in collision energy by increasing the temperature can improve the collision frequency.

- Define terms catalyst, catalysis, homogeneous catalysis and heterogeneous

catalysis.

- Explain that a catalyst provides a reaction pathway that has low activation energy.

- Describe enzymes as biological catalysts.

- Explain why powdered zinc reacts faster.

- Draw energy diagrams that represent the activation energy and show the effect of a catalyst.

- Calculate initial rate using concentration data.

- Deduce the order of a reaction using the method of initial rates.

8 Reaction kinetics

8.1 Rate of reaction

1 explain and use the term rate of reaction, frequency of collisions, effective collisions and non-effective

collisions

2 explain qualitatively, in terms of frequency of effective collisions, the effect of concentration and pressure

changes on the rate of a reaction

3 use experimental data to calculate the rate of a reaction

8.2 Effect of temperature on reaction rates and the concept of activation energy

1 define activation energy, E

, as the minimum energy required for a collision to be effective

2 sketch and use the Boltzmann distribution to explain the significance of activation energy

3 explain qualitatively, in terms both of the Boltzmann distribution and of frequency of effective collisions, the

A

effect of temperature change on the rate of a reaction

8.3 Homogeneous and heterogeneous catalysts

1 explain and use the terms catalyst and catalysis

(a) explain that, in the presence of a catalyst, a reaction has a different mechanism, i.e. one of lower

activation energy

(b) explain this catalytic effect in terms of the Boltzmann distribution

(c) construct and interpret a reaction pathway diagram, for a reaction in the presence and absence of an

effective catalyst

23.4 Gibbs free energy change, ΔG

1 state and use the Gibbs equation ΔG

2 perform calculations using the equation ΔG

= ΔH

– TΔS

= ΔH

– TΔS

3 state whether a reaction or process will be feasible by using the sign of ΔG

4 predict the effect of temperature change on the feasibility of a reaction, given standard enthalpy and

entropy changes

26.1 Simple rate equations, orders of reaction and rate constants

1 explain and use the terms rate equation, order of reaction, overall order of reaction, rate constant, half-life,

rate-determining step and intermediate

2 (a) understand and use rate equations of the form rate = k [A]

m

[B]

n

(for which m and n are 0, 1 or 2)

(b) deduce the order of a reaction from concentration-time graphs or from experimental data relating to

the initial rates method and half-life method

(c) interpret experimental data in graphical form, including concentration-time and rate-concentration

graphs

(d) calculate an initial rate using concentration data

(e) construct a rate equation

3 (a) show understanding that the half-life of a first-order reaction is independent of concentration

(b) use the half-life of a first-order reaction in calculations

4 calculate the numerical value of a rate constant, for example by:

(a) using the initial rates and the rate equation

(b) using the half-life, t

1/2

, and the equation k = 0.693/t

1/2

5 for a multi-step reaction:

(a) suggest a reaction mechanism that is consistent with the rate equation and the equation for the overall

reaction

(b) predict the order that would result from a given reaction mechanism and rate-determining step

(c) deduce a rate equation using a given reaction mechanism and rate-determining step for a given reaction

(d) identify an intermediate or catalyst from a given reaction mechanism

(e) identify the rate determining step from a rate equation and a given reaction mechanism

6 describe qualitatively the effect of temperature change on the rate constant and hence the rate of a reaction

"1 Collision theory and rates of reaction

Understandings:

Species react as a result of collisions of sufficient energy and proper orientation.

- The rate of reaction is expressed as the change in concentration of a particular reactant/product per unit time.

- Concentration changes in a reaction can be followed indirectly by monitoring changes in mass, volume and colour.

- Activation energy (Ea) is the minimum energy that colliding molecules need in order to have successful collisions leading to a reaction.

- By decreasing Ea, a catalyst increases the rate of a chemical reaction, without itself being permanently chemically changed."

1 Understand the concept of collision theory and how it relates to the rate of chemical reactions

2 Explain how changes in concentration and pressure affect the rate of a reaction in terms of frequency of effective collisions

3 Use experimental data to calculate the rate of a reaction

4 Understand the concept of activation energy and its role in chemical reactions

5 Use the Boltzmann distribution to explain the effect of temperature on the rate of a reaction

6 Understand the concept of catalysts and how they increase the rate of a reaction by lowering the activation energy

7 Interpret and construct reaction pathway diagrams, including in the presence and absence of catalysts

8 Understand the relationship between Gibbs free energy change, ΔG, and the feasibility of a reaction

9 Understand and use rate equations, including orders of reaction and rate constants

10 Calculate the numerical value of a rate constant using the initial rates and half-life method

11 Suggest a reaction mechanism that is consistent with a given rate equation and rate-determining step

12 Describe the effect of temperature change on the rate constant and rate of a reaction.

Calculation of reaction rates from tangents of graphs of concentration, volume or mass against time should be covered.

- Students should be familiar with the interpretation of graphs of changes in concentration, volume or mass against time

Usage:

-What is the relationship between collision probability and reaction rate?

-How are rate expressions determined?

-What is the relationship between rate equations and reaction mechanisms?

-What is the relationship between activation energy and temperature's effect on reaction rate?

-What is the activation energy of a reaction and what affects it?

The chemical kinetics topic has been retained . The curriculum focuses on understanding collision theory and how it relates to the rate of chemical reactions. Students are expected to explain how changes in concentration, pressure, and temperature affect the rate of a reaction. They will also learn about activation energy, catalysts, and rate equations, including orders of reaction and rate constants. Additionally, they will understand the relationship between Gibbs free energy change and the feasibility of a reaction. The curriculum emphasizes the use of experimental data to calculate the rate of a reaction and the ability to interpret and construct reaction pathway diagrams.

 

Equilibria
(Basic, Acid-Base, Ionic)
(Physical Chemistry)

- Define chemical equilibrium in terms of a reversible reaction.

- Write both forward and reverse reactions and describe the macroscopic characteristics of each.

- State the necessary conditions for equilibrium and the ways that equilibrium can be recognized.

- Describe the microscopic events that occur when a chemical system is in equilibrium.

- Write the equilibrium expression for a given chemical reaction.

- Relate the equilibrium expression in terms of concentration, partial pressure, number of moles and mole fraction.

- Write expression for reaction quotient.

- Determine if the equilibrium constant will increase or decrease when temperature is changed, given the equation for the reaction.

- Propose microscopic events that account for observed macroscopic changes that

take place during a shift in equilibrium.

- Determine if the reactants or products are favored in a chemical reaction, given the equilibrium constant.

- State Le Chatelier's Principle and be able to apply it to systems in equilibrium with changes in concentration, pressure, temperature, or the addition of catalyst.

- Explain industrial applications of Le Chatelier's Principle using Haber's process as an example.

- Define and explain solubility product.

- Define and explain common ion effect giving suitable examples.

- Calculate the equilibrium constant for a reaction given the equilibrium concentrations of reactants and products.

- Calculate the concentration specified, given the equilibrium constant and appropriate information about the equilibrium concentrations.

- Define Bronsted and Lowery concepts for acids and bases

- Define salts, conjugate acids and conjugate bases.

- Identify conjugate acid-base pairs of Bronsted-Lowery acid and base(

- Explain ionization constant of water and calculate pH and pOH in aqueous medium using given Kw values.

- Use the extent of ionization and the acid dissociation constant, Ka, to distinguish between strong and weak acids.

- Use the extent of ionization and the base dissociation constant, Kb, to distinguish between strong and weak bases.

- Define a buffer, and show with equations how a buffer system works.

- Make a buffered solution and explain how such a solution maintains a constant pH, even with the addition of small amounts of strong acid or strong base.

- Use the concept of hydrolysis to explain why aqueous solutions of some salts are acidic or basic.

- Use concept of hydrolysis to explain why the solution of a salt is not necessarily neutral.

- Define and explain leveling effect.

- Calculate the fourth parameter when given three of four parameters- molarity of base, volume of base, molality of acid, volume of acid - used in a titration experiment, assuming a strong acid and strong base reaction.

- Calculate the [H301. given the Ka and molar concentration of weak acid.

- Calculate concentrations of ions of slightly soluble salts.

- Calculate Ka for the system, given the equilibrium concentrations of a weak acid and the [H30+] in the solution.

- Perform acid-base titrations to calculate molality and strength of given sample solutions.

7 Equilibria

7.1 Chemical equilibria: reversible reactions, dynamic equilibrium

1 (a) understand what is meant by a reversible reaction

(b) understand what is meant by dynamic equilibrium in terms of the rate of forward and reverse reactions

being equal and the concentration of reactants and products remaining constant

(c) understand the need for a closed system in order to establish dynamic equilibrium

2 define Le Chatelier’s principle as: if a change is made to a system at dynamic equilibrium, the position of

equilibrium moves to minimise this change

3 use Le Chatelier’s principle to deduce qualitatively (from appropriate information) the effects of changes in

temperature, concentration, pressure or presence of a catalyst on a system at equilibrium

4 deduce expressions for equilibrium constants in terms of concentrations, K

5 use the terms mole fraction and partial pressure

6 deduce expressions for equilibrium constants in terms of partial pressures, K,

7 deduce expressions to carry out calculations (such calculations will not require the solving of

quadratic equations)

8 calculate the quantities present at equilibrium, given appropriate data

9 state whether changes in temperature, concentration or pressure or the presence of a catalyst affect the

value of the equilibrium constant for a reaction

10 describe and explain the conditions used in the Haber process and the Contact process, as examples of the

importance of an understanding of dynamic equilibrium in the chemical industry and the application of Le

Chatelier’s principle

7.2 Brønsted–Lowry theory of acids and bases

1 state the names and formulae of the common acids, limited to hydrochloric acid, sulfuric acid, nitric acid, ethanoic acid,

2 state the names and formulae of the common alkalis, limited to sodium hydroxide, NaOH, potassium hydroxide, ammonia,

7.2 Brønsted–Lowry theory of acids and bases (continued)

3 describe the Brønsted–Lowry theory of acids and bases

4 describe strong acids and strong bases as fully dissociated in aqueous solution and weak acids and weak

bases as partially dissociated in aqueous solution

5 appreciate that water has pH of 7, acid solutions pH of below 7 and alkaline solutions pH of above 7

6 explain qualitatively the differences in behaviour between strong and weak acids including the reaction with

a reactive metal and difference in pH values by use of a pH meter, universal indicator or conductivity

7 understand that neutralisation reactions occur when H+(aq) and OH–(aq) form H2O(l)

8 understand that salts are formed in neutralisation reactions

9 sketch the pH titration curves of titrations using combinations of strong and weak acids with strong and weak alkalis

10 select suitable indicators for acid-alkali titrations, given appropriate data (pKa values will not be used)

25.1 Acids and bases

1 understand and use the terms conjugate acid and conjugate base

2 define conjugate acid–base pairs, identifying such pairs in reactions

3 define mathematically the terms pH, K

a

, pK

a

and K

w

and use them in calculations (K

and the equation

K

w

= K

a

× K

will not be tested)

4 calculate [H

b

+

(aq)] and pH values for:

(a) strong acids

(b) strong alkalis

(c) weak acids

5 (a) define a buffer solution

(b) explain how a buffer solution can be made

(c) explain how buffer solutions control pH; use chemical equations in these explanations

(d) describe and explain the uses of buffer solutions, including the role of HCO

3

b

in controlling pH in blood

6 calculate the pH of buffer solutions, given appropriate data

7 understand and use the term solubility product, K

8 write an expression for K

9 calculate K

sp

sp

sp

from concentrations and vice versa

10 (a) understand and use the common ion effect to explain the different solubility of a compound in a

solution containing a common ion

(b) perform calculations using K

25.2 Partition coefficients

sp

values and concentration of a common ion

1 state what is meant by the term partition coefficient, K

pc

2 calculate and use a partition coefficient for a system in which the solute is in the same physical state in the

two solvents

3 understand the factors affecting the numerical value of a partition coefficient in terms of the polarities of

the solute and the solvents used

26.2 Homogeneous and heterogeneous catalysts

1 explain that catalysts can be homogeneous or heterogeneous

2 describe the mode of action of a heterogeneous catalyst to include adsorption of reactants, bond weakening

and desorption of products, for example:

(a) iron in the Haber process

(b) palladium, platinum and rhodium in the catalytic removal of oxides of nitrogen from the exhaust gases

of car engines

3 describe the mode of action of a homogeneous catalyst by being used in one step and reformed in a later

step, for example:

(a) atmospheric oxides of nitrogen in the oxidation of atmospheric sulfur dioxide

(b) Fe

2+

or Fe

3+

in the I

/S

2

O

8

2–

reaction

1 Equilibrium

Understandings:

A state of equilibrium is reached in a closed system when the rates of the forward and reverse reactions are equal.

- The equilibrium law describes how the equilibrium constant (Kc) can be determined for a particular chemical reaction.

- The magnitude of the equilibrium constant indicates the extent of a reaction at equilibrium and is temperature dependent.

- The reaction quotient (Q) measures the relative amount of products and reactants present during a reaction at a particular point in time. Q is the equilibrium expression with non-equilibrium concentrations. The position of the equilibrium changes with changes in concentration, pressure, and temperature.

- A catalyst has no effect on the position of equilibrium or the equilibrium constant

2 Theories of acids and bases

Understandings:

A Brønsted–Lowry acid is a proton/H+ donor and a Brønsted–Lowry base is a proton/H+ acceptor.

- Amphiprotic species can act as both Brønsted–Lowry acids and bases.

- A pair of species differing by a single proton is called a conjugate acid-base pair.

3 Properties of acids and bases

Understandings:

Most acids have observable characteristic chemical reactions with reactive metals, metal oxides, metal hydroxides, hydrogen carbonates and carbonates.

- Salt and water are produced in exothermic neutralization reactions.

4 The pH scale

Understandings:

pH = − log[H+(aq)] and [H+] = 10−pH.

- A change of one pH unit represents a 10-fold change in the hydrogen ion concentration [H+].

- pH values distinguish between acidic, neutral and alkaline solutions.

- The ionic product constant,

𝐾𝐾𝑤𝑤

= [H+][OH−] = 10−14 at 298 K

5 Strong and weak acids and bases

Understandings:

Strong and weak acids and bases differ in the extent of ionization.

- Strong acids and bases of equal concentrations have higher conductivities than weak acids and bases.

- A strong acid is a good proton donor and has a weak conjugate base.

- A strong base is a good proton acceptor and has a weak conjugate acid

6 Acid deposition

Understandings:

Rain is naturally acidic because of dissolved CO2 and has a pH of 5.6. Acid deposition has a pH below 5.6.

- Acid deposition is formed when nitrogen or sulfur oxides dissolve in water to form HNO3, HNO2, H2SO4 and H2SO3.

- Sources of the oxides of sulfur and nitrogen and the effects of acid deposition should be covered.

1. Understand what is meant by a reversible reaction and dynamic equilibrium in terms of the rate of forward and reverse reactions being equal and the concentration of reactants and products remaining constant

2. State the necessary conditions for equilibrium and the ways that equilibrium can be recognized.

3. Describe the microscopic events that occur when a chemical system is in equilibrium.

4. Write the equilibrium expression for a given chemical reaction in terms of concentration, partial pressure, number of moles and mole fraction.

5. Propose microscopic events that account for observed macroscopic changes that take place during a shift in equilibrium.

6. Determine if the equilibrium constant will increase or decrease when temperature is changed, given the equation for the reaction.

7. State Le Chatelier's Principle and be able to apply it to systems in equilibrium with changes in concentration, pressure, temperature, or the addition of catalyst.

8. Explain industrial applications of Le Chatelier's Principle using Haber's process as an example.

9. Define and explain solubility product.

10. Define and explain common ion effect giving suitable examples.

11. Define acids and bases in terms of Arrhenius, Bronsted-Lowry and Lewis theory.

12. Use the extent of ionization and the acid dissociation constant, Ka, to distinguish between strong and weak acids. Use the extent of ionization and the base dissociation constant, Kb, to distinguish between strong and weak bases.

13. Define a buffer, and show how a buffer system works.

14. Make a buffered solution and explain how such a solution maintains a constant pH, even with the addition of small amounts of strong acid or strong base.

15. Use the concept of hydrolysis to explain why aqueous solutions of some salts are acidic or basic.

16. Use concept of hydrolysis to explain why the solution of a salt is not necessarily neutral.

17. Define and explain leveling effect.

18. Calculate the fourth parameter when given three of four parameters in a titration experiment, assuming a strong acid and strong base reaction.

19. Calculate the [H30+] given the Ka and molar concentration of weak acid.

20. Calculate concentrations of ions of slightly soluble salts.

21. Perform acid-base titrations to calculate molality and strength of given sample solutions.

22 sketch the pH titration curves of titrations using combinations of strong and weak acids with strong and weak alkalis

23 select suitable indicators for acid-alkali titrations, given appropriate data (pKa values will not be used)

Acid-Base Theory

1 understand and use the terms conjugate acid and conjugate base

2 define conjugate acid–base pairs, identifying such pairs in reactions

3 define mathematically the terms pH, Ka, pKa and Kw and use them in calculations (Kb and the equation

Kw = Ka × Kb will not be tested)

4 calculate [H+(aq)] and pH values for:

(a) strong acids

(b) strong alkalis

(c) weak acids

5 (a) define a buffer solution

(b) explain how a buffer solution can be made

(c) explain how buffer solutions control pH; use chemical equations in these explanations

(d) describe and explain the uses of buffer solutions, including the role of HCO3

– in controlling pH in blood

6 calculate the pH of buffer solutions, given appropriate data

7 understand and use the term solubility product, Ksp

8 write an expression for Ksp

9 calculate Ksp from concentrations and vice versa

10 (a) understand and use the common ion effect to explain the different solubility of a compound in a

solution containing a common ion

(b) perform calculations using Ksp values and concentration of a common ion

The pH scale

Understandings:

pH = − log[H+(aq)] and [H+] = 10−pH.

- A change of one pH unit represents a 10-fold change in the hydrogen ion concentration [H+].

- pH values distinguish between acidic, neutral and alkaline solutions.

- The ionic product constant,

𝐾𝐾𝑤𝑤

= [H+][OH−] = 10−14 at 298 K

Partition Coefficient

1 state what is meant by the term partition coefficient, Kpc

2 calculate and use a partition coefficient for a system in which the solute is in the same physical state in the two solvents

3 understand the factors affecting the numerical value of a partition coefficient in terms of the polarities of the solute and the solvents used

1 Physical and chemical systems should be covered.

- Relationship between Kc values for reactions that are multiples or inverses of one another should be covered.

- Specific details of any industrial process are not required

2 Lewis theory is not required here.

- The location of the proton transferred should be clearly indicated. For example, CH3COOH/CH3COO– rather than C2H4O2/C2H3O2–.

- Students should know the representation of a proton in aqueous solution as both H+ (aq) and H3O+ (aq).

- The difference between the terms amphoteric and amphiprotic should be covered.

3 Bases which are not hydroxides, such as ammonia, soluble carbonates and hydrogen carbonates should be covered.

- The colour changes of different indicators are given in the data booklet.

4 Students will not be assessed on pOH values.

- Students should be concerned only with strong acids and bases in this sub- topic.

- Knowing the temperature dependence of

𝐾𝐾𝑤𝑤

is not required.

- Equations involving H3O+ instead of H+ may be applied.

5 The terms ionization and dissociation can be used interchangeably.

- A data booklet for a list of weak acids and bases is being developed.

Usage:

"1 Square brackets are used in chemistry in a range of contexts eg concentrations (topic 1.3), Lewis (electron dot) structures (topic 4.3) and complexes (topic 14.1)

3 A number of acids and bases are used in our everyday life from rust removers to oven cleaners, from foods to toothpastes, from treatments for bee stings to treatment of wasp stings

1 Obtaining evidence for scientific theories—isotopic labelling and its use in defining equilibrium. (1.8) Common language across different disciplines—the term dynamic equilibrium is used in other contexts, but not necessarily with the chemistry definition in mind.

2 Falsification of theories—HCN altering the theory that oxygen was the element which gave a compound its acidic properties allowed for other acid–base theories to develop. (2.5) Theories being superseded—one early theory of acidity derived from the sensation of a sour taste, but this had been proven false. (1.9) Public understanding of science—outside of the arena of chemistry, decisions are sometimes referred to as ""acid test"" or ""litmus test"". (5.5)

3 Obtaining evidence for theories—observable properties of acids and bases have led to the modification of acid–base theories. (1.9)

4 Occam’s razor—the pH scale is an attempt to scale the relative acidity over a wide range of H+ concentrations into a very simple number. (2.7)

5 Improved instrumentation—the use of advanced analytical techniques has allowed the relative strength of different acids and bases to be quantified. (1.8) Looking for trends and discrepancies—patterns and anomalies in relative strengths of acids and bases can be explained at the molecular level. (3.1) The outcomes of experiments or models may be used as further evidence for a claim—data for a particular type of reaction supports the idea that weak acids exist in equilibrium. (1.9)

6 Risks and problems—oxides of metals and non-metals can be characterized by their acid–base properties. Acid deposition is a topic that can be discussed from different perspectives. Chemistry allows us to understand and to reduce the environmental impact of human activities.

What is equilibrium and what factors affect the position of equilibrium in a reversible reaction?

How are acid-base reactions defined and how do they involve proton transfer?

How is the acidity of a solution characterized and what factors affect the strength of acids and bases?

How does industrialization affect acid rain and what measures can be taken to reduce its impact?

How is equilibrium constant related to acid-base reactions and how can it be used to predict the feasibility of a reaction?

How do pH curves reflect the dissociation constants of acids and bases and how can they be used in experiments?

The curriculums retain the core concepts of chemical equilibrium, acid-base theory, and the pH scale. Revised curriculum adds more emphasis on the practical applications of these concepts in industrial and biological settings, as well as the use of modern technology and data analysis. Some topics, such as solubility product calculations, have been removed from FSC 2006 as they are covered in more detail in higher-level courses. The changes reflect the need to update the curriculum to reflect advances in the field and to prepare students for the challenges and opportunities of the future.

 

Electrochemistry and Redox
(Physical Chemistry)

- Give the characteristics of a Redox reaction.

- Determine the oxidation number of an atom of any element in a pure substance.

- Define oxidation and reduction in terms of a change in oxidation number.

- Use the oxidation-number change method to identify atoms being oxidized or reduced in redox reactions.

- Use the oxidation-number change method to balance redox equations.

- Balance redox reactions that take place in acid solutions.

- Break a redox reaction into oxidation and reduction half reactions.

- When given an unbalanced redox equation, use the half reaction method to balance the equation.

- Define cathode, anode, electrode potential and S.H.E. (Standard Hydrogen Electrode).

- Identify the substance oxidized and the substance reduced in a dry cell.

- Use the activity series of metals to predict the products of single replacement reactions. (Analysis)

- Define cell potential, and describe how it is determined.

- Describe the reaction that occurs when a lead storage battery is recharged.

- Explain how a fuel cell produces electrical energy.

- Define the standard electrode potential of an electrode.

- Distinguish between electrical terms such as coulomb, ampere, and volt.

- State and explain Faraday's laws.

- Describe how a dry cell supplies electricity.

- Explain how a lead storage battery produces electricity.

- Define corrosion and describe simple methods like electroplating and galvanizing for its prevention.

- Use standard electrode potentials to calculate the standard emf of cell

- Predict the feasibility of an electrochemical reaction from emf data.

- Calculate the amount of substance reduced when a quantity of another substance is oxidized in an electrochemical cell.

- Calculate the cell potential for an electrochemical cell under standard conditions.

- Deduce the direction of flow of electrons in an electrochemical cell.

- Calculate the quantity of charge passed in an electrochemical cell during electrolysis.

- Calculate the mass or volume of substance liberated during electrolysis.

6 Electrochemistry

6.1 Redox processes: electron transfer and changes in oxidation number (oxidation state)

1 calculate oxidation numbers of elements in compounds and ions

2 use changes in oxidation numbers to help balance chemical equations

3 explain and use the terms redox, oxidation, reduction and disproportionation in terms of electron transfer

and changes in oxidation number

4 explain and use the terms oxidising agent and reducing agent

5 use a Roman numeral to indicate the magnitude of the oxidation number of an element

24.1 Electrolysis

1 predict the identities of substances liberated during electrolysis from the state of electrolyte (molten or

aqueous), position in the redox series (electrode potential) and concentration

2 state and apply the relationship F = Le between the Faraday constant, F, the Avogadro constant, L, and the

charge on the electron, e

3 calculate:

(a) the quantity of charge passed during electrolysis, using Q = It

(b) the mass and/or volume of substance liberated during electrolysis

4 describe the determination of a value of the Avogadro constant by an electrolytic method

24.2 Standard electrode potentials E

; standard cell potentials E

and the Nernst equation

1 define the terms:

(a) standard electrode (reduction) potential

(b) standard cell potential

2 describe the standard hydrogen electrode

3 describe methods used to measure the standard electrode potentials of:

(a) metals or non-metals in contact with their ions in aqueous solution

(b) ions of the same element in different oxidation states

4 calculate a standard cell potential by combining two standard electrode potentials

5 use standard cell potentials to:

cell

(a) deduce the polarity of each electrode and hence explain/deduce the direction of electron flow in the

external circuit of a simple cell

(b) predict the feasibility of a reaction

6 deduce from E

reducing agents

values the relative reactivity of elements, compounds and ions as oxidising agents or as

7 construct redox equations using the relevant half-equations

8 predict qualitatively how the value of an electrode potential, E, varies with the concentrations of the

aqueous ions

9 use the Nernst equation, to predict quantitatively how the value of an electrode potential varies with the concentrations of the

aqueous ions;

10 understand and use the equation for gibbs free energy

Oxidation and reduction:

Oxidation and reduction can be considered in terms of oxygen gain/hydrogen loss, electron transfer or change in oxidation number.

- An oxidizing agent is reduced and a reducing agent is oxidized.

- Variable oxidation numbers exist for transition metals and for most main-group non-metals.

- The activity series ranks metals according to the ease with which they undergo oxidation.

- The Winkler Method can be used to measure biochemical oxygen demand (BOD), used as a measure of the degree of pollution in a water sample.

Electrochemical cells:

Voltaic (Galvanic) cells:

- Voltaic cells convert energy from spontaneous, exothermic chemical processes to electrical energy.

- Oxidation occurs at the anode (negative electrode) and reduction occurs at the cathode (positive electrode) in a voltaic cell. Electrolytic cells:

- Electrolytic cells convert electrical energy to chemical energy, by bringing about non-spontaneous processes.

- Oxidation occurs at the anode (positive electrode) and reduction occurs at the cathode (negative electrode) in an electrolytic cell

1. Understand and use the concept of oxidation numbers in identifying oxidation and reduction reactions

2. Use changes in oxidation numbers to balance chemical equations

3. Understand the terms redox, oxidation, reduction, and disproportionation in terms of electron transfer and changes in oxidation number

4. Understand the concepts of oxidizing and reducing agents, and their role in redox reactions

5. Use Roman numerals to indicate the magnitude of the oxidation number of an element

6. Understand the concept of the activity series of metals and how it relates to the ease of oxidation

7. Understand the use of the Winkler Method to measure biochemical oxygen demand (BOD) and its use as a measure of water pollution

8. Understand how voltaic (galvanic) cells convert energy from spontaneous, exothermic chemical processes to electrical energy, with oxidation at the anode and reduction at the cathode

9. Understand how electrolytic cells convert electrical energy to chemical energy, with oxidation at the anode and reduction at the cathode.

10. Students should be able to predict the identities of substances liberated during electrolysis based on the state of the electrolyte, position in the redox series, and concentration.

11. Students should understand and be able to apply the relationship between the Faraday constant, Avogadro constant, and the charge on the electron.

12. Students should be able to calculate the quantity of charge passed during electrolysis and the mass or volume of substance liberated during electrolysis.

13. Students should understand the determination of the Avogadro constant by an electrolytic method.

14. Students should be able to define and describe the terms "standard electrode potential" and "standard cell potential"

15. Students should be able to describe the standard hydrogen electrode and methods used to measure standard electrode potentials.

16. Students should be able to calculate standard cell potentials by combining two standard electrode potentials and use them to predict the feasibility of a reaction and the direction of electron flow in a simple cell.

17. Students should be able to deduce the relative reactivity of elements, compounds, and ions as oxidizing agents or reducing agents from their electrode potential values.

18. Students should be able to construct redox equations using relevant half-equations.

19. Students should understand how electrode potentials vary with the concentrations of aqueous ions and use the Nernst equation to predict this quantitatively.

20. Students should understand and use the equation for Gibbs free energy.

Oxidation number and oxidation state are often used interchangeably, though IUPAC does formally distinguish between the two terms. Oxidation numbers are represented by Roman numerals according to IUPAC.

- Oxidation states should be represented with the sign given before the number, eg +2 not 2+.

- The oxidation state of hydrogen in metal hydrides (-1) and oxygen in peroxides (-1) should be covered.

- A simple activity series is given in the data booklet in section 25. For voltaic cells, a cell diagram convention should be covered.

Usage:

-What are redox reactions and how do they play a role in chemical and biochemical processes?

-What is the difference between a voltaic cell and an electrolytic cell and how do they convert energy?

-What is the relationship between electrical and chemical energy in electrochemical cells?

-How do changes in oxidation state relate to redox reactions?

-What is the purpose of electrolysis and how can the products be predicted based on the electrolyte and position in the redox series?

-What is the relationship between Faraday's constant, Avogadro's constant, and the charge on the electron and how is it used in electrolysis calculations?

-What is the Nernst equation and how is it used to predict electrode potentials?

-What factors affect the feasibility of a reaction in an electrochemical cell?

-How can the relative reactivity of elements, compounds, and ions be determined as oxidizing or reducing agents using electrode potentials?

-What is the relationship between Gibbs free energy and electrode potentials in electrochemistry?

The NCC 2023 curriculum builds upon the previous curriculum by further emphasizing the concepts of oxidation numbers, redox reactions, and electrochemistry. It introduces new concepts such as the Winkler Method, Faraday constant, Avogadro constant, and Gibbs free energy. Additionally, it places greater emphasis on understanding the practical applications of these concepts, such as the use of the Winkler Method to measure water pollution and the construction of voltaic and electrolytic cells for energy conversion.

 

Periodicity
(Inorganic Chemistry)

- Recognize the demarcation of the Periodic Table into s block, p block, d block, and f block.

- Describe how physical properties like atomic radius, ionization energy, electronegativity, electrical conductivity and melting and boiling points of elements change within a group and within a period in the Periodic Table.

- Describe reactions of period 3 elements with water, oxygen and chlorine.

- Describe physical properties and acid-base behavior of oxides, chlorides and hydroxides of period 3 elements.

- Describe reactions of oxides and chlorides of period 3 elements with water.

- Describe reactions of Group IV elements with water.

- Discuss the chlorides and oxides of group IV elements.

- Explain the relative behaviour of halogens as oxidizing agents and reducing agents.

- Compare the acidity of hydrogen halides.

- Distinguish between an oxide and a peroxide.

- Write representative equations for the formation of oxides and sulphides.

- Compare the outermost s and p orbital system of an element with its chemical properties.

9 The Periodic Table: chemical periodicity

9.1 Periodicity of physical properties of the elements in Period 3

1 describe qualitatively (and indicate the periodicity in) the variations in atomic radius, ionic radius, melting

point and electrical conductivity of the elements

2 explain the variation in melting point and electrical conductivity in terms of the structure and bonding of the

elements

9.2 Periodicity of chemical properties of the elements in Period 3

1 describe, and write equations for, the reactions of the elements with oxygen, chlorine and water (Na and Mg only)

2 state and explain the variation in the oxidation number of the oxides and chlorides (NaCl, MgCl in terms of their outer shell (valence shell)electrons

3 describe, and write equations for, the reactions, if any, of the oxides with water including the likely pHs of the solutions obtained

4 describe, explain, and write equations for, the acid / base behaviour of the oxides and the hydroxides NaOH, Mg(OH) including, where relevant, amphoteric behaviour in reactions with acids and bases (sodium hydroxide only)

5 describe, explain, and write equations for, the reactions of the chlorides with water including the likely pHs of the solutions obtained

6 explain the variations and trends in 9.2.2, 9.2.3, 9.2.4 and 9.2.5 in terms of bonding and electronegativity

7 suggest the types of chemical bonding present in the chlorides and oxides from observations of their chemical and physical properties

9.3 Chemical periodicity of other elements

1 predict the characteristic properties of an element in a given group by using knowledge of chemical periodicity

2 deduce the nature, possible position in the Periodic Table and identity of unknown elements from given information about physical and chemical properties

3.1 Periodic table

Understandings:

• The periodic table is arranged into four blocks associated with the four sublevels—s, p, d, and f.

• The periodic table consists of groups (vertical columns) and periods (horizontal

rows).

• The period number (n) is the outer energy level that is occupied by electrons.

• The number of the principal energy level and the number of the valence

electrons in an atom can be deduced from its position on the periodic table.

• The periodic table shows the positions of metals, non-metals and metalloids.

Applications and skills:

• Deduction of the electron configuration of an atom from the element’s position

on the periodic table, and vice versa.

• The terms alkali metals, halogens, noble gases, transition metals, lanthanoids

and actinoids should be known.

• The group numbering scheme from group 1 to group 18, as recommended by

IUPAC, should be used.

3.2 Periodic trends

Understandings:

• Vertical and horizontal trends in the periodic table exist for atomic radius, ionic

radius, ionization energy, electron affinity and electronegativity.

• Trends in metallic and non-metallic behaviour are due to the trends above.

• Oxides change from basic through amphoteric to acidic across a period.

Applications and skills:

• Prediction and explanation of the metallic and non-metallic behaviour of an

element based on its position in the periodic table.

• Discussion of the similarities and differences in the properties of elements in

the same group, with reference to alkali metals (group 1) and halogens (group

17).

• Construction of equations to explain the pH changes for reactions of, and the oxides of nitrogen and sulfur with water.

• Only examples of general trends across periods and down groups are required.

For ionization energy the discontinuities in the increase across a period should

be covered.

Guidance:

• Group trends should include the treatment of the reactions of alkali metals with

water, alkali metals with halogens and halogens with halide ions.

1. The periodic table consists of groups (vertical columns) and periods (horizontal rows)

2. The periodic table is arranged into four blocks associated with the four sublevels—s, p, d, and f.

3. The period number (n) is the outer energy level that is occupied by electrons.

4. The number of the principal energy level and the number of the valence electrons in an atom can be deduced from its position on the periodic table.

5. The periodic table shows the positions of metals, non-metals and metalloids.

6. Vertical and horizontal trends in the periodic table exist for atomic radius, ionic radius, ionization energy, electron affinity and electronegativity.

7. Trends in metallic and non-metallic behavior are due to the trends in valence electrons.

8. The terms alkali metals, halogens, noble gases, transition metals, lanthanoids and actinoids should be known.

9. The group numbering scheme from group 1 to group 18, as recommended by IUPAC, should be used.

10. Deduction of the electron configuration of an atom from the element’s position on the periodic table, and vice versa.

11. describe, and write equations for, the reactions of the elements with oxygen, chlorine and water (Na and Mg only)

12. state and explain the variation in the oxidation number of the oxides and chlorides (NaCl, MgCl in terms of their outer shell (valence shell)electrons

13. describe, and write equations for, the reactions, if any, of the oxides with water including the likely pHs of the solutions obtained

14. describe, explain, and write equations for, the acid / base behaviour of the oxides and the hydroxides NaOH, Mg(OH)2 including, where relevant, amphoteric behaviour in reactions with acids and bases (sodium hydroxide only)

15. describe, explain, and write equations for, the reactions of the chlorides with water including the likely pHs of the solutions obtained

16. explain the variations and trends in terms of bonding and electronegativity

17. suggest the types of chemical bonding present in the chlorides and oxides from observations of their chemical and physical properties

18. predict the characteristic properties of an element in a given group by using knowledge of chemical periodicity

19. deduce the nature, possible position in the Periodic Table and identity of unknown elements from given information about physical and chemical properties

• The terms alkali metals, halogens, noble gases, transition metals, lanthanoids

and actinoids should be known.

• The group numbering scheme from group 1 to group 18, as recommended by

IUPAC, should be used.

Usage:

- anomalies in first ionization energy values can be connected to stability

in electron configuration

- production of acid rain

-What is the periodic table and how is it arranged?

-How does the arrangement of elements in the periodic table help predict their electron configuration?

-What trends in physical and chemical properties do elements show across periods and down groups in the periodic table?

The curricula cover the Periodic Table and its classification of elements. NCC 2023 expands upon the topics covered previously by including the use of the periodic table to predict properties of elements and identify unknown elements based on their properties. Former curriculum includes specific reactions of period 3 elements with water, oxygen, and chlorine, which are not explicitly mentioned in the NCC 2023 curriculum.

• Other scientific subjects also use the periodic table to understand the structure

and reactivity of elements as it applies to their own disciplines.

Group 2
(Inorganic Chemistry)

- Explain the trends in physical properties and oxidation states in groups I, II, IV and VII of the Periodic Table.

- Describe reactions of Group Ielements with water, oxygen and chlorine.

- Explain effect of heat on nitrates, carbonates and hydrogen carbonates of Group I elements.

- Describe reactions of Group II elements with water, oxygen and nitrogen.

- Discuss the trend in solubility of the hydroxides, sulphates and carbonates of Group II elements.

- Discuss the trends in thermal stability of the nitrates and carbonates of Group II elements.

- Differentiate beryllium from other members of its group.

10 Group 2

10.1 Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their

compounds

1 describe, and write equations for, the reactions of the elements with oxygen, water and dilute hydrochloric

and sulfuric acids

2 describe, and write equations for, the reactions of the oxides, hydroxides and carbonates with water and

dilute hydrochloric and sulfuric acids

3 describe, and write equations for, the thermal decomposition of the nitrates and carbonates, to include the

trend in thermal stabilities

4 describe, and make predictions from, the trends in physical and chemical properties of the elements involved

in the reactions in 10.1.1 and the compounds involved in 10.1.2, 10.1.3 and 10.1.5

5 state the variation in the solubilities of the hydroxides and sulfates

27 Group 2

27.1 Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their

compounds

1 describe and explain qualitatively the trend in the thermal stability of the nitrates and carbonates including

the effect of ionic radius on the polarisation of the large anion

2 describe and explain qualitatively the variation in solubility and of enthalpy change of solution, ΔH

, of the

hydroxides and sulfates in terms of relative magnitudes of the enthalpy change of hydration and the lattice

energy

sol

 

1. Understand the properties and trends of Group 2 elements, including their electron configurations, reactivity, and common compounds such as oxides, hydroxides and carbonates

2. Describe the chemical reactivity of Group 2 elements, including their reactions with oxygen, water, and acids.

3. Explain the reactivity of Group 2 elements in terms of their electron configuration and valence electrons.

4. Describe the industrial and everyday uses of Group 2 compounds, including their role in medicine and agriculture.

5. Understand and use the term reactivity series and its application in predicting the outcome of chemical reactions.

6. Explain the extraction and purification process of Group 2 elements and their compounds.

7. Understand and use the term thermal decomposition and its application in the analysis of Group 2 compounds especially carbonates and nitrates.

8. Explain the trend in solubility of group 2 sulfates and hydroxides using terms enthalpy of hydration and enthalpy of solution

9. Compare and contrast the properties and reactivity of Group 2 elements with other groups in the periodic table.

10. Understand and use the term complex ion and its application in the formation of Group 2 compounds.

11. Understand and use the term basic oxide and its application in the formation of Group 2 compounds.

 

What are the characteristics of Group 2 compounds?

How do the properties of Group 2 oxides differ from one another?

How do Group 2 carbonates and nitrates differ in terms of their chemical properties?

How do Group 2 compounds react with acids?

What are the industrial uses of Group 2 compounds?

How do Group 2 compounds contribute to environmental issues such as air pollution?

How do the solubility of Group 2 compounds change with temperature and pressure?

What is the role of Group 2 compounds in biological systems?

How are Group 2 compounds used in medicine?

What are the safety precautions that should be taken when handling Group 2 compounds?

it appears that much of the content related to the properties and reactivity of Group 1 and Group 2 elements has been retained, but there are some additions and removals. In the 2023 curriculum, there is an increased emphasis on the application of concepts such as enthalpy and complex ions, while some more basic content, such as the description of specific reactions, has been removed. Additionally, the 2023 curriculum appears to place a greater emphasis on developing students' understanding of scientific concepts and principles, rather than just memorizing facts.

 

Group 17
(Inorganic Chemistry)

 

11 Group 17

11.1 Physical properties of the Group 17 elements

1 describe the colours and the trend in volatility of chlorine, bromine and iodine

2 describe and explain the trend in the bond strength of the halogen molecules

3 interpret the volatility of the elements in terms of instantaneous dipole–induced dipole forces

11.2 The chemical properties of the halogen elements and the hydrogen halides

1 describe the relative reactivity of the elements as oxidising agents

2 describe the reactions of the elements with hydrogen and explain their relative reactivity in these reactions

3 describe the relative thermal stabilities of the hydrogen halides and explain these in terms of bond strengths

11.3 Some reactions of the halide ions

1 describe the relative reactivity of halide ions as reducing agents

2 describe and explain the reactions of halide ions with:

(a) aqueous silver ions followed by aqueous ammonia (the formation and formula of thecomplex is not required)

(b) concentrated sulfuric acid, to include balanced chemical equations

11.4 The reactions of chlorine

1 describe and interpret, in terms of changes in oxidation number, the reaction of chlorine with cold and with hot aqueous sodium hydroxide and recognise these as disproportionation reactions

2 explain, including by use of an equation, the use of chlorine in water purification to include the production of the active species HOCl and ClO– which kill bacteria.

 

1. Describe the colors and trend in volatility of chlorine, bromine and iodine

2. Describe and explain the trend in bond strength of halogen molecules

3. Interpret the volatility of the elements in terms of instantaneous dipole-induced dipole forces

4. Describe the relative reactivity of the halogen elements as oxidizing agents

5. Describe the reactions of the elements with hydrogen and explain their relative reactivity in these reactions

6. Describe the relative thermal stabilities of the hydrogen halides and explain these in terms of bond strengths

7. Describe the relative reactivity of halide ions as reducing agents

8. Describe and explain the reactions of halide ions with aqueous silver ions and concentrated sulfuric acid

9. Describe and interpret the reaction of chlorine with cold and hot aqueous sodium hydroxide as disproportionation reactions

10. Explain the use of chlorine in water purification, including the production of the active species HOCl and ClO- which kill bacteria.

some other important reaction to include would be the reactions of halogens with metals to form halides, with alkenes to form haloalkanes, with alcohols to form haloalkanes, with carboxylic acids to form haloalkanoic acids, amines to form haloamines, and with sulfur and phosphorus compounds to form halides of these elements.

 

Group 17 elements were not separately discussed in 2006 curriculum and have been added here. Additional information on their reactivities, stability and usage of halogens as oxidizind and reducing specie for environmental clealiness will certainly impact profoundly.

 

N & S (Inorganic Chemistry)

 

12 Nitrogen and sulfur

12.1 Nitrogen and sulfur

1 explain the lack of reactivity of nitrogen, with reference to triple bond strength and lack of polarity

2 describe and explain:

(a) the basicity of ammonia, using the Brønsted–Lowry theory

(b) the structure of the ammonium ion and its formation by an acid–base reaction

(c) the displacement of ammonia from ammonium salts by an acid–base reaction

3 state and explain the natural and man-made occurrences of oxides of nitrogen and their catalytic removal

from the exhaust gases of internal combustion engines

4 understand that atmospheric oxides of nitrogen (NO and NO

) can react with unburned hydrocarbons to

form peroxyacetyl nitrate, PAN, which is a component of photochemical smog

5 describe the role of NO and NO

2

2

in the formation of acid rain both directly and in their catalytic role in the

oxidation of atmospheric sulfur dioxide

 

Nitrogen

1. Explain the lack of reactivity of nitrogen due to its triple bond strength and lack of polarity

2. Describe and explain the basicity of ammonia using the Brønsted–Lowry theory

3. Understand the structure of the ammonium ion and how it is formed by an acid-base reaction

4. Describe how ammonia can be displaced from ammonium salts through acid-base reactions

5. Understand the natural and man-made occurrences of oxides of nitrogen and their catalytic removal from exhaust gases of internal combustion engines

6. Explain the role of NO and NO2 in the formation of photochemical smog, specifically in the reaction with unburned hydrocarbons to form peroxyacetyl nitrate (PAN)

7. Describe the role of NO and NO2 in the formation of acid rain, both directly and through their catalytic role in the oxidation of atmospheric sulfur dioxide.

Sulfur

8. Explain the lack of reactivity of sulfur, with reference to its bonding and stability of its compounds.

9. Describe and explain the different oxidation states of sulfur and their relative stability.

10. Understand the properties and uses of sulfuric acid, including its production and industrial applications.

11. Describe the role of sulfur in the formation of acid rain and its impact on the environment.

12. Identify and describe the chemical reactions and processes involving sulfur, such as combustion and oxidation.

13. Understand the uses of sulfur compounds in industry and everyday life, such as in fertilizers, gunpowder and rubber, and in the Synthetic organic chemistry, including the synthesis of dyes, drugs and fragrances.

 

 

They were not in 2006 curriculum so it is added to explain their role in industry and and everyday lives.

 

Transition Metals
(Inorganic Chemistry)

- Describe electronic structures of elements and ions of d-block elements.

- Explain why the electronic configuration for chromium and copper differ from those assigned using the Aufbau principle.

- Describe important reactions and uses of Vanadium, Chromium, Manganese, Iron and Copper.

- Explain shapes, origin of colors and nomenclature of coordination compounds.

- Relate the coordination number of ions to the crystal structure of the compound of which they are a part.

- Define an alloy and describe some properties of an alloy that are different from the metals that compose it.

- Describe the reactions of potassium dichromate with oxalic acid and Mohr's salt.

- Describe the reactions of potassium manganate VII with ferrous sulphate, oxalic acid and Mohr's salt.

- Calculate concentration of iron (II) ions in solution by titration with KMn04-

- Explain the reaction of hexaaquacopper II) ions with iodide and determine the concentration of copper (II) ions in the solution.

28.1 General physical and chemical properties of the first row of transition elements, titanium to copper

1 define a transition element as a d-block element which forms one or more stable ions with incomplete d

orbitals

2 sketch the shape of a 3dxy orbital and 3dz2 orbital

3 understand that transition elements have the following properties:

(a) they have variable oxidation states

(b) they behave as catalysts

(c) they form complex ions

(d) they form coloured compounds

4 explain why transition elements have variable oxidation states in terms of the similarity in energy of the 3d

and the 4s sub-shells

5 explain why transition elements behave as catalysts in terms of having more than one stable oxidation state,

and vacant d orbitals that are energetically accessible and can form dative bonds with ligands

6 explain why transition elements form complex ions in terms of vacant d orbitals that are energetically

accessible

28.2 General characteristic chemical properties of the first set of transition elements, titanium to copper

1 describe and explain the reactions of transition elements with ligands to form complexes, including the

complexes of copper(II) and cobalt(II) ions with water and ammonia molecules and hydroxide and chloride

ions

2 define the term ligand as a species that contains a lone pair of electrons that forms a dative covalent bond to

a central metal atom / ion

3 understand and use the terms

(a) monodentate ligand including as examples H2O, NH3, Cl – and CN–

(b) bidentate ligand including as examples 1,2-diaminoethane, en, H2NCH2CH2NH2 and the ethanedioate

ion, C2O42–

(c) polydentate ligand including as an example EDTA4–

4 define the term complex as a molecule or ion formed by a central metal atom / ion surrounded by one or

more ligands

5 describe the geometry (shape and bond angles) of transition element complexes which are linear, square

planar, tetrahedral or octahedral

6 (a) state what is meant by coordination number

(b) predict the formula and charge of a complex ion, given the metal ion, its charge or oxidation state, the

ligand and its coordination number or geometry

7 explain qualitatively that ligand exchange can occur, including the complexes of copper(II) ions and

cobalt(II) ions with water and ammonia molecules and hydroxide and chloride ions

8 predict, using E

values, the feasibility of redox reactions involving transition elements and their ions

9 describe the reactions of, and perform calculations involving:

(a) MnO4

– / C2O4

2– in acid solution given suitable data

(b) MnO4

– / Fe2+ in acid solution given suitable data

(c) Cu2+ / I– given suitable data

10 perform calculations involving other redox systems given suitable data

28.3 Colour of complexes

1 define and use the terms degenerate and non-degenerate d orbitals

2 describe the splitting of degenerate d orbitals into two non-degenerate sets of d orbitals of higher energy,

and use of Δ E in:

(a) octahedral complexes, two higher and three lower d orbitals

(b) tetrahedral complexes, three higher and two lower d orbitals

3 explain why transition elements form coloured compounds in terms of the frequency of light absorbed as an

electron is promoted between two non-degenerate d orbitals

4 describe, in qualitative terms, the effects of different ligands on Δ E, frequency of light absorbed, and hence

the complementary colour that is observed

5 use the complexes of copper(II) ions and cobalt(II) ions with water and ammonia molecules and hydroxide

and chloride ions as examples of ligand exchange affecting the colour observed

28.4 Stereoisomerism in transition element complexes

1 describe the types of stereoisomerism shown by complexes, including those associated with bidentate

ligands:

(a) geometrical (cis-trans) isomerism, e.g. square planar such as [Pt(NH3)2Cl 2] and octahedral such as

[Co(NH3)4(H2O)2]2+ and [Ni(H2NCH2CH2NH2)2(H2O)2]2+

(b) optical isomerism, e.g. [Ni(H2NCH2CH2NH2)3]2+ and [Ni(H2NCH2CH2NH2)2(H2O)2]2+

2 deduce the overall polarity of complexes such as those described in 28.4.1(a) and 28.4.1(b)

28.5 Stability constants, Kstab

1 define the stability constant, Kstab, of a complex as the equilibrium constant for the formation of the

complex ion in a solvent (from its constituent ions or molecules)

2 write an expression for a Kstab of a complex ([H2O] should not be included)

3 use Kstab expressions to perform calculations

4 describe and explain ligand exchanges in terms of Kstab values and understand that a large Kstab is due to the

formation of a stable complex ion

1 First-row d-block elements

Understandings:

Transition elements have variable oxidation states, form complex ions with ligands, have coloured compounds, and display catalytic and magnetic properties.

- Zn is not considered to be a transition element as it does not form ions with incomplete d-orbitals.

- Transition elements show an oxidation state of +2 when the s-electrons are removed.

2 Coloured complexes

Understandings:

- The d sub-level splits into two sets of orbitals of different energy in a complex ion.

- Complexes of d-block elements are coloured, as light is absorbed when an electron is excited between the d-orbitals.

- The colour absorbed is complementary to the colour observed.

28.1 General physical and chemical properties of the first row of transition elements, titanium to copper

1 define a transition element as a d-block element which forms one or more stable ions with incomplete d

orbitals

2 sketch the shape of a 3dxy orbital and 3dz2 orbital

3 understand that transition elements have the following properties:

(a) they have variable oxidation states

(b) they behave as catalysts

(c) they form complex ions

(d) they form coloured compounds

4 explain why transition elements have variable oxidation states in terms of the similarity in energy of the 3d

and the 4s sub-shells

5 explain why transition elements behave as catalysts in terms of having more than one stable oxidation state,

and vacant d orbitals that are energetically accessible and can form dative bonds with ligands

6 explain why transition elements form complex ions in terms of vacant d orbitals that are energetically

accessible

28.2 General characteristic chemical properties of the first set of transition elements, titanium to copper

1 describe and explain the reactions of transition elements with ligands to form complexes, including the

complexes of copper(II) and cobalt(II) ions with water and ammonia molecules and hydroxide and chloride

ions

2 define the term ligand as a species that contains a lone pair of electrons that forms a dative covalent bond to

a central metal atom / ion

3 understand and use the terms

(a) monodentate ligand including as examples H2O, NH3, Cl – and CN–

(b) bidentate ligand including as examples 1,2-diaminoethane, en, H2NCH2CH2NH2 and the ethanedioate

ion, C2O42–

(c) polydentate ligand including as an example EDTA4–

4 define the term complex as a molecule or ion formed by a central metal atom / ion surrounded by one or

more ligands

5 describe the geometry (shape and bond angles) of transition element complexes which are linear, square

planar, tetrahedral or octahedral

6 (a) state what is meant by coordination number

(b) predict the formula and charge of a complex ion, given the metal ion, its charge or oxidation state, the

ligand and its coordination number or geometry

7 explain qualitatively that ligand exchange can occur, including the complexes of copper(II) ions and

cobalt(II) ions with water and ammonia molecules and hydroxide and chloride ions

8 predict, using E

values, the feasibility of redox reactions involving transition elements and their ions

9 describe the reactions of, and perform calculations involving:

(a) MnO4

– / C2O4

2– in acid solution given suitable data

(b) MnO4

– / Fe2+ in acid solution given suitable data

(c) Cu2+ / I– given suitable data

10 perform calculations involving other redox systems given suitable data

28.3 Colour of complexes

1 define and use the terms degenerate and non-degenerate d orbitals

2 describe the splitting of degenerate d orbitals into two non-degenerate sets of d orbitals of higher energy,

and use of Δ E in:

(a) octahedral complexes, two higher and three lower d orbitals

(b) tetrahedral complexes, three higher and two lower d orbitals

3 explain why transition elements form coloured compounds in terms of the frequency of light absorbed as an

electron is promoted between two non-degenerate d orbitals

4 describe, in qualitative terms, the effects of different ligands on Δ E, frequency of light absorbed, and hence

the complementary colour that is observed

5 use the complexes of copper(II) ions and cobalt(II) ions with water and ammonia molecules and hydroxide

and chloride ions as examples of ligand exchange affecting the colour observed

28.4 Stereoisomerism in transition element complexes

1 describe the types of stereoisomerism shown by complexes, including those associated with bidentate

ligands:

(a) geometrical (cis-trans) isomerism, e.g. square planar such as [Pt(NH3)2Cl 2] and octahedral such as

[Co(NH3)4(H2O)2]2+ and [Ni(H2NCH2CH2NH2)2(H2O)2]2+

(b) optical isomerism, e.g. [Ni(H2NCH2CH2NH2)3]2+ and [Ni(H2NCH2CH2NH2)2(H2O)2]2+

2 deduce the overall polarity of complexes such as those described in 28.4.1(a) and 28.4.1(b)

28.5 Stability constants, Kstab

1 define the stability constant, Kstab, of a complex as the equilibrium constant for the formation of the

complex ion in a solvent (from its constituent ions or molecules)

2 write an expression for a Kstab of a complex ([H2O] should not be included)

3 use Kstab expressions to perform calculations

4 describe and explain ligand exchanges in terms of Kstab values and understand that a large Kstab is due to the

formation of a stable complex ion

Common oxidation charges on transition metal ions are given in section 9 of the data booklet, and common oxidation states are given in section 14. The spectrochemical series is given in the data booklet in section 15. A list of polydentate ligands is given in the data booklet in section 16.

- Students are not expected to recall the colour of specific complex ions.The relation between the colour observed and absorbed is illustrated by the colour wheel in the data booklet in section 17. - Students are not expected to know the different splitting patterns and their relation to the coordination number. Only the splitting of the 3d orbitals in an octahedral crystal field is required.

The transition elements have characteristic properties; these properties are related to their all having incomplete d sublevels.

d-orbitals have the same energy in an isolated atom, but split into two sub-levels in a complex ion. The electric field of ligands may cause the d-orbitals in

complex ions to split so that the energy of an electron transition between them corresponds to a photon of visible light.

The present curricula covers the general physical and chemical properties, characteristic chemical properties, colour of complexes, stereoisomerism, and stability constants of the first row of transition elements, titanium to copper. It includes concepts such as transition elements having variable oxidation states, forming complex ions and coloured compounds, behaving as catalysts due to the presence of vacant d orbitals, and exhibiting stereoisomerism with bidentate ligands. also covers the use of Kstab expressions to perform calculations and understand ligand exchanges.

 

Air
(Envioronmental Chemistry)

- Recognize various chemical reactions occurring in the atmosphere.

- Recognize that the release of COx, SOx, NOx, VOCs are associated with the combustion of hydrocarbon based fuels.

- Outline problems associated with release of pollutants including acid rain and the

formation by free radical reactions of hazardous inorganic and organic compounds e.g., PAN.

- Describe causes and impacts of urban smog.

- Explain greenhouse effect and global warming as resulting in climate change.

- Explain the build up to and recognize the adverse effects of ozone in the troposphere.

- Describe the role of CFCs in destroying ozone in the stratosphere.

- Describe the role of ozone in the stratosphere in reducing the intensity of harmful UV

radiation reaching the earth.

- List possible alternatives to the use of CFCs.

- Recognize and describe various water pollutants.

- Explain the various parameters of water analysis.

- List some major products of the petrochemicals industry, together with their uses.

- Describe how properties of gases promote greenhouse effect.

- Make connections between Halons and CFCs and their effects on ozone depletion.

- Predict effects of radiation pollution.

- Explain the need to work in a well-ventilated area when working with toxic solvents as used in adhesives.

 

 

1. Understanding of the properties and composition of air and the factors that affect air quality

2. Knowledge of the sources and effects of air pollution, including both natural and human-caused pollutants including Carbon monoxide (CO), Sulfur dioxide (SO2),Nitrogen oxides (NOx), Particulate matter (PM), Ozone (O3), Lead (Pb), Mercury (Hg), Polycyclic aromatic hydrocarbons (PAHs), Persistent organic pollutants (POPs), Greenhouse gases (such as carbon dioxide, methane, and nitrous oxide), Chlorofluorocarbons (CFCs) and other ozone-depleting substances, Volatile organic compounds (VOCs), Heavy metals (such as lead, mercury, and cadmium))

3. Familiarity with the methods and techniques used to measure and monitor air quality

4. Understanding of the impact of human activities on the atmosphere, including the effects of burning fossil fuels and deforestation

5. Knowledge of the chemical reactions and processes that occur in the atmosphere, such as the formation of smog and acid rain

6. Familiarity with the laws and regulations related to air quality and the measures used to control air pollution

7. Ability to analyze data and interpret air quality measurements and trends

8. Understanding of the link between air quality and human health and the ability to evaluate the potential health risks associated with air pollution

9. Knowledge of the technologies and strategies used to reduce air pollution and improve air quality, such as emissions control and renewable energy sources.

10. Ability to design experiments and collect data to test hypotheses about air quality

11. Familiarity with the global scale problems of air pollution, such as global warming and the greenhouse effect.

12. Ability to think critically about the economic, social and political issues related to air pollution and air quality management.

13. Familiarity with light pollution, microplastics, noise pollution, toxic waste and plastic pollution.

It would be pertinent to include

-The role of government and industry in protecting air and the environment

-Case studies on specific air pollution events and their causes and effects

-Field studies and investigations of local air and environmental conditions

-Discussion of current policy and legislation related to air and the environment

-Critical evaluation of scientific information and media coverage of air and environmental issues.

-The composition and behavior of the atmosphere, including the impacts of human activity.

-The greenhouse effect and global warming

-Air pollution, sources and effects

-Ozone depletion and the Antarctic ozone hole

-Acid rain and its causes and effects

-Air quality and regulations

-Alternative energy sources and their impact on air and the environment

-Climate change and its causes and effects

-Air and water pollution control technology

-The role of government and industry in protecting air and the environment

it looks that the 2023 curriculum has a broader and more comprehensive focus on air quality and pollution. The 2006 curriculum primarily focuses on atmospheric chemistry and the effects of pollution on the environment and human health. In contrast, the updated curriculum covers a wider range of pollutants, including microplastics and plastic pollution, as well as light and noise pollution. Additionally, the 2023 curriculum places greater emphasis on the impact of human activities on the atmosphere and strategies for reducing air pollution. Overall, the 2023 curriculum appears to be more up-to-date and relevant to current environmental issues.

 

Water
(Envioronmental Chemistry)

- Estimate chloride ions in tap water using titration technique.

- Describe how rain water seepage through hazardous wastes dumpsites can dissolve and reach drinking water supplies.

- Describe three ways in which water is purified naturally.

- Explain how photochemical reactions contribute to air pollution.

- Identify ways in which air pollution resulting from auto exhausts can be alleviated.

- Recognize the use of catalytic converters in reducing pollutant emissions from petrol driven cars.

- Differentiate between ozone at the earth's surface and ozone formation and depletion in the atmosphere.

- Realize that dumping waste water from household and industry without treatment to the rivers and creeks is dangerous for the environment.

 

 

1. Understanding of different types of water pollution, such as point source and nonpoint source pollution

2. Familiarity with common water pollutants, such as oil, pesticides, and heavy metals

3. Knowledge of the sources and effects of water pollution on human health and the environment

4. Understanding of water treatment methods and technologies, such as filtration and purification

5. Familiarity with laws and regulations related to water pollution and conservation

6. Understanding of the impact of human activities on water resources, such as agriculture and industrial processes

7. Knowledge of conservation and management strategies for protecting and preserving water resources

8. Understanding of the chemical properties of water and how they relate to water quality and pollution.

Understanding of sources of pollution in Pakistan, such as industrial waste, agricultural practices, urbanization and transportation

Knowledge of the effects of pollution on human health, including respiratory issues, cardiovascular disease and cancer

Understanding of the impacts of pollution on the environment, including air and water pollution, deforestation and loss of biodiversity

Familiarity with common pollutants found in Pakistan, such as lead, sulfur dioxide, and nitrogen oxides

Knowledge of government policies and regulations in place to address pollution in Pakistan

Understanding of the importance of community engagement and individual actions in addressing and reducing pollution

Familiarity with technologies and practices used in pollution control and treatment, such as filtration, neutralization and bioremediation

Ability to analyze data and evaluate the effectiveness of pollution control measures

Familiarity with international conventions and agreements related to pollution, such as the Paris Agreement, and how they relate to Pakistan.

Intermediates of aerobic respiration and photosynthesis are not required. The names and structural formulas of the amino acids are given.

- Reference should be made to alpha helix and beta pleated sheet, and to fibrous and globular proteins with examples of each.

- In paper chromatography the use of Rf values and locating agents should be covered.

- In enzyme kinetics Km and Vmax are not required

 

The NCC 2023 curriculum retains the core concepts of water pollution and its impact on human health and the environment, along with water treatment methods and laws related to water pollution and conservation. The curriculum adds practical skills to estimate chloride ions in tap water, knowledge of natural water purification methods, and a deeper understanding of air pollution and its mitigation strategies. The curriculum also emphasizes the chemical properties of water and its relationship with water quality and pollution. Overall, the curriculum covers a wider range of topics related to water pollution and conservation, including the impact of human activities on water resources and conservation and management strategies.

 

Green chemistry and Sustainability
(Envioronmental Chemistry)

 

 

 

The goal of this section is to introduce the concepts of green chemistry and sustainability, and to develop a sense of relation and responsibility in individuals towards the world and envioronment.

Candidates are expected to

1. Understand the principles and practices of green chemistry, including reducing or eliminating the use and generation of hazardous substances in the design, manufacture, and use of chemical products.

2. Understand the concept of sustainability and its relationship to green chemistry.

3. Understand the environmental and human health impacts of traditional chemical processes and products.

4. Understand the potential benefits of using green chemistry in chemical manufacturing, including reduced waste and pollution, increased efficiency, and cost savings.

5. Understand the role of government, industry, and individuals in promoting and implementing green chemistry and sustainability practices.

6. Understand the use of renewable resources and the reduction of waste and carbon footprint.

7. Understand the concept of life-cycle assessment and its application in evaluating the environmental impact of chemical products and processes.

8. Understand the importance of collaboration and interdisciplinary approaches in promoting and implementing green chemistry and sustainability practices.

9. Understand the role of chemists in the development and implementation of green chemistry and sustainability practices.

10. Understand the importance of consumer awareness and education in promoting green chemistry and sustainability.

11. Understand the impact of green chemistry and sustainability on economic, environmental and social aspects.

Care should be taken to not reduce this section ot a rote learning or memorization exercise. Teachers may do any of the following and teach this through lens of reasoning and critical thinking

Incorporate real-world examples and case studies of sustainable practices in industry and everyday life to make the subject more relatable and engaging for students.

Use hands-on activities and laboratory experiments that demonstrate the principles of green chemistry and sustainability in action.

Encourage critical thinking and problem-solving by having students analyze and evaluate the environmental impact of various chemical processes and products.

Provide opportunities for students to conduct independent research and explore current developments and advances in the field of green chemistry and sustainability.

Foster a culture of environmental responsibility and stewardship in the classroom by promoting sustainable practices such as recycling, energy conservation, and reducing waste.

Use technology such as virtual simulations to visualize the impact of industrial practices on the environment.

Collaborate with other teachers in the school and incorporate green chemistry and sustainability topics into interdisciplinary curriculum.

Invite experts from industry or research institutions to speak to the class or participate in webinars on these topics.

Use online resources and interactive simulations to supplement classroom instruction.

Encourage students to think about the long-term implications of their actions and how to make responsible choices as consumers and citizens.

A few examples:

Real world case studies:

The use of biofuels as a sustainable alternative to fossil fuels

The implementation of green roofs and urban gardening to reduce the urban heat island effect

The use of recycled materials in product manufacturing and packaging

The development of closed loop systems in industry to minimize waste and conserve resources

Webinars:

"Green Chemistry: Innovations in Sustainable Science" by the American Chemical Society

"Sustainability in Action: Implementing Green Chemistry in Industry" by the Royal Society of Chemistry

"Sustainable Chemistry: From Research to Practice" by the European Chemical Society

Sustainability in action:

Visiting a local recycling facility or waste-to-energy plant to learn about waste management practices

Inviting a sustainability expert or green chemist to speak to the class about their work and experiences

Partnering with a local organization to implement a sustainability project in the community, such as a community garden or clean-up initiative.

Organizing a field trip to a sustainable business or organization to see green chemistry and sustainability practices in action.

How can the principles of green chemistry be applied to improve the sustainability of industrial processes and products?

What are some specific examples of companies or industries that have successfully implemented sustainable practices through the use of green chemistry?

How can individuals and consumers play a role in promoting and supporting sustainable chemistry practices in their daily lives?

This topic was not previously discussed in 2006 curriculum but is added with the objective to foster and inculcate scientific literacy in sustainability and to simultaneously develop such skills among the present and future generations

 

Introduction to Organic Chemistry
(Nomenclature, Functional Group, Isomerism, Formulae)
(Organic Chemistry)

- Define organic chemistry and organic compounds.

- Explain why there is such a diversity and magnitude of organic compounds.

- Classify organic compounds on structural basis.

- Explain the use of coal as a source of both aliphatic and aromatic hydrocarbons.

- Explain the use of plants as a source of organic compounds.

- Explain that organic compounds are also synthesized in the lab.

- Define functional groups and homologous series.

13.1 Formulae, functional groups and the naming of organic compounds

1 define the term hydrocarbon as a compound made up of C and H atoms only

2 understand that alkanes are simple hydrocarbons with no functional group

3 understand that the compounds in the table on page 26 and 27 contain a functional group which dictates

their physical and chemical properties

4 interpret and use the general, structural, displayed and skeletal formulae of the classes of compound stated

in the table on page 26 and 27

5 understand and use systematic nomenclature of simple aliphatic organic molecules with functional groups

detailed in the table on page 26 and 27, up to six carbon atoms (six plus six for esters, straight chains only

for esters and nitriles)

6 deduce the molecular and/or empirical formula of a compound, given its structural, displayed or skeletal

formula

13.2 Characteristic organic reactions

1 interpret and use the following terminology associated with types of organic compounds and reactions:

(a) homologous series

(b) saturated and unsaturated

(c) homolytic and heterolytic fission

(d) free radical, initiation, propagation, termination (the use of arrows to show movement of single

electrons is not required)

(e) nucleophile, electrophile, nucleophilic, electrophilic

(f) addition, substitution, elimination, hydrolysis, condensation

(g) oxidation and reduction

(in equations for organic redox reactions, the symbol [O] can be used to represent one atom of oxygen from

an oxidising agent and the symbol [H] one atom of hydrogen from a reducing agent)

2 understand and use the following terminology associated with types of organic mechanisms:

(a) free-radical substitution

(b) electrophilic addition

(c) nucleophilic substitution

(d) nucleophilic addition

(in organic reaction mechanisms, the use of curly arrows to represent movement of electron pairs is

expected; the arrow should begin at a bond or a lone pair of electrons)

13.3 Shapes of organic molecules; σ and π bonds

1 describe organic molecules as either straight-chained, branched or cyclic

2 describe and explain the shape of, and bond angles in, molecules containing sp, sp

2

and sp

hybridised atoms

3 describe the arrangement of σ and π bonds in molecules containing sp, sp

2

and sp

3

hybridised atoms

4 understand and use the term planar when describing the arrangement of atoms in organic molecules, for

example ethene

3

13.4 Isomerism: structural and stereoisomerism

1 describe structural isomerism and its division into chain, positional and functional group isomerism

2 describe stereoisomerism and its division into geometrical (cis/trans) and optical isomerism (use of E, Z

nomenclature is acceptable but is not required)

3 describe geometrical (cis/trans) isomerism in alkenes, and explain its origin in terms of restricted rotation

due to the presence of π bonds

4 explain what is meant by a chiral centre and that such a centre gives rise to two optical isomers

(enantiomers)

(Students should appreciate that compounds can contain more than one chiral centre, but knowledge of

meso compounds, or nomenclature such as diastereoisomers is not required)

5 identify chiral centres and geometrical (cis/trans) isomerism in a molecule of given structural formula

including cyclic compounds

6 deduce the possible isomers for an organic molecule of known molecular formula

29.1 Formulae, functional groups and the naming of organic compounds

1 understand that the compounds in the table on page 42 contain a functional group which dictates their

physical and chemical properties

2 interpret and use the general, structural, displayed and skeletal formulae of the classes of compound stated

in the table on page 42

3 understand and use systematic nomenclature of simple aliphatic organic molecules (including cyclic

compounds containing a single ring of up to six carbon atoms) with functional groups detailed in the table

on page 42, up to six carbon atoms (six plus six for esters and amides, straight chains only for esters and

nitriles)

4 understand and use systematic nomenclature of simple aromatic molecules with one benzene ring and one

or more simple substituents, for example 3-nitrobenzoic acid or 2,4,6-tribromophenol

29.2 Characteristic organic reactions

1 understand and use the following terminology associated with types of organic mechanisms:

(a) electrophilic substitution

(b) addition-elimination

29.3 Shapes of aromatic organic molecules; σ and π bonds

1 describe and explain the shape of benzene and other aromatic molecules, including sp

hybridisation, in

terms of σ bonds and a delocalised π system

2

29.4 Isomerism: optical

1 understand that enantiomers have identical physical and chemical properties apart from their ability to

rotate plane polarised light and their potential biological activity

2 understand and use the terms optically active and racemic mixture

3 describe the effect on plane polarised light of the two optical isomers of a single substance

4 explain the relevance of chirality to the synthetic preparation of drug molecules including:

(a) the potential different biological activity of the two enantiomers

(b) the need to separate a racemic mixture into two pure enantiomers

(c) the use of chiral catalysts to produce a single pure optical isomer

(Students should appreciate that compounds can contain more than one chiral centre, but knowledge of

meso compounds and nomenclature such as diastereoisomers is not required.)

A homologous series is a series of compounds of the same family, with the same general formula, which differ from each other by a common structural unit.

- Structural formulas can be represented in full and condensed format.

- Structural isomers are compounds with the same molecular formula but different arrangements of atoms.

- Functional groups are the reactive parts of molecules.

- Saturated compounds contain single bonds only and unsaturated compounds contain double or triple bonds.

- Benzene is an aromatic, unsaturated hydrocarbon.

Alkanes:

- Alkanes have low reactivity and undergo free-radical substitution reactions. Alkenes:

- Alkenes are more reactive than alkanes and undergo addition reactions. Bromine water can be used to distinguish between alkenes and alkanes. Alcohols:

- Alcohols undergo esterification (or condensation) reactions with acids and some undergo oxidation reactions. Halogenoalkanes:

- Halogenoalkanes are more reactive than alkanes. They can undergo (nucleophilic) substitution reactions. A nucleophile is an electron-rich species containing a lone pair that it donates to an electron-deficient carbon. Polymers:

- Addition polymers consist of a wide range of monomers and form the basis of the plastics industry. Benzene:

- Benzene does not readily undergo addition reactions but does undergo electrophilic substitution reactions

Types of organic reactions:

Nucleophilic Substitution Reactions:

- SN1 represents a nucleophilic unimolecular substitution reaction and SN2 represents a nucleophilic bimolecular substitution reaction. SN1 involves a carbocation intermediate. SN2 involves a concerted reaction with a transition state.

- For tertiary halogenoalkanes the predominant mechanism is SN1 and for primary halogenoalkanes it is SN2. Both mechanisms occur for secondary halogenoalkanes.

- The rate determining step (slow step) in an SN1 reaction depends only on the concentration of the halogenoalkane, rate = k[halogenoalkane]. For SN2, rate = k[halogenoalkane][nucleophile]. SN2 is stereospecific with an inversion of configuration at the carbon.

- SN2 reactions are best conducted using aprotic, non-polar solvents and SN1 reactions are best conducted using protic, polar solvents. Electrophilic

Addition Reactions:

- An electrophile is an electron-deficient species that can accept electron pairs from a nucleophile. Electrophiles are Lewis acids. Markovnikov’s rule can be applied to predict the major product in electrophilic addition reactions of unsymmetrical alkenes with hydrogen halides and interhalogens. The formation of the major product can be explained in terms of the relative stability of possible carbocations in the reaction mechanism. Electrophilic

Substitution Reactions:

- Benzene is the simplest aromatic hydrocarbon compound (or arene) and has a delocalized structure of π bonds around its ring. Each carbon to carbon bond has a bond order of 1.5. Benzene is susceptible to attack by electrophiles. Reduction Reactions:

- Carboxylic acids can be reduced to primary alcohols (via the aldehyde). Ketones can be reduced to secondary alcohols. Typical reducing agents are lithium aluminium hydride (used to reduce carboxylic acids) and sodium borohydride.

1. Understand that hydrocarbons are compounds made up of C and H atoms only

2. Understand that alkanes are simple hydrocarbons with no functional group

3. Understand that compounds in a table contain a functional group which dictates their physical and chemical properties

4. Interpret and use the general, structural, displayed and skeletal formulae of the classes of compounds

5. Understand and use systematic nomenclature of simple aliphatic organic molecules with functional groups

6. Deduce the molecular and/or empirical formula of a compound, given its structural, displayed or skeletal formula

7. Understand and use terminology associated with types of organic compounds and reactions: homologous series, saturated and unsaturated, homolytic and heterolytic fission, free radical, initiation, propagation, termination, nucleophile, electrophile, nucleophilic, electrophilic, addition, substitution, elimination, hydrolysis, condensation, oxidation and reduction

8. Understand and use terminology associated with types of organic mechanisms: free-radical substitution, electrophilic addition, nucleophilic substitution, nucleophilic addition

9. Describe organic molecules as either straight-chained, branched or cyclic

10. Describe and explain the shape of, and bond angles in, molecules containing sp, sp2, and sp3 hybridized atoms

11. Describe the arrangement of σ and π bonds in molecules containing sp, sp2, and sp3 hybridized atoms

12. Understand and use the term planar when describing the arrangement of atoms in organic molecules

13. Describe structural isomerism and its division into chain, positional and functional group isomerism

14. Describe stereoisomerism and its division into geometrical (cis/trans) and optical isomerism

15. Describe geometrical (cis/trans) isomerism in alkenes, and explain its origin in terms of restricted rotation due to the presence of π bonds

16. Describe and explain the shape of benzene and other aromatic molecules, including sp hybridisation, in terms of σ bonds and a delocalised π system

17. Explain what is meant by a chiral center and that such a center gives rise to two optical isomers (enantiomers)

18. Identify chiral centers and geometrical and deduce possible isomers

19. Understand that enantiomers have identical physical and chemical properties except for their ability to rotate plane-polarized light and potential biological activity.

20. Understand and use the terms optically active and racemic mixture.

21. Describe the effect on plane-polarized light of the two optical isomers of a single substance.

22. Explain the significance of chirality in the synthetic preparation of drug molecules, including the potential different biological activity of enantiomers, the need to separate racemic mixtures, and the use of chiral catalysts to produce a single pure optical isomer.

23. Note that compounds can have more than one chiral center, but knowledge of meso compounds and nomenclature such as diastereoisomers is not required.

Skeletal formulas should be discussed in the course.

- The general formulas (eg CnH2n+2) of alkanes, alkenes, alkynes, ketones, alcohols, aldehydes and carboxylic acids should be known.

- The distinction between class names and functional group names needs to be made. Eg for OH, hydroxyl is the functional group whereas alcohol is the class name.

- The following nomenclature should be covered– non-cyclic alkanes and halogenoalkanes up to halohexanes.

– alkenes up to hexene and alkynes up to hexyne.

– compounds up to six carbon atoms (in the basic chain for nomenclature purposes) containing only one of the functional groupshydroxyl, ether, carbonyl (from aldehydes or ketones), ester and carbox

Reference should be made to initiation, propagation and termination steps in free-radical substitution reactions.

Free radicals should be represented by a single dot.

A homologous series is a series of compounds of the same family, with the same general formula, which differ from each other by a common structural unit.

- Structural formulas can be represented in full and condensed format.

- Structural isomers are compounds with the same molecular formula but different arrangements of atoms.

- Functional groups are the reactive parts of molecules.

- Saturated compounds contain single bonds only and unsaturated compounds contain double or triple bonds.

- Benzene is an aromatic, unsaturated hydrocarbon.

Alkanes:

- Alkanes have low reactivity and undergo free-radical substitution reactions. Alkenes:

- Alkenes are more reactive than alkanes and undergo addition reactions. Bromine water can be used to distinguish between alkenes and alkanes. Alcohols:

- Alcohols undergo esterification (or condensation) reactions with acids and some undergo oxidation reactions. Halogenoalkanes:

- Halogenoalkanes are more reactive than alkanes. They can undergo (nucleophilic) substitution reactions. A nucleophile is an electron-rich species containing a lone pair that it donates to an electron-deficient carbon. Polymers:

- Addition polymers consist of a wide range of monomers and form the basis of the plastics industry. Benzene:

- Benzene does not readily undergo addition reactions but does undergo electrophilic substitution reactions

Types of organic reactions:

Nucleophilic Substitution Reactions:

- SN1 represents a nucleophilic unimolecular substitution reaction and SN2 represents a nucleophilic bimolecular substitution reaction. SN1 involves a carbocation intermediate. SN2 involves a concerted reaction with a transition state.

- For tertiary halogenoalkanes the predominant mechanism is SN1 and for primary halogenoalkanes it is SN2. Both mechanisms occur for secondary halogenoalkanes.

- The rate determining step (slow step) in an SN1 reaction depends only on the concentration of the halogenoalkane, rate = k[halogenoalkane]. For SN2, rate = k[halogenoalkane][nucleophile]. SN2 is stereospecific with an inversion of configuration at the carbon.

- SN2 reactions are best conducted using aprotic, non-polar solvents and SN1 reactions are best conducted using protic, polar solvents. Electrophilic

Addition Reactions:

- An electrophile is an electron-deficient species that can accept electron pairs from a nucleophile. Electrophiles are Lewis acids. Markovnikov’s rule can be applied to predict the major product in electrophilic addition reactions of unsymmetrical alkenes with hydrogen halides and interhalogens. The formation of the major product can be explained in terms of the relative stability of possible carbocations in the reaction mechanism. Electrophilic

Substitution Reactions:

- Benzene is the simplest aromatic hydrocarbon compound (or arene) and has a delocalized structure of π bonds around its ring. Each carbon to carbon bond has a bond order of 1.5. Benzene is susceptible to attack by electrophiles. Reduction Reactions:

- Carboxylic acids can be reduced to primary alcohols (via the aldehyde). Ketones can be reduced to secondary alcohols. Typical reducing agents are lithium aluminium hydride (used to reduce carboxylic acids) and sodium borohydride.

Usage:

"Fractional distillation makes great use of many petrochemicals.

- Dyes, pesticides, herbicides, explosives, soap, cosmetics, synthetic scents and flavourings.

- Alkane usage as fuels.

- The role of ethene in fruit ripening.

- Alcohols, usage as fuel additives.

- Alcohols, role in the breathalyser.

- Esters, varied uses—perfumes, food flavourings, solvents, nitroglycerin, biofuels and painkiller"

Organic chemistry focuses on the chemistry of compounds containing carbon.

Structure, bonding and chemical reactions involving functional group interconversions are key strands in organic chemistry.

Key organic reaction types include nucleophilic substitution, electrophilic addition, electrophilic substitution and redox reactions. Reaction mechanisms vary

and help in understanding the different types of reaction taking place.

Stereoisomerism involves isomers which have different arrangements of atoms in space but do not differ in connectivity or bond multiplicity (ie whether single,

double or triple) between the isomers themselves.

The curriculums cover the basics of organic chemistry, including hydrocarbons, functional groups, isomerism, and stereochemistry. However, the NCC 2023 curriculum expands on these topics by including more in-depth knowledge on terminology, molecular structure, and chemical reactions. Additionally, the NCC 2023 curriculum places a greater emphasis on the significance of organic chemistry in drug synthesis and biological activity. Some topics from the FSC 2006 curriculum, such as the use of coal as a source of organic compounds and nomenclature for diastereoisomers, are no longer included in the NCC 2023 curriculum. This reflects changes in the field and the need to prioritize relevant and current information for students.

1.2—empirical and molecular formulas Topics 4.2 and 4.3—Lewis (electron dot) structures, multiple bonds, VSEPR theory, resonance and bond and molecular polarity 4.4—intermolecular forces 5.3—exothermic reactions and bond enthalpies 8.4—weak acids Option A.5—materials and polymers Options B.2 and B.7—proteins Option D.9—organic structure in medici

9.1—redox processes Option A.5—polymers Option B.3—lipids

Hydrocarbons
(Alkanes, Alkenes, Alkynes, Benzene)
(Organic Chemistry)

- Classify hydrocarbons as aliphatic and aromatic.

- Describe nomenclature of alkanes and cycloalkanes.

- Explain the shapes of alkanes and cycloalkanes exemplified by ethane and cyclopropane.

- Explain unreative nature of alkanes towards polar reagents.

- Define homolytic and heterolytic fission, free radical initiation, propagation and termination.

- Describe the mechanism of free radical substitution in alkanes exemplified by methane and ethane.

- Identify organic redox reactions.

- Explain what is meant by a chiral centre and show that such a centre gives rise to optical isomerism.

- Identify chiral centers in given structural formula of a molecule.

- Explain the nomenclature of alkenes.

- Explain shape of ethene molecule in terms of sigma and pi C-C bonds.

- Describe the structure and reactivity of alkenes as exemplified by ethene.

- Define and explain with suitable examples the terms isomerism, stereoisomerism and structural isomerism.

- Explain dehydration of alcohols and dehydrohalogenation of RX for the preparation of ethene.

- Describe the chemistry of alkenes by the following reactions of ethene:

- Hydrogenation, hydrohalogenation, hydration, halogenation, halohydration, epoxidation, ozonolysis, polymerization.

- Explain the concept of conjugation in alkenes having alternate double bonds.

- Use the IUPAC naming system for alkenes.

- Explain the shape of benzene molecule (molecular orbital aspect).

- Define resonance, resonance energy and relative stability.

- Compare the reactivity of benzene with alkanes and alkenes.

- Describe what is meant by the term delocalized electrons in the context of the

benzene ring.

- Describe addition reactions of benzene and methyl benzene.

- Describe the mechanism of electrophilic substitution in benzene.

- Discuss chemistry of benzene and methyl benzene by nitration, sulphonation, halogenation, Friedal Craft's alkylation and acylation.

- Apply the knowledge of positions of substituents in the electrophilic substitution of benzene.

- Use the IUPAC naming system for alkynes.

- Compare the reactivity of alkynes with alkanes, alkenes and arenes.

- Discuss the shape of alkynes in terms of sigma and pi C-C bonds.

- Describe the preparation of alkynes using elimination reactions.

- Describe acidity of alkynes.

- Discuss chemistry of alkynes by hydrogenation, hydrohalogenation, hydration, bromination, ozonolysis, and reaction with metals.

- Describe and differentiate between substitution and addition reactions.

- Explain isomerism in alkanes, alkenes, alkynes and substituted benzene.

- Draw different possible ring structures of benzene (Kekule structures).

- Draw straight chain structures of alkanes, alkanes and alkynes up to 10 carbon atoms.

- Identify and link uses of various hydrocarbons used in daily life. ( understanding )

- Identify various hydrocarbons which will be important as fuels for the future energy needs of Pakistan understandin

1 recall the reactions (reagents and conditions) by which alkanes can be produced:

(a) addition of hydrogen to an alkene in a hydrogenation reaction, H2(g) and Pt/Ni catalyst and heat

(b) cracking of a longer chain alkane, heat with Al 2O3

2 describe:

(a) the complete and incomplete combustion of alkanes

(b) the free-radical substitution of alkanes by Cl 2 or Br2 in the presence of ultraviolet light, as exemplified by

the reactions of ethane

3 describe the mechanism of free-radical substitution with reference to the initiation, propagation and

termination steps

4 suggest how cracking can be used to obtain more useful alkanes and alkenes of lower Mr from heavier crude

oil fractions

5 understand the general unreactivity of alkanes, including towards polar reagents in terms of the strength of

the C–H bonds and their relative lack of polarity

6 recognise the environmental consequences of carbon monoxide, oxides of nitrogen and unburnt

hydrocarbons arising from the combustion of alkanes in the internal combustion engine and of their catalytic

removal

1 recall the reactions (including reagents and conditions) by which alkenes can be produced:

(a) elimination of HX from a halogenoalkane by ethanolic NaOH and heat

(b) dehydration of an alcohol, by using a heated catalyst (e.g. Al 2O3) or a concentrated acid

(c) cracking of a longer chain alkane

2 describe the following reactions of alkenes:

(a) the electrophilic addition of

(i) hydrogen in a hydrogenation reaction, H2(g) and Pt/Ni catalyst and heat

(ii) steam, H2O(g) and H3PO4 catalyst

(iii) a hydrogen halide, HX(g) at room temperature

(iv) a halogen, X2

(b) the oxidation by cold dilute acidified KMnO4 to form the diol

(c) the oxidation by hot concentrated acidified KMnO4 leading to the rupture of the carbon–carbon double

bond and the identities of the subsequent products to determine the position of alkene linkages in larger

molecules

(d) addition polymerisation exemplified by the reactions of ethene and propene

3 describe the use of aqueous bromine to show the presence of a C=C bond

4 describe the mechanism of electrophilic addition in alkenes, using bromine / ethene and hydrogen

bromide / propene as examples

5 describe and explain the inductive effects of alkyl groups on the stability of primary, secondary and tertiary

cations formed during electrophilic addition (this should be used to explain Markovnikov addition)

1 describe the chemistry of arenes as exemplified by the following reactions of benzene and methylbenzene:

(a) substitution reactions with Cl 2 and with Br2 in the presence of a catalyst, AlCl 3 or Al Br3, to form

halogenoarenes (aryl halides)

(b) nitration with a mixture of concentrated HNO3 and concentrated H2SO4 at a temperature between

25 °C and 60 °C

(c) Friedel–Crafts alkylation by CH3Cl and AlCl 3 and heat

(d) Friedel–Crafts acylation by CH3COCl and AlCl 3 and heat

(e) complete oxidation of the side-chain using hot alkaline KMnO4 and then dilute acid to give a benzoic

acid

(f) hydrogenation of the benzene ring using H2 and Pt/Ni catalyst and heat to form a cyclohexane ring

2 describe the mechanism of electrophilic substitution in arenes:

(a) as exemplified by the formation of nitrobenzene and bromobenzene

(b) with regards to the effect of delocalisation (aromatic stabilisation) of electrons in arenes to explain the

predomination of substitution over addition

3 predict whether halogenation will occur in the side-chain or in the aromatic ring in arenes depending on

reaction conditions

4 describe that in the electrophilic substitution of arenes, different substituents direct to different ring

positions (limited to the directing effects of –NH2, –OH, –R, –NO2, –COOH and –COR)

Alkanes:

- Alkanes have low reactivity and undergo free-radical substitution reactions. Alkenes:

- Alkenes are more reactive than alkanes and undergo addition reactions. Bromine water can be used to distinguish between alkenes and alkanes.

Classify hydrocarbons as aliphatic and aromatic.

Describe nomenclature of alkanes and cycloalkanes.

Explain the shapes of alkanes and cycloalkanes exemplified by ethane and cyclopropane.

Explain unreactive nature of alkanes towards polar reagents.

Define homolytic and heterolytic fission, free radical initiation, propagation and termination.

Describe the mechanism of free radical substitution in alkanes exemplified by methane and ethane.

Identify organic redox reactions.

Explain what is meant by a chiral center and show that such a center gives rise to optical isomerism.

Identify chiral centers in given structural formula of a molecule.

Explain the nomenclature of alkenes.

Explain shape of ethene molecule in terms of sigma and pi C-C bonds.

Describe the structure and reactivity of alkenes as exemplified by ethene.

Define and explain with suitable examples the terms isomerism, stereoisomerism and structural isomerism.

Explain dehydration of alcohols and dehydrohalogenation of RX for the preparation of ethene.

Describe the chemistry of alkenes by the following reactions of ethene: hydrogenation, hydrohalogenation, hydration, halogenation, halohydration, epoxidation, ozonolysis, polymerization.

Explain the concept of conjugation in alkenes having alternate double bonds.

Use the IUPAC naming system for alkenes.

Explain the shape of benzene molecule (molecular orbital aspect).

Define resonance, resonance energy and relative stability.

Compare the reactivity of benzene with alkanes and alkenes.

1 recall the reactions (reagents and conditions) by which alkanes can be produced:

(a) addition of hydrogen to an alkene in a hydrogenation reaction, H2(g) and Pt/Ni catalyst and heat

(b) cracking of a longer chain alkane, heat with Al 2O3

2 describe:

(a) the complete and incomplete combustion of alkanes

(b) the free-radical substitution of alkanes by Cl 2 or Br2 in the presence of ultraviolet light, as exemplified by

the reactions of ethane

3 describe the mechanism of free-radical substitution with reference to the initiation, propagation and

termination steps

4 suggest how cracking can be used to obtain more useful alkanes and alkenes of lower Mr from heavier crude

oil fractions

5 understand the general unreactivity of alkanes, including towards polar reagents in terms of the strength of

the C–H bonds and their relative lack of polarity

6 recognise the environmental consequences of carbon monoxide, oxides of nitrogen and unburnt

hydrocarbons arising from the combustion of alkanes in the internal combustion engine and of their catalytic

removal

1 recall the reactions (including reagents and conditions) by which alkenes can be produced:

(a) elimination of HX from a halogenoalkane by ethanolic NaOH and heat

(b) dehydration of an alcohol, by using a heated catalyst (e.g. Al 2O3) or a concentrated acid

(c) cracking of a longer chain alkane

2 describe the following reactions of alkenes:

(a) the electrophilic addition of

(i) hydrogen in a hydrogenation reaction, H2(g) and Pt/Ni catalyst and heat

(ii) steam, H2O(g) and H3PO4 catalyst

(iii) a hydrogen halide, HX(g) at room temperature

(iv) a halogen, X2

(b) the oxidation by cold dilute acidified KMnO4 to form the diol

(c) the oxidation by hot concentrated acidified KMnO4 leading to the rupture of the carbon–carbon double

bond and the identities of the subsequent products to determine the position of alkene linkages in larger

molecules

(d) addition polymerisation exemplified by the reactions of ethene and propene

3 describe the use of aqueous bromine to show the presence of a C=C bond

4 describe the mechanism of electrophilic addition in alkenes, using bromine / ethene and hydrogen

bromide / propene as examples

5 describe and explain the inductive effects of alkyl groups on the stability of primary, secondary and tertiary

cations formed during electrophilic addition (this should be used to explain Markovnikov addition)

1 describe the chemistry of arenes as exemplified by the following reactions of benzene and methylbenzene:

(a) substitution reactions with Cl 2 and with Br2 in the presence of a catalyst, AlCl 3 or Al Br3, to form

halogenoarenes (aryl halides)

(b) nitration with a mixture of concentrated HNO3 and concentrated H2SO4 at a temperature between

25 °C and 60 °C

(c) Friedel–Crafts alkylation by CH3Cl and AlCl 3 and heat

(d) Friedel–Crafts acylation by CH3COCl and AlCl 3 and heat

(e) complete oxidation of the side-chain using hot alkaline KMnO4 and then dilute acid to give a benzoic

acid

(f) hydrogenation of the benzene ring using H2 and Pt/Ni catalyst and heat to form a cyclohexane ring

2 describe the mechanism of electrophilic substitution in arenes:

(a) as exemplified by the formation of nitrobenzene and bromobenzene

(b) with regards to the effect of delocalisation (aromatic stabilisation) of electrons in arenes to explain the

predomination of substitution over addition

3 predict whether halogenation will occur in the side-chain or in the aromatic ring in arenes depending on

reaction conditions

4 describe that in the electrophilic substitution of arenes, different substituents direct to different ring

positions (limited to the directing effects of –NH2, –OH, –R, –NO2, –COOH and –COR)

Usage:

- Alkane usage as fuels.

- The role of ethene in fruit ripening.

 

The topics covered in 2006 and NCC 2023 include hydrocarbons, their classification, nomenclature, shapes, reactivity, and isomerism. However, NCC 2023 also covers additional topics such as the concept of conjugation in alkenes, the shape of benzene molecule in terms of molecular orbital aspect, resonance, resonance energy, and relative stability. In addition, it compares the reactivity of benzene with alkanes and alkenes.

 

Halogenalkanes
(Organic Chemistry)

- Name alkyl halides using IUPAC system.

- Discuss the structure and reactivity of RX.

- Describe the preparation of RX by the reaction of alcohols with HX, SOCI2 and PX3 and by radical halogenation of alkanes.

- Describe the mechanism and types of nucleophilic substitution reactions.

- Describe the mechanism and types of elimination reactions.

- Describe the preparation and reactivity of Grignard's Reagents.

- Discuss chemistry of Grignard's reagent by the addition of aldehydes, ketones, esters and carbon dioxide.

- Identify organometallic compounds in medicines.

- Compare haemoglobin and chlorophyll.

- Recognize alkyl halides as precursors of many organic compounds.

1 recall the reactions (reagents and conditions) by which halogenoalkanes can be produced:

(a) the free-radical substitution of alkanes by Cl 2 or Br2 in the presence of ultraviolet light, as exemplified by

the reactions of ethane

(b) electrophilic addition of an alkene with a halogen, X2, or hydrogen halide, HX(g), at room temperature

(c) substitution of an alcohol, e.g. by reaction with HX or KBr with H2SO4 or H3PO4; or with PCl 3 and heat;

or with PCl 5; or with SOCl 2

2 classify halogenoalkanes into primary, secondary and tertiary

3 describe the following nucleophilic substitution reactions:

(a) the reaction with NaOH(aq) and heat to produce an alcohol

(b) the reaction with KCN in ethanol and heat to produce a nitrile

(c) the reaction with NH3 in ethanol heated under pressure to produce an amine

(d) the reaction with aqueous silver nitrate in ethanol as a method of identifying the halogen present as

exemplified by bromoethane

4 describe the elimination reaction with NaOH in ethanol and heat to produce an alkene as exemplified by

bromoethane

5 describe the SN1 and SN2 mechanisms of nucleophilic substitution in halogenoalkanes including the inductive

effects of alkyl groups

6 recall that primary halogenoalkanes tend to react via the SN2 mechanism; tertiary halogenoalkanes via the

SN1 mechanism; and secondary halogenoalkanes by a mixture of the two, depending on structure

7 describe and explain the different reactivities of halogenoalkanes (with particular reference to the relative

strengths of the C–X bonds as exemplified by the reactions of halogenoalkanes with aqueous silver nitrates)

1 recall the reactions by which halogenoarenes can be produced: substitution of an arene with Cl2 or Br2

in the presence of a catalyst, Al Cl3 or Al Br3 to form a halogenoarene, exemplified by benzene to form

chlorobenzene and methylbenzene to form 2-chloromethylbenzene and 4-chloromethylbenzene

2 explain the difference in reactivity between a halogenoalkane and a halogenoarene as exemplified by

chloroethane and chlorobenzene

- Halogenoalkanes are more reactive than alkanes. They can undergo (nucleophilic) substitution reactions. A nucleophile is an electron-rich species containing a lone pair that it donates to an electron-deficient carbon.

Recall the reactions (reagents and conditions) by which halogenoalkanes can be produced:

(a) the free-radical substitution of alkanes by Cl 2 or Br2 in the presence of ultraviolet light, as exemplified by

the reactions of ethane

(b) electrophilic addition of an alkene with a halogen, X2, or hydrogen halide, HX(g), at room temperature

(c) substitution of an alcohol, e.g. by reaction with HX or KBr with H2SO4 or H3PO4; or with PCl 3 and heat;

or with PCl 5; or with SOCl 2

2 classify halogenoalkanes into primary, secondary and tertiary

3 describe the following nucleophilic substitution reactions:

(a) the reaction with NaOH(aq) and heat to produce an alcohol

(b) the reaction with KCN in ethanol and heat to produce a nitrile

(c) the reaction with NH3 in ethanol heated under pressure to produce an amine

(d) the reaction with aqueous silver nitrate in ethanol as a method of identifying the halogen present as

exemplified by bromoethane

4 describe the elimination reaction with NaOH in ethanol and heat to produce an alkene as exemplified by

bromoethane

5 describe the SN1 and SN2 mechanisms of nucleophilic substitution in halogenoalkanes including the inductive

effects of alkyl groups

6 recall that primary halogenoalkanes tend to react via the SN2 mechanism; tertiary halogenoalkanes via the

SN1 mechanism; and secondary halogenoalkanes by a mixture of the two, depending on structure

7 describe and explain the different reactivities of halogenoalkanes (with particular reference to the relative

strengths of the C–X bonds as exemplified by the reactions of halogenoalkanes with aqueous silver nitrates)

 

 

curriculums cover the chemistry of alkyl halides and their reactions, including nucleophilic substitution and elimination reactions. The NCC 2023 curriculum includes additional topics such as the chemistry of Grignard reagents, organometallic compounds in medicines, and a more detailed discussion of halogenoalkane reactions and mechanisms. Some topics that were present in the 2006 curriculum, such as the preparation of alkyl halides by radical halogenation of alkanes, are not explicitly mentioned in the NCC 2023 curriculum. These changes reflect updates in the field of organic chemistry and a shift towards a more detailed and integrated understanding of the subject.

 

Hydroxy Compounds
(alcohols and phenols)
(Organic Chemistry)

- Explain nomenclature, structure and acidity of alcohols as exemplified by ethanol.

- Describe the preparation of alcohols by reduction of aldehydes, ketones, carboxylic acids and esters.

- Explain reactivity of alcohols.

- Describe the chemistry of alcohols by preparation of ethers and esters, oxidative cleavage of 1, 2-diols.

- Discuss thiols (RSH).

- Explain the nomenclature, structure and acidity of phenols.

- Describe the preparation of phenol from benzene sulphonic acid, chlorobenzene, acidic oxidation of cumene and hydrolysis of diazonium salts.

- Discuss the reactivity of phenol and their chemistry by electrophilic aromatic substitution, reaction with Na metal and oxidation.

- Differentiate between alcohol and phenol.

- Describe isomerism in alcohols and phenols.

- Identify ethers from their formula.

- Identify alcohols using appropriate laboratory tests.

- Identify phenols using appropriate laboratory tests.

- Determine boiling points of alcohols and phenols in laboratory.

- Explain the role of disinfectants in hygiene.

- Differentiate between disinfectants and antiseptics.

- Reco nize that ethers are used in anesthesia

1 recall the reactions (reagents and conditions) by which alcohols can be produced:

(a) electrophilic addition of steam to an alkene, H2O(g) and H3PO4 catalyst

(b) reaction of alkenes with cold dilute acidified potassium manganate(VII) to form a diol

(c) substitution of a halogenoalkane using NaOH(aq) and heat

(d) reduction of an aldehyde or ketone using NaBH4 or LiAl H4

(e) reduction of a carboxylic acid using LiAl H4

(f) hydrolysis of an ester using dilute acid or dilute alkali and heat

2 describe:

(a) the reaction with oxygen (combustion)

(b) substitution to halogenoalkanes, e.g. by reaction with HX or KBr with H2SO4 or H3PO4; or with PCl 3 and

heat; or with PCl 5; or with SOCl 2

(c) the reaction with Na(s)

(d) oxidation with acidified K2Cr2O7 or acidified KMnO4 to:

(i) carbonyl compounds by distillation

(ii) carboxylic acids by refluxing

(primary alcohols give aldehydes which can be further oxidised to carboxylic acids, secondary alcohols

give ketones, tertiary alcohols cannot be oxidised)

(e) dehydration to an alkene, by using a heated catalyst, e.g. Al 2O3 or a concentrated acid

(f) formation of esters by reaction with carboxylic acids and concentrated H2SO4 or H3PO4 as catalyst as

exemplified by ethanol

3 (a) classify alcohols as primary, secondary and tertiary alcohols, to include examples with more than one

alcohol group

(b) state characteristic distinguishing reactions, e.g. mild oxidation with acidified K2Cr2O7, colour change

from orange to green

4 deduce the presence of a CH3CH(OH)– group in an alcohol, CH3CH(OH)–R, from its reaction with alkaline

I2(aq) to form a yellow precipitate of tri-iodomethane and an ion, RCO2

5 explain the acidity of alcohols compared with water

1 describe the reaction with acyl chlorides to form esters using ethyl ethanoate

1 recall the reactions (reagents and conditions) by which phenol can be produced:

(a) reaction of phenylamine with HNO2 or NaNO2 and dilute acid below 10 °C to produce the diazonium

salt; further warming of the diazonium salt with H2O to give phenol

2 recall the chemistry of phenol, as exemplified by the following reactions:

(a) with bases, for example NaOH(aq) to produce sodium phenoxide

(b) with Na(s) to produce sodium phenoxide and H2(g)

(c) in NaOH(aq) with diazonium salts, to give azo compounds

(d) nitration of the aromatic ring with dilute HNO3(aq) at room temperature to give a mixture of

2-nitrophenol and 4-nitrophenol

(e) bromination of the aromatic ring with Br2(aq) to form 2,4,6-tribromophenol

3 explain the acidity of phenol

4 describe and explain the relative acidities of water, phenol and ethanol

5 explain why the reagents and conditions for the nitration and bromination of phenol are different from those

for benzene

6 recall that the hydroxyl group of a phenol directs to the 2-, 4- and 6-positions

7 apply knowledge of the reactions of phenol to those of other phenolic compounds, e.g. naphthol

- Alcohols undergo esterification (or condensation) reactions with acids and some undergo oxidation reactions.

1 recall the reactions (reagents and conditions) by which alcohols can be produced:

(a) electrophilic addition of steam to an alkene, H2O(g) and H3PO4 catalyst

(b) reaction of alkenes with cold dilute acidified potassium manganate(VII) to form a diol

(c) substitution of a halogenoalkane using NaOH(aq) and heat

(d) reduction of an aldehyde or ketone using NaBH4 or LiAl H4

(e) reduction of a carboxylic acid using LiAl H4

(f) hydrolysis of an ester using dilute acid or dilute alkali and heat

2 describe:

(a) the reaction with oxygen (combustion)

(b) substitution to halogenoalkanes, e.g. by reaction with HX or KBr with H2SO4 or H3PO4; or with PCl 3 and

heat; or with PCl 5; or with SOCl 2

(c) the reaction with Na(s)

(d) oxidation with acidified K2Cr2O7 or acidified KMnO4 to:

(i) carbonyl compounds by distillation

(ii) carboxylic acids by refluxing

(primary alcohols give aldehydes which can be further oxidised to carboxylic acids, secondary alcohols

give ketones, tertiary alcohols cannot be oxidised)

(e) dehydration to an alkene, by using a heated catalyst, e.g. Al 2O3 or a concentrated acid

(f) formation of esters by reaction with carboxylic acids and concentrated H2SO4 or H3PO4 as catalyst as

exemplified by ethanol

3 (a) classify alcohols as primary, secondary and tertiary alcohols, to include examples with more than one

alcohol group

(b) state characteristic distinguishing reactions, e.g. mild oxidation with acidified K2Cr2O7, colour change

from orange to green

4 deduce the presence of a CH3CH(OH)– group in an alcohol, CH3CH(OH)–R, from its reaction with alkaline

I2(aq) to form a yellow precipitate of tri-iodomethane and an ion, RCO2

5 explain the acidity of alcohols compared with water

1 describe the reaction with acyl chlorides to form esters using ethyl ethanoate

1 recall the reactions (reagents and conditions) by which phenol can be produced:

(a) reaction of phenylamine with HNO2 or NaNO2 and dilute acid below 10 °C to produce the diazonium

salt; further warming of the diazonium salt with H2O to give phenol

2 recall the chemistry of phenol, as exemplified by the following reactions:

(a) with bases, for example NaOH(aq) to produce sodium phenoxide

(b) with Na(s) to produce sodium phenoxide and H2(g)

(c) in NaOH(aq) with diazonium salts, to give azo compounds

(d) nitration of the aromatic ring with dilute HNO3(aq) at room temperature to give a mixture of

2-nitrophenol and 4-nitrophenol

(e) bromination of the aromatic ring with Br2(aq) to form 2,4,6-tribromophenol

3 explain the acidity of phenol

4 describe and explain the relative acidities of water, phenol and ethanol

5 explain why the reagents and conditions for the nitration and bromination of phenol are different from those

for benzene

6 recall that the hydroxyl group of a phenol directs to the 2-, 4- and 6-positions

7 apply knowledge of the reactions of phenol to those of other phenolic compounds, e.g. naphthol

Usage:

Alcohols, role in the breathalyser.

 

it can be observed that there is a significant overlap in the topics covered, such as the nomenclature, structure, and reactivity of alcohols and phenols, their preparation and identification using laboratory tests, and their chemistry by various reactions. However, NCC 2023 adds some additional topics, such as the hydrolysis of esters and oxidative cleavage of 1,2-diols, whereas the 2006 curriculum includes the chemistry of aldehydes and ketones. The changes in the curriculum reflect the evolving needs and demands of the scientific community, as well as advancements in the field of chemistry.

 

Carbonyl Compounds
(Carboxylic Acids, Aldehydes, Ketones, Esters)
(Organic Chemistry)

- Explain nomenclature and structure of aldehydes and ketones.

- Discuss the preparation of aldehydes and ketones by ozonolysis of alkenes, hydration of alkynes, oxidation of alcohols and Friedal Craft's acylation of aromatics.

- Describe reactivity of aldehydes and ketones and their comparison.

- Describe acid and base catalysed nucleophilic addition reactions of aldehydes and ketones.

- Discuss the chemistry of aldehydes and ketones by their reduction to hydrocarbons, alcohols, by using carbon nucleophiles, nitrogen nucleophiles and oxygen nucleophiles.

- Describe oxidation reactions of aldehydes and ketones.

- Describe isomerism in aldehydes and ketones.

- Identify aldehydes in the laboratory tests.

- Identify ketones using appropriate laboratory tests.

- Determine melting or boiling points of aldehydes and ketones in laboratory.

- Explain how oxidation and reduction alters the structure of organic compounds.

- Explain the need to limit exposure to formaldehyde vapors as used in adhesives, varnishes, paints, foam insulation and permanent press clothing.

- Describe glucose and fructose as examples of aldehydes and ketones

- Describe preparation of carboxylic acids by carbonation of Grignard's Reagent, hydrolysis of nitriles, oxidation of primary alcohols, oxidation of aldehydes and oxidation of alkyl benzenes.

- Discuss reactivity of carboxylic acids.

- Describe the chemistry of carboxylic acids by conversion to carboxylic acid

derivatives: acyl halides, acid anhydrides, esters, amides and reactions involving inter conversion of these.

- Describe reactions of carboxylic acid derivatives.

- Describe isomerism in carboxylic acids.

- Identify carboxylic acids in the laboratory

- Determine melting or boiling points of carboxylic acids in laboratory.

- List carboxylic acids present in fruits, vegetables and other natural products.

- Link different carboxylic acids with their characteristic taste.

- Recognize carboxylic acids used as preservatives in food and food products

1 recall the reactions (reagents and conditions) by which aldehydes and ketones can be produced:

(a) the oxidation of primary alcohols using acidified K2Cr2O7 or acidified KMnO4 and distillation to produce

aldehydes

(b) the oxidation of secondary alcohols using acidified K2Cr2O7 or acidified KMnO4 and distillation to

produce ketones

2 describe:

(a) the reduction of aldehydes and ketones, using NaBH4 or LiAl H4 to produce alcohols

(b) the reaction of aldehydes and ketones with HCN, KCN as catalyst, and heat to produce hydroxynitriles

exemplified by ethanal and propanone

3 describe the mechanism of the nucleophilic addition reactions of hydrogen cyanide with aldehydes and

ketones in 17.1.2(b)

4 describe the use of 2,4-dinitrophenylhydrazine (2,4-DNPH reagent) to detect the presence of carbonyl

compounds

5 deduce the nature (aldehyde or ketone) of an unknown carbonyl compound from the results of simple tests

(Fehling’s and Tollens’ reagents; ease of oxidation)

6 deduce the presence of a CH3CO – group in an aldehyde or ketone, CH3CO–R, from its reaction with alkaline

I2(aq) to form a yellow precipitate of tri-iodomethane and an ion, RCO2

1 recall the reactions by which carboxylic acids can be produced:

(a) oxidation of primary alcohols and aldehydes with acidified K2Cr2O7 or acidified KMnO4 and refluxing

(b) hydrolysis of nitriles with dilute acid or dilute alkali followed by acidification

(c) hydrolysis of esters with dilute acid or dilute alkali and heat followed by acidification

2 describe:

(a) the redox reaction with reactive metals to produce a salt and H2(g)

(b) the neutralisation reaction with alkalis to produce a salt and H2O(l )

(c) the acid–base reaction with carbonates to produce a salt and H2O(l) and CO2(g)

(d) esterification with alcohols with concentrated H2SO4 as catalyst

(e) reduction by LiAl H4 to form a primary alcohol

1 recall the reaction (reagents and conditions) by which esters can be produced:

(a) the condensation reaction between an alcohol and a carboxylic acid with concentrated H2SO4 as catalyst

2 describe the hydrolysis of esters by dilute acid and by dilute alkali and heat

 

1 recall the reactions (reagents and conditions) by which aldehydes and ketones can be produced:

(a) the oxidation of primary alcohols using acidified K2Cr2O7 or acidified KMnO4 and distillation to produce

aldehydes

(b) the oxidation of secondary alcohols using acidified K2Cr2O7 or acidified KMnO4 and distillation to

produce ketones

2 describe:

(a) the reduction of aldehydes and ketones, using NaBH4 or LiAl H4 to produce alcohols

(b) the reaction of aldehydes and ketones with HCN, KCN as catalyst, and heat to produce hydroxynitriles

exemplified by ethanal and propanone

3 describe the mechanism of the nucleophilic addition reactions of hydrogen cyanide with aldehydes and

ketones in 17.1.2(b)

4 describe the use of 2,4-dinitrophenylhydrazine (2,4-DNPH reagent) to detect the presence of carbonyl

compounds

5 deduce the nature (aldehyde or ketone) of an unknown carbonyl compound from the results of simple tests

(Fehling’s and Tollens’ reagents; ease of oxidation)

6 deduce the presence of a CH3CO – group in an aldehyde or ketone, CH3CO–R, from its reaction with alkaline

I2(aq) to form a yellow precipitate of tri-iodomethane and an ion, RCO2

1 recall the reactions by which carboxylic acids can be produced:

(a) oxidation of primary alcohols and aldehydes with acidified K2Cr2O7 or acidified KMnO4 and refluxing

(b) hydrolysis of nitriles with dilute acid or dilute alkali followed by acidification

(c) hydrolysis of esters with dilute acid or dilute alkali and heat followed by acidification

2 describe:

(a) the redox reaction with reactive metals to produce a salt and H2(g)

(b) the neutralisation reaction with alkalis to produce a salt and H2O(l )

(c) the acid–base reaction with carbonates to produce a salt and H2O(l) and CO2(g)

(d) esterification with alcohols with concentrated H2SO4 as catalyst

(e) reduction by LiAl H4 to form a primary alcohol

1 recall the reaction (reagents and conditions) by which esters can be produced:

(a) the condensation reaction between an alcohol and a carboxylic acid with concentrated H2SO4 as catalyst

2 describe the hydrolysis of esters by dilute acid and by dilute alkali and heat

1 recall the reaction by which benzoic acid can be produced:

(a) reaction of an alkylbenzene with hot alkaline KMnO4 and then dilute acid, exemplified by

methylbenzene

2 describe the reaction of carboxylic acids with PCl 3 and heat, PCl 5, or SOCl 2 to form acyl chlorides

3 recognise that some carboxylic acids can be further oxidised:

(a) the oxidation of methanoic acid, HCOOH, with Fehling’s reagent or Tollens’ reagent or acidified KMnO4

or acidified K2Cr2O7 to carbon dioxide and water

(b) the oxidation of ethanedioic acid, HOOCCOOH, with warm acidified KMnO4 to carbon dioxide

4 describe and explain the relative acidities of carboxylic acids, phenols and alcohols

5 describe and explain the relative acidities of chlorine-substituted carboxylic acids

recall the reaction by which esters can be produced:

(a) reaction of alcohols with acyl chlorides using the formation of ethyl ethanoate and phenyl benzoate as

examples

Acyl Chlorides

1 recall the reactions (reagents and conditions) by which acyl chlorides can be produced:

(a) reaction of carboxylic acids with PCl 3 and heat, PCl 5, or SOCl 2

2 describe the following reactions of acyl chlorides:

(a) hydrolysis on addition of water at room temperature to give the carboxylic acid and HCl

(b) reaction with an alcohol at room temperature to produce an ester and HCl

(c) reaction with phenol at room temperature to produce an ester and HCl

(d) reaction with ammonia at room temperature to produce an amide and HCl

(e) reaction with a primary or secondary amine at room temperature to produce an amide and HCl

3 describe the addition-elimination mechanism of acyl chlorides in reactions in 33.3.2(a) – (e)

4 explain the relative ease of hydrolysis of acyl chlorides, alkyl chlorides and halogenoarenes (aryl chlorides)"

 

 

From the prior curriculum, students retained a foundational understanding of organic chemistry, including the nomenclature, structure, preparation, and reactions of aldehydes, ketones, and carboxylic acids.Presently, there is an increased emphasis on practical applications of these concepts in areas such as organic synthesis, biochemistry, and environmental chemistry. The new curriculum also includes updated information and approaches, such as green chemistry and the use of computational tools. Some topics from the former curriculum, such as the chemistry of dyes and pigments, have been removed from the NCC 2023 curriculum.

 

Nitrogen Compounds
(Organic Chemistry)

- Discuss nomenclature, structure and basicity of amines.

- Describe the preparation of amines by alkylation of ammonia to RX and reduction of nitriles, nitro and amide functional groups.

- Discuss reactivity of amines.

- Describe chemistry of amines by alkylation of amines with RX, reactions with aldehydes, ketones, preparation of amides and diazonium salts.

- Describe isomerism in alkyl halides and amines.

- Identify amines in the laboratory by carrying out appropriate tests.

- Perform tests to detect nitrogen in organic compounds.

1 recall the reactions by which amines can be produced:

(a) reaction of a halogenoalkane with NH3 in ethanol heated under pressure

1 recall the reactions by which nitriles can be produced:

(a) reaction of a halogenoalkane with KCN in ethanol and heat

2 recall the reactions by which hydroxynitriles can be produced:

(a) the reaction of aldehydes and ketones with HCN, KCN as catalyst, and heat

3 describe the hydrolysis of nitriles with dilute acid or dilute alkali followed by acidification to produce a

carboxylic acid

 

1 recall the reactions by which amines can be produced:

(a) reaction of a halogenoalkane with NH3 in ethanol heated under pressure

1 recall the reactions by which nitriles can be produced:

(a) reaction of a halogenoalkane with KCN in ethanol and heat

2 recall the reactions by which hydroxynitriles can be produced:

(a) the reaction of aldehydes and ketones with HCN, KCN as catalyst, and heat

3 describe the hydrolysis of nitriles with dilute acid or dilute alkali followed by acidification to produce a

carboxylic acid

Primary and secondary amines

1 recall the reactions (reagents and conditions) by which primary and secondary amines are produced:

(a) reaction of halogenoalkanes with NH3 in ethanol heated under pressure

(b) reaction of halogenoalkanes with primary amines in ethanol, heated in a sealed tube / under pressure

(c) the reduction of amides with LiAl H4

(d) the reduction of nitriles with LiAl H4 or H2 / Ni

2 describe the condensation reaction of ammonia or an amine with an acyl chloride at room temperature to

give an amide

3 describe and explain the basicity of aqueous solutions of amines

Phenylamine and azo compounds

1 describe the preparation of phenylamine via the nitration of benzene to form nitrobenzene followed by

reduction with hot Sn/concentrated HCl , followed by NaOH(aq)

2 describe:

(a) the reaction of phenylamine with Br2(aq) at room temperature

(b) the reaction of phenylamine with HNO2 or NaNO2 and dilute acid below 10 °C to produce the

diazonium salt; further warming of the diazonium salt with H2O to give phenol

3 describe and explain the relative basicities of aqueous ammonia, ethylamine and phenylamine

4 recall the following about azo compounds:

(a) describe the coupling of benzenediazonium chloride with phenol in NaOH(aq) to form an azo compound

(b) identify the azo group

(c) state that azo compounds are often used as dyes

(d) that other azo dyes can be formed via a similar route

Amides

1 recall the reactions (reagents and conditions) by which amides are produced:

(a) the reaction between ammonia and an acyl chloride at room temperature

(b) the reaction between a primary amine and an acyl chloride at room temperature

2 describe the reactions of amides:

(a) hydrolysis with aqueous alkali or aqueous acid

(b) the reduction of the CO group in amides with LiAl H4 to form an amine

3 state and explain why amides are much weaker bases than amines

 

 

Both curriculums cover similar topics related to amines, such as their nomenclature, structure, basicity, preparation methods, and reactions with various functional groups. However, the NCC 2023 curriculum includes some additional topics such as the synthesis and reactions of azo compounds and the chemistry of heterocyclic compounds, while previous ones includes topics such as isomerism in alkyl halides and amines, and the identification of amines and nitrogen in organic compounds through laboratory tests. Additionally, NCC 2023 curriculum provides more detail on certain topics such as the preparation of amines and amides by reduction reactions, and the basicity of aqueous solutions of amines.

 

Polymerization
(Organic Chemistry)

- Describe the formation and uses of PVC and Nylon.

- Discuss the importance of the chemical industries in the economy of Pakistan.

- Describe the raw materials available in Pakistan for various chemical industries.

- Describe the chemical processes of addition and condensation polymerization.

20.1 Addition polymerisation

1 describe addition polymerisation as exemplified by poly(ethene) and poly(chloroethene), PVC

2 deduce the repeat unit of an addition polymer obtained from a given monomer

3 identify the monomer(s) present in a given section of an addition polymer molecule

4 recognise the difficulty of the disposal of poly(alkene)s, i.e. non-biodegradability and harmful combustion

products

35.1 Condensation polymerisation

1 describe the formation of polyesters:

(a) the reaction between a diol and a dicarboxylic acid or dioyl chloride

(b) the reaction of a hydroxycarboxylic acid

2 describe the formation of polyamides:

(a) the reaction between a diamine and a dicarboxylic acid or dioyl chloride

(b) the reaction of an aminocarboxylic acid

(c) the reaction between amino acids

3 deduce the repeat unit of a condensation polymer obtained from a given monomer or pair of monomers

4 identify the monomer(s) present in a given section of a condensation polymer molecule

35.2 Predicting the type of polymerisation

1 predict the type of polymerisation reaction for a given monomer or pair of monomers

2 deduce the type of polymerisation reaction which produces a given section of a polymer molecule

35.3 Degradable polymers

1 recognise that poly(alkenes) are chemically inert and can therefore be difficult to biodegrade

2 recognise that some polymers can be degraded by the action of light

3 recognise that polyesters and polyamides are biodegradable by acidic and alkaline hydrolysis

 

Explain the chemical processes and properties of PVC and Nylon, and the applications of these polymers in the industry.

Discuss the importance of chemical industries in the economy of Pakistan, and describe the raw materials that are available in the country for various chemical industries.

Describe the chemical processes of addition and condensation polymerization and the differences between them.

Examples include

- additon polymers such as poly(ethene) and poly(chloroethene), PVC,

- polyesters (from reactions of diol and dicarboxylic or dioyl acid, and from hydroxycarboxylic acid),

- polyamides (from reactions of a diamine and a dicarboxylic acid or dioyl chloride, of an aminocarboxylic acid, or between amino acids)

student should be able to identify the polymer formed, the monomer prsent in a section of polumer, and classify them as one of the two polymers.

Deduce the repeat unit of a polymer obtained from a given monomer or pair of monomers and identify the monomers present in a given section of a polymer molecule.

Predict the type of polymerization reaction for a given monomer or pair of monomers, and explain the challenges associated with the disposal of non-biodegradable polymers.

- recognise that poly(alkenes) are chemically inert and can therefore be difficult to biodegrade

- recognise that some polymers can be degraded by the action of light

- recognise that polyesters and polyamides are biodegradable by acidic and alkaline hydrolysis

 

Polymers are made up of repeating monomer units which can be manipulated in various ways to give structures with desired properties.

Condensation polymers are formed by the loss of small molecules as functional groups from monomers join.

The topics of PVC and Nylon formation and uses, the importance of chemical industries in Pakistan's economy, and the raw materials available for chemical industries are retained from FSC 2006. The NCC 2023 added new topics on the chemical processes of addition and condensation polymerization, identifying monomers in a polymer molecule, predicting polymerization reactions, and challenges associated with the disposal of non-biodegradable polymers. No topics were removed from FSC 2006.

 

Organic Synthesis
(Organic Chemistry)

 

21.1 Organic synthesis

1 for an organic molecule containing several functional groups:

(a) identify organic functional groups using the reactions in the syllabus

(b) predict properties and reactions

2 devise multi-step synthetic routes for preparing organic molecules using the reactions in the syllabus

3 analyse a given synthetic route in terms of type of reaction and reagents used for each step of it, and

possible by-products

36.1 Organic synthesis

1 for an organic molecule containing several functional groups:

(a) identify organic functional groups using the reactions in the syllabus

(b) predict properties and reactions

2 devise multi-step synthetic routes for preparing organic molecules using the reactions in the syllabus

3 analyse a given synthetic route in terms of type of reaction and reagents used for each step of it, and

possible by-products

Synthetic routes:

The synthesis of an organic compound stems from a readily available starting material via a series of discrete steps. Functional group interconversions are the basis of such synthetic routes.

- Retro-synthesis of organic compounds.

1. Understand the concept of organic synthesis and functional group interconversions.

2. Identify organic functional groups using the reactions in the syllabus.

3. Predict properties and reactions of organic molecules based on functional group presence.

4. Devise multi-step synthetic routes for preparing organic molecules using the reactions in the syllabus.

5. Analyze a given synthetic route in terms of type of reaction and reagents used for each step of it, and possible by-products.

6. Understand the concept of retro-synthesis and its application in organic synthesis.

Conversions with more than four stages will not be assessed in synthetic routes.

- Reaction types can cover any of the reactions covered in 10 and sub-20.1

Usage:

Natural products are compounds isolated from natural sources and include taxol, mescaline and capsaicin.

Organic synthesis is the systematic preparation of a compound from a widely available starting material or the synthesis of a compound via a synthetic route

that often can involve a series of different steps.

This domain was not previously taught in 2006 curriculum The 2023 curriculum places a greater emphasis on organic synthesis, functional group interconversions, and multi-step synthetic routes. Retro-synthesis, a key concept in organic synthesis, is included in the 2023 curriculum . The 2023 curriculum aims to equip students with a deeper understanding of organic reactions and their applications in synthesis.

 

Biochemistry
(Organic Chemistry)

- Explain the basis of classification and structure-Function relationship of Carbohydrates

- Explain the role of various Carbohydrates in health and diseases

- Identify the nutritional importance and their role as energy storage

- Explain the basis of classification and structure-function relationship of proteins

- Describe the role of various proteins in maintaining body functions and their nutritional

importance

- Describe the role of enzyme as biocatalyst and relate this role to various functions such as digestion of food

- Identify factors that affect enzyme activity such as effect of temperature and pH.

- Explain the role of inhibitors of enzyme catalyzed reactions

- Describe the basis of classification and structure-Function relationship of Lipids

- Identify the nutritional and Biological importance of lipids

- Identify the structural components of DNA and RNA

- Recognize the structural differences between DNA polymer (double strand) and RNA (single strand).

- Relate DNA sequences to its function as storage of genetic information

- Relate RNA sequence (transcript) to its role in transfer of information to protein (Translation)

- Identify the sources of minerals such as Iron, Calcium, Phosphorous and Zinc

- Describe the role of Iron, Calcium, Phos horous and Zinc in nutrition

- Explain why animals and humans have large glycogen deposits for sustainable muscular activities. Hibernating animals (polar bear, reptiles and amphibians) accumulate fat to meet energy resources during hibernation

- Identify complex Carbohydrates which provide lubrication to elbow and Knee.

- Describe fibrous proteins from hair and silk

- Explain how Cholesterol and amino acid serve as hormones

- Identify insulin as a protein hormone whose deficiency leads to diabetes mellitus

- Explain the role of minerals in structure and function

- Identify Calcium as a requirement for coagulation

- Identify how milk proteins can be precipitated by lowering the pH using lemon juice

34.4 Amino acids

1 describe the acid / base properties of amino acids and the formation of zwitterions, to include the isoelectric

point

2 describe the formation of amide (peptide) bonds between amino acids to give di- and tripeptides

3 interpret and predict the results of electrophoresis on mixtures of amino acids and dipeptides at varying pHs

(the assembling of the apparatus will not be tested)

 

- Explain the basis of classification and structure-Function relationship of Carbohydrates

- Explain the role of various Carbohydrates in health and diseases

- Identify the nutritional importance and their role as energy storage

- Explain the basis of classification and structure-function relationship of proteins

- Describe the role of various proteins in maintaining body functions and their nutritional

importance

- Describe the role of enzyme as biocatalyst and relate this role to various functions such as digestion of food

- Identify factors that affect enzyme activity such as effect of temperature and pH.

- Explain the role of inhibitors of enzyme catalyzed reactions

- Describe the basis of classification and structure-Function relationship of Lipids

- Identify the nutritional and Biological importance of lipids

- Identify the structural components of DNA and RNA

- Recognize the structural differences between DNA polymer (double strand) and RNA (single strand).

- Relate DNA sequences to its function as storage of genetic information

- Relate RNA sequence (transcript) to its role in transfer of information to protein (Translation)

- Identify the sources of minerals such as Iron, Calcium, Phosphorous and Zinc

- Describe the role of Iron, Calcium, Phos horous and Zinc in nutrition

- Explain why animals and humans have large glycogen deposits for sustainable muscular activities. Hibernating animals (polar bear, reptiles and amphibians) accumulate fat to meet energy resources during hibernation

- Identify complex Carbohydrates which provide lubrication to elbow and Knee.

- Describe fibrous proteins from hair and silk

- Explain how Cholesterol and amino acid serve as hormones

- Identify insulin as a protein hormone whose deficiency leads to diabetes mellitus

- Explain the role of minerals in structure and function

- Identify Calcium as a requirement for coagulation

- Identify how milk proteins can be precipitated by lowering the pH using lemon juice

 

Toxicity and carcinogenic properties of heavy metals are the result of their ability to form coordinated compounds, have various oxidation states and act as

catalysts in the human body.

Metabolic reactions involve a complex interplay between many different components in highly controlled environments.

Proteins are the most diverse of the biopolymers responsible for metabolism and structural integrity of living organisms.

Lipids are a broad group of biomolecules that are largely non-polar and therefore insoluble in water.

Carbohydrates are oxygen-rich biomolecules, which play a central role in metabolic reactions of energy transfer.

Vitamins are organic micronutrients with diverse functions that must be obtained from the diet.

Our increasing knowledge of biochemistry has led to several environmental problems, while also helping to solve others.

Analyses of protein activity and concentration are key areas of biochemical research.

The topics covered in both curriculums are quite similar, with a focus on biochemistry and nutrition. For now, there is more emphasis on understanding the role of nutrients in disease prevention and management, as well as the latest research on nutritional science. Additionally, this includes topics such as food processing, food safety, and food sustainability. However, the topics covered in FSC 2006 are still relevant and provide a good foundation for understanding the basics of biochemistry and nutrition.

 

Combustion Analysis
(Lab and Analysis Skills)

- Compare the classical method of analysis with modern methods.

- Discuss the procedure of combustion analysis.

 

 

Solve simple problems involving combustion analysis

 

Analytical techniques can be used to determine the structure of a compound, analyse the composition of a substance or determine the purity of a compound.

Spectroscopic techniques are used in the structural identification of organic and inorganic compounds.

In spite of discussion on the procedure of combustion analysis which can be made in theory. Students should get themselves familiarize practically on how to solve problems pertaining to it.

 

Mass spectrometry
(Lab and Analysis Skills)

- Explain instrumentation and working of MS.

- Outline the use of MS in determination of relative isotopic masses and isotopic abundance.

- Calculate the average atomic mass of an element from isotopic data.

- Calculate percentage of C, H and 0 from given data and determine empirical and molecular formula.

- Describe how mass spectrometer Is used to determine the relative masses of , atoms and the abundances of isotopes.

- Explain how different instruments help in the study of chemistry.

- Explain how forensic chemists use the MS to identify small amounts of unknown material.

22.2 Mass spectrometry

1 analyse mass spectra in terms of m/e values and isotopic abundances (knowledge of the working of the mass

spectrometer is not required)

2 calculate the relative atomic mass of an element given the relative abundances of its isotopes, or its mass

spectrum

3 deduce the molecular mass of an organic molecule from the molecular ion peak in a mass spectrum

4 suggest the identity of molecules formed by simple fragmentation in a given mass spectrum

5 deduce the number of carbon atoms, n, in a compound using the M +1 peak and the formula

n =100 × (abundance of M +1 ion) / (1.1 × abundance of M + ion)

6 deduce the presence of bromine and chlorine atoms in a compound using the M +2 peak

 

1 analyse mass spectra in terms of m/e values and isotopic abundances (knowledge of the working of the mass

spectrometer is not required)

2 calculate the relative atomic mass of an element given the relative abundances of its isotopes, or its mass

spectrum

3 deduce the molecular mass of an organic molecule from the molecular ion peak in a mass spectrum

4 suggest the identity of molecules formed by simple fragmentation in a given mass spectrum

5 deduce the number of carbon atoms, n, in a compound using the M +1 peak and the formula

n =100 × (abundance of M +1 ion) / (1.1 × abundance of M + ion)

6 deduce the presence of bromine and chlorine atoms in a compound using the M +2 peak

• Radioisotopes are used in nuclear medicine for diagnostics, treatment and research, as tracers in biochemical and pharmaceutical research, and as “chemical clocks” in geological and archaeological dating. • PET (positron emission tomography) scanners give three-dimensional images of tracer concentration in the body, and can be used to detect cancers. Absorption and emission spectra are widely used in astronomy to analyse light from stars. Atomic absorption spectroscopy is a very sensitive means of determining the presence and concentration of metallic elements. Fireworks—emission spectra.

 

Comparatively, gives more comprehensive details on how to structure the formula of a compound using formulas and techniques irrespective of explaining the instrumentation,working and uses of MS

 

Spectrocopy
(Lab and Analysis Skills)

- Define spectroscopy and discuss its applications in analytical Chemistry

- State the regions of electromagnetic spectrum used in IR and UV/vis spectroscopy

- Explain the origin of IR absorption of simple molecules.

- Determine structures of phenol, toluene, acetone and ethanol from its IR spectrum.

- Predict whether a given molecule will absorb in the UV/visible region.

- Predict the color of a transition metal complex from its UV/visible spectrum.

- Define and explain atomic emission and atomic absorption spectrum.

22.1 Infrared spectroscopy

1 analyse an infrared spectrum of a simple molecule to identify functional groups (see the Data section for the

functional groups required)

 

1 analyse an infrared spectrum of a simple molecule to identify functional groups (see the Data section for the functional groups required)

- Determine structures of phenol, toluene, acetone and ethanol from its IR spectrum.

- Predict whether a given molecule will absorb in the UV/visible region.

- Predict the color of a transition metal complex from its UV/visible spectrum.

- Define and explain atomic emission and atomic absorption spectrum.

 

Although spectroscopic characterization techniques form the backbone of structural identification of compounds, typically no one technique results in a full structural identification of a molecule.

The topics covered includes the analysis of infrared and UV/visible spectra, the prediction of the color of transition metal complexes from their UV/visible spectra, and the definition and explanation of atomic emission and absorption spectra. These topics build on the foundations of spectroscopy and analytical chemistry, which were also covered in the former curriculum. However, the revised curriculum places a greater emphasis on the practical applications of spectroscopy in identifying functional groups in simple molecules, predicting the properties of transition metal complexes, and understanding the mechanisms of atomic emission and absorption.

 

NMR
(Lab and Analysis Skills)

- Outline in simple terms the principles of proton NMR spectroscopy.

- Explain how chemical environment of a Proton affects the magnetic field it experiences and hence the absorption of energy at resonance. frequency

- Describe standard scales used in proton NMR.

37.3 Carbon-13 NMR spectroscopy

Learning outcomes

Candidates should be able to:

1 analyse and interpret a carbon-13 NMR spectrum of a simple molecule to deduce:

(a) the different environments of the carbon atoms present

(b) the possible structures for the molecule

2 predict or explain the number of peaks in a carbon-13 NMR spectrum for a given molecule

37.4 Proton (1H) NMR spectroscopy

Learning outcomes

Candidates should be able to:

1 analyse and interpret a proton (1H) NMR spectrum of a simple molecule to deduce:

(a) the different environments of proton present using chemical shift values

(b) the relative numbers of each type of proton present from relative peak areas

(c) the number of equivalent protons on the carbon atom adjacent to the one to which the given proton is

attached from the splitting pattern, using the n + 1 rule (limited to singlet, doublet, triplet, quartet and

multiplet)

(d) the possible structures for the molecule

2 predict the chemical shifts and splitting patterns of the protons in a given molecule

3 describe the use of tetramethylsilane, TMS, as the standard for chemical shift measurements

4 state the need for deuterated solvents, e.g. CDCl 3, when obtaining a proton NMR spectrum

5 describe the identification of O–H and N–H protons by proton exchange using D2O

 

1. Understand and analyze the different environments of carbon atoms present in a simple molecule using a carbon-13 NMR spectrum.

2. Use a carbon-13 NMR spectrum to deduce possible structures of a simple molecule.

3. Predict the number of peaks in a carbon-13 NMR spectrum for a given molecule.

4. Understand and analyze the different environments of protons present in a simple molecule using a proton (1H) NMR spectrum.

5. Use a proton (1H) NMR spectrum to deduce relative numbers of each type of proton present, the number of equivalent protons on the carbon atom adjacent to the one to which the given proton is attached, and possible structures of a simple molecule.

37.3 Carbon-13 NMR spectroscopy

Learning outcomes

Candidates should be able to:

1 analyse and interpret a carbon-13 NMR spectrum of a simple molecule to deduce:

(a) the different environments of the carbon atoms present

(b) the possible structures for the molecule

2 predict or explain the number of peaks in a carbon-13 NMR spectrum for a given molecule

37.4 Proton (1H) NMR spectroscopy

Learning outcomes

Candidates should be able to:

1 analyse and interpret a proton (1H) NMR spectrum of a simple molecule to deduce:

(a) the different environments of proton present using chemical shift values

(b) the relative numbers of each type of proton present from relative peak areas

(c) the number of equivalent protons on the carbon atom adjacent to the one to which the given proton is

attached from the splitting pattern, using the n + 1 rule (limited to singlet, doublet, triplet, quartet and

multiplet)

(d) the possible structures for the molecule

2 predict the chemical shifts and splitting patterns of the protons in a given molecule

3 describe the use of tetramethylsilane, TMS, as the standard for chemical shift measurements

4 state the need for deuterated solvents, e.g. CDCl 3, when obtaining a proton NMR spectrum

5 describe the identification of O–H and N–H protons by proton exchange using D2O

 

Analysis of C-13 NMR spectrum in addition to the elementry proton NMR statistics studied previously.This provides students enriching experience to study spectrum and deduce structures coherently.

 

Chromatography
(Lab and Analysis Skills)

- Explain why forensic chemists must have strong problem-solving skills and a broad background in analytical chemistry.

- Recognize the link between chemical instrumentation and technology

- Make connections between chromatography and MS as used in the analysis of small amounts of unknown materials.

37.1 Thin-layer chromatography

1 describe and understand the terms

(a) stationary phase, for example aluminium oxide (on a solid support)

(b) mobile phase; a polar or non-polar solvent

(c) Rf value

(d) solvent front and baseline

2 interpret Rf values

3 explain the differences in Rf

values in terms of interaction with the stationary phase and of relative solubility

in the mobile phase

37.2 Gas / liquid chromatography

1 describe and understand the terms

(a) stationary phase; a high boiling point non-polar liquid (on a solid support)

(b) mobile phase; an unreactive gas

(c) retention time

2 interpret gas / liquid chromatograms in terms of the percentage composition of a mixture

3 explain retention times in terms of interaction with the stationary phase

 

1. Explain the principles and applications of thin-layer chromatography and gas/liquid chromatography in forensic chemistry and analysis of unknown materials.

2. Identify and interpret Rf values and retention times in chromatograms to determine the composition of a mixture.

3. Understand the importance of selecting the appropriate stationary and mobile phases in chromatography and their impact on the separation of compounds.

4. Describe the use of mass spectrometry in combination with chromatography for identifying and quantifying small amounts of unknown materials in forensic analysis.

 

 

connecting two techniques (chromatography + MS) to pinpoint unknown substance in forensics. In hope, that students recognize and appreciates the addition of TLC ,GLC including few terms like Rf, retention time, to decipher stationary and mobile phases to bring out maximum separation of components which makes elucidation of a mixture efficient.

 

Materials
(Chemistry in Context)

- Discuss types and applications of hair dyes.

- Describe preparation and applications of various cosmetics like nail varnish, nail polish remover and lipsticks.

- Describe types and applications of synthetic adhesives.

- List the safety measures and precautions in process industries.

- List various petrochemicals and their functions.

- Identify risks associated with the manufacturing of chemicals.

- Trace the development and uses of different synthetic fibers.

 

1 Materials science introduction

Understandings:

Materials are classified based on their uses, properties, or bonding and structure.

- The properties of a material based on the degree of covalent, ionic or metallic character in a compound can be deduced from its position on a bonding triangle.

- Composites are mixtures in which materials are composed of two distinct phases, a reinforcing phase that is embedded in a matrix phase

2 Metals and inductively coupled plasma (ICP) spectroscopy

Understandings:

Reduction by coke (carbon), a more reactive metal, or electrolysis are means of obtaining some metals from their ores.

- The relationship between charge and the number of moles of electrons is given by Faraday’s constant, F.

- Alloys are homogeneous mixtures of metals with other metals or non-metals.

- Diamagnetic and paramagnetic compounds differ in electron spin pairing and their behaviour in magnetic fields.

- Trace amounts of metals can be identified and quantified by ionizing them with argon gas plasma in Inductively Coupled Plasma (ICP) Spectroscopy using Mass Spectroscopy ICP-MS and Optical Emission Spectroscopy ICP-OES.

3 Catalysts

Understandings:

Reactants adsorb onto heterogeneous catalysts at active sites and the products desorb.

- Homogeneous catalysts chemically combine with the reactants to form a temporary activated complex or a reaction intermediate.

- Transition metal catalytic properties depend on the adsorption/absorption properties of the metal and the variable oxidation states.

- Zeolites act as selective catalysts because of their cage structure.

- Catalytic particles are nearly always nanoparticles that have large surface areas per unit mass

4 Liquid crystals

Understandings:

Thermoplastics soften when heated and harden when cooled.

- A thermosetting polymer is a prepolymer in a soft solid or viscous state that changes irreversibly into a hardened thermoset by curing.

- Elastomers are flexible and can be deformed under force but will return to nearly their original shape once the stress is released.

- High density polyethene (HDPE) has no branching allowing chains to be packed together.

- Low density polyethene (LDPE) has some branching and is more flexible.

- Plasticizers added to a polymer increase the flexibility by weakening the intermolecular forces between the polymer chains.

- Atom economy is a measure of efficiency applied in green chemistry.

- Isotactic addition polymers have substituents on the same side.

- Atactic addition polymers have the substituents randomly placed

5 Nanotechnology

Understandings:

Molecular self-assembly is the bottom-up assembly of nanoparticles and can occur by selectively attaching molecules to specific surfaces. Self-assembly can also occur spontaneously in solution.

- Possible methods of producing nanotubes are arc discharge, chemical vapour deposition (CVD) and high pressure carbon monoxide (HIPCO).

- Arc discharge involves either vaporizing the surface of one of the carbon electrodes, or discharging an arc through metal electrodes submersed in a hydrocarbon solvent, which forms a small rod-shaped deposit on the anod

6 Environmental impact—plastics

Understandings:

Plastics do not degrade easily because of their strong covalent bonds.

- Burning of polyvinyl chloride releases dioxins, HCl gas and incomplete hydrocarbon combustion products.

- Dioxins contain unsaturated six-member heterocyclic rings with two oxygen atoms, usually in positions 1 and 4.

- Chlorinated dioxins are hormone disrupting, leading to cellular and genetic damage.

- Plastics require more processing to be recycled than other materials.

- Plastics are recycled based on different resin types

7 Superconducting metals and X-ray crystallography

Understandings:

Superconductors are materials that offer no resistance to electric currents below a critical temperature.

- The Meissner effect is the ability of a superconductor to create a mirror image magnetic field of an external field, thus expelling it.

- Resistance in metallic conductors is caused by collisions between electrons and positive ions of the lattice.

- The Bardeen–Cooper–Schrieffer (BCS) theory explains that below the critical temperature electrons in superconductors form Cooper pairs which move freely through the superconductor.

- Type 1 superconductors have sharp transitions to superconductivity whereas Type 2 superconductors have more gradual transitions.

- X-ray diffraction can be used to analyse structures of metallic and ionic compounds.

- Crystal lattices contain simple repeating unit cells.

- Atoms on faces and edges of unit cells are shared.

- The number of nearest neighbours of an atom/ion is its coordination numb

8 Condensation polymers

Understandings:

Condensation polymers require two functional groups on each monomer.

- NH3, HCl and H2O are possible products of condensation reactions.

- Kevlar® is a polyamide with a strong and ordered structure. The hydrogen bonds between O and N can be broken with the use of concentrated sulfuric acid.

9 Environmental impact—heavy metals

Toxic doses of transition metals can disturb the normal oxidation/reduction balance in cells through various mechanisms.

- Some methods of removing heavy metals are precipitation, adsorption, and chelation.

- Polydentate ligands form more stable complexes than similar monodentate ligands due to the chelate effect, which can be explained by considering

1. Understand the properties of different materials and how they can be applied to desired structures.

2. Explain the process of extracting metals from ores and alloying them to achieve desired characteristics.

3. Understand the mechanism of catalysts and how they increase the rate of a reaction while remaining unchanged at the end.

4. Explain the challenges associated with recycling and toxicity of some materials produced through materials science.

5. Explain the use of X-ray crystallography in analyzing structures.

Permeability to moisture should be considered with respect to bonding and simple packing arrangements.

- Consider properties of metals, polymers and ceramics in terms of metallic, covalent, and ionic bonding.

Consider catalytic properties such as selectivity for only the desired product, efficiency, ability to work in mild/severe conditions, environmental impact and impurities.

- The use of carbon nanocatalysts should be covered

Possible implications of nanotechnology include uncertainty as to toxicity levels on a nanoscale, unknown health risks with new materials, concern that human defence systems are not effective against particles on the nanoscale, responsibilities of the industries and governments involved in this research. Conductivity of graphene and fullerenes can be explained in terms of delocalization of electrons. An explanation based on hybridization is not required

Consider green chemistry polymers.

The American Chemical Society (ACS) provides a wealth of resources for teachers, including lesson plans and activities on materials science, such as "Exploring the Properties of Materials" and "Polymers in Daily Life." These resources can be found at https://www.acs.org/content/acs/en/education/resources/highschool/chemmatters/teacher-resources/lesson-plans.html

The National Science Teachers Association (NSTA) also has a collection of resources for teaching materials science, including "Materials Science for High School Students" and "The Science of Superconductors." These resources can be found at https://www.nsta.org/publications/subjects/materials-science/

The Materials Research Society (MRS) offers a range of educational resources, including "Materials Science and Engineering Education" and "Introduction to Materials Science." These resources can be found at https://www.mrs.org/education-careers/k-12-education

The National Science Foundation (NSF) provides information on materials science research and education, including "Materials Science and Engineering: An Introduction" and "Materials Science in the Classroom." These resources can be found at https://www.nsf.gov/materialsresearch/

The Materials Research Science and Engineering Center (MRSEC) at the University of Wisconsin-Madison offers resources for high school teachers including "Materials Science Curriculum" and "Materials Science Research in the Classroom." These resources can be found at https://www.mrsec.wisc.edu/Edetc/highschool.html

Khan Academy provides free online materials and videos on various topics in materials science such as "Introduction to Materials Science" and "Properties of Materials." These resources can be found at https://www.khanacademy.org/science/materials-science

Usage:

Protein synthesis in cells is a form of nanotechnology with ribosomes acting as molecular assemblers.

The products of science and technology can have a negative impact on the environment. Are scientists ethically responsible for the impact of their products?

Does science, economics or politics play the most essential role in research, such as the development of new polymers?"

1. What are the properties of a material and how can they be used to create desired structures?

2. How are metals extracted from ores and alloyed to achieve specific characteristics?

3. What is the role of catalysts in chemical reactions and how do they increase the rate of the reaction?

4. What are the benefits and challenges of recycling materials used in material science?

The FSC 2006 curriculum included topics such as hair dyes, cosmetics, synthetic adhesives, petrochemicals, and synthetic fibers. These topics have now been removed from the curriculum. However, the updated curriculum has introduced new topics such as materials science, metal extraction and alloying, catalysts, recycling, and X-ray crystallography. Some topics from FSC 2006, such as safety measures in process industries, have been retained in the NCC 2023 curriculum.

 

Energy
(Chemistry in Context)

- Interpret difference between petrochemical and chemicals derived from them.

- Describe the fractional distillation and refining of Petroleum

- List the various raw materials for Petrochemical industry.

- Identify the important fractions.

- Describe the basic building block processes in Petrochemical technology.

- Describe the Petrochemical process technology.

- List some major petrochemicals.

 

1. Energy sources

Understanding:

A useful energy source releases energy at a reasonable rate and produces minimal pollution.

- The quality of energy is degraded as heat is transferred to the surroundings. Energy and materials go from a concentrated into a dispersed form. The quantity of the energy available for doing work decreases.

- Renewable energy sources are naturally replenished. Non-renewable energy sources are finite.

- Energy density = energy released from fuel volume of fuel consumed .

- Specific energy = energy released from fuel mass of fuel consumed .

- The effeciency of an energy transfer = useful output energy total input energy x 10

2. Fossil fuels

Understanding:

Fossil fuels were formed by the reduction of biological compounds that contain carbon, hydrogen, nitrogen, sulfur and oxygen.

- Petroleum is a complex mixture of hydrocarbons that can be split into different component parts called fractions by fractional distillation.

- Crude oil needs to be refined before use. The different fractions are separated by a physical process in fractional distillation.

- The tendency of a fuel to auto-ignite, which leads to “knocking” in a car engine, is related to molecular structure and measured by the octane number.

- The performance of hydrocarbons as fuels is improved by the cracking and catalytic reforming reactions.

- Coal gasification and liquefaction are chemical processes that convert coal to gaseous and liquid hydrocarbons.

- A carbon footprint is the total amount of greenhouse gases produced during human activities. It is generally expressed in equivalent tons of carbon dioxide.

3. Nuclear fusion and fission

Understanding:

Nuclear fusion

- Light nuclei can undergo fusion reactions as this increases the binding energy per nucleon.

- Fusion reactions are a promising energy source as the fuel is inexpensive and abundant, and no radioactive waste is produced.

- Absorption spectra are used to analyse the composition of stars. Nuclear fission

- Heavy nuclei can undergo fission reactions as this increases the binding energy per nucleon.

- 235U undergoes a fission chain reaction: U235 92 + n10 → U236 92 → X + Y + neutrons.

- The critical mass is the mass of fuel needed for the reaction to be self- sustaining.

- 239Pu, used as a fuel in “breeder reactors”, is produced from 238U by neutron capture.

- Radioactive waste may contain isotopes with long and short half-lives.

- Half-life is the time it takes for half the number of atoms to decay.

4. Solar energy

Understanding:

Light can be absorbed by chlorophyll and other pigments with a conjugated electronic structure.

- Photosynthesis converts light energy into chemical energy: 6CO2 + 6H2O ◊ C6H12O6 + 6O2

- Fermentation of glucose produces ethanol which can be used as a biofuel: C6H12O6 ◊ 2C2H5OH + 2CO2

- Energy content of vegetable oils is similar to that of diesel fuel but they are not used in internal combustion engines as they are too viscous.

- Transesterification between an ester and an alcohol with a strong acid or base catalyst produces a different ester: RCOOR1 + R2OH ◊ RCOOR2 + R1OH

- In the transesterification process, involving a reaction with an alcohol in the presence of a strong acid or base, the triglyceride vegetable oils are converted to a mixture mainly comprising of alkyl esters and glycerol, but with some fatty acids.

- Transesterification with ethanol or methanol produces oils with lower viscosity that can be used in diesel engines.

5. Environmental impact—global warming

Understanding:

Greenhouse gases allow the passage of incoming solar short wavelength radiation but absorb the longer wavelength radiation from the Earth. Some of the absorbed radiation is re-radiated back to Earth.

- There is a heterogeneous equilibrium between concentration of atmospheric carbon dioxide and aqueous carbon dioxide in the oceans.

- Greenhouse gases absorb IR radiation as there is a change in dipole moment as the bonds in the molecule stretch and bend.

- Particulates such as smoke and dust cause global dimming as they reflect sunlight, as do clouds.

6. Electrochemistry, rechargeable batteries and fuel cells

Understanding:

An electrochemical cell has internal resistance due to the finite time it takes for ions to diffuse. The maximum current of a cell is limited by its internal resistance.

- The voltage of a battery depends primarily on the nature of the materials used while the total work that can be obtained from it depends on their quantity.

- In a primary cell the electrochemical reaction is not reversible. Rechargeable cells involve redox reactions that can be reversed using electricity.

- A fuel cell can be used to convert chemical energy, contained in a fuel that is consumed, directly to electrical energy.

- Microbial fuel cells (MFCs) are a possible sustainable energy source using different carbohydrates or substrates present in waste waters as the fuel.

- The Nernst equation, E = E0- (RTnF) ln Q, can be used to calculate the potential of a half-cell in an electrochemical cell, under non-standard conditions.

- The electrodes in a concentration cell are the same but the concentration of the electrolyte solutions at the cathode and anode are different.

7. Nuclear fusion and nuclear fission

Understanding:

Nuclear fusion:

- The mass defect (∆m) is the difference between the mass of the nucleus and the sum of the masses of its individual nucleons.

- The nuclear binding energy (ΔE) is the energy required to separate a nucleus into protons and neutrons. Nuclear fission:

- The energy produced in a fission reaction can be calculated from the mass difference between the products and reactants using the Einstein mass–energy equivalence relationship

𝐸𝐸 = 𝑚𝑚𝑚𝑚

2.

- The different isotopes of uranium in uranium hexafluoride can be separated, using diffusion or centrifugation causing fuel enrichment.

- The effusion rate of a gas is inversely proportional to the square root of the molar mass (Graham’s Law).

- Radioactive decay is kinetically a first order process with the half-life related to the decay constant by the equation

𝜆𝜆 = ln 2 𝑡𝑡

12 .

- The dangers of nuclear energy are due to the ionizing nature of the radiation it produces which leads to the production of oxygen free radicals such as superoxide (O2-), and hydroxyl (HO·). These free radicals can initiate chain reactions that can damage DNA and enzymes in living cells.

8. Photovoltaic cells and dye-sensitized solar cells (DSSC)

Understanding:

Molecules with longer conjugated systems absorb light of longer wavelength.

- The electrical conductivity of a semiconductor increases with an increase in temperature whereas the conductivity of metals decreases.

- The conductivity of silicon can be increased by doping to produce n-type and p- type semiconductors.

- Solar energy can be converted to electricity in a photovoltaic cell.

- DSSCs imitate the way in which plants harness solar energy. Electrons are "injected" from an excited molecule directly into the TiO2 semiconductor.

- The use of nanoparticles coated with light-absorbing dye increases the effective surface area and allows more light over a wider range of the visible spectrum to be absorbed.

1. Understand the difference between petrochemical and chemicals derived from them, and identify the various raw materials for the petrochemical industry.

2. Explain the process of fractional distillation and refining of petroleum, and identify the important fractions.

3. Describe the basic building block processes in petrochemical technology, and explain the petrochemical process technology.

4. List some major petrochemicals, and understand the importance of petrochemicals in the modern world.

5. Understand the energy density and specific energy of different energy sources, and explain the efficiency of energy transfer.

6. Understand the formation, properties, and uses of fossil fuels, and explain the importance of fossil fuels in the modern world.

7. Understand the mechanism and importance of nuclear fusion and fission, and explain the importance of nuclear energy in the modern world.

8. Understand the importance and mechanism of solar energy, and explain the importance of renewable energy in the modern world.

9. Understand the environmental impact of energy consumption, particularly in relation to global warming, and be able to explain the importance of reducing carbon footprint and moving towards sustainable energy sources.

10. Apply their knowledge of energy sources and their properties to critically evaluate the advantages and disadvantages of different energy sources and make informed decisions about energy consumption.

1. Energy sources

Understanding:

A useful energy source releases energy at a reasonable rate and produces minimal pollution.

- The quality of energy is degraded as heat is transferred to the surroundings. Energy and materials go from a concentrated into a dispersed form. The quantity of the energy available for doing work decreases.

- Renewable energy sources are naturally replenished. Non-renewable energy sources are finite.

- Energy density = energy released from fuel volume of fuel consumed .

- Specific energy = energy released from fuel mass of fuel consumed .

- The effeciency of an energy transfer = useful output energy total input energy x 10

2. Fossil fuels

Understanding:

Fossil fuels were formed by the reduction of biological compounds that contain carbon, hydrogen, nitrogen, sulfur and oxygen.

- Petroleum is a complex mixture of hydrocarbons that can be split into different component parts called fractions by fractional distillation.

- Crude oil needs to be refined before use. The different fractions are separated by a physical process in fractional distillation.

- The tendency of a fuel to auto-ignite, which leads to “knocking” in a car engine, is related to molecular structure and measured by the octane number.

- The performance of hydrocarbons as fuels is improved by the cracking and catalytic reforming reactions.

- Coal gasification and liquefaction are chemical processes that convert coal to gaseous and liquid hydrocarbons.

- A carbon footprint is the total amount of greenhouse gases produced during human activities. It is generally expressed in equivalent tons of carbon dioxide.

3. Nuclear fusion and fission

Understanding:

Nuclear fusion

- Light nuclei can undergo fusion reactions as this increases the binding energy per nucleon.

- Fusion reactions are a promising energy source as the fuel is inexpensive and abundant, and no radioactive waste is produced.

- Absorption spectra are used to analyse the composition of stars. Nuclear fission

- Heavy nuclei can undergo fission reactions as this increases the binding energy per nucleon.

- 235U undergoes a fission chain reaction: U235 92 + n10 → U236 92 → X + Y + neutrons.

- The critical mass is the mass of fuel needed for the reaction to be self- sustaining.

- 239Pu, used as a fuel in “breeder reactors”, is produced from 238U by neutron capture.

- Radioactive waste may contain isotopes with long and short half-lives.

- Half-life is the time it takes for half the number of atoms to decay.

4. Solar energy

Understanding:

Light can be absorbed by chlorophyll and other pigments with a conjugated electronic structure.

- Photosynthesis converts light energy into chemical energy: 6CO2 + 6H2O ◊ C6H12O6 + 6O2

- Fermentation of glucose produces ethanol which can be used as a biofuel: C6H12O6 ◊ 2C2H5OH + 2CO2

- Energy content of vegetable oils is similar to that of diesel fuel but they are not used in internal combustion engines as they are too viscous.

- Transesterification between an ester and an alcohol with a strong acid or base catalyst produces a different ester: RCOOR1 + R2OH ◊ RCOOR2 + R1OH

- In the transesterification process, involving a reaction with an alcohol in the presence of a strong acid or base, the triglyceride vegetable oils are converted to a mixture mainly comprising of alkyl esters and glycerol, but with some fatty acids.

- Transesterification with ethanol or methanol produces oils with lower viscosity that can be used in diesel engines.

5. Environmental impact—global warming

Understanding:

Greenhouse gases allow the passage of incoming solar short wavelength radiation but absorb the longer wavelength radiation from the Earth. Some of the absorbed radiation is re-radiated back to Earth.

- There is a heterogeneous equilibrium between concentration of atmospheric carbon dioxide and aqueous carbon dioxide in the oceans.

- Greenhouse gases absorb IR radiation as there is a change in dipole moment as the bonds in the molecule stretch and bend.

- Particulates such as smoke and dust cause global dimming as they reflect sunlight, as do clouds.

6. Electrochemistry, rechargeable batteries and fuel cells

Understanding:

An electrochemical cell has internal resistance due to the finite time it takes for ions to diffuse. The maximum current of a cell is limited by its internal resistance.

- The voltage of a battery depends primarily on the nature of the materials used while the total work that can be obtained from it depends on their quantity.

- In a primary cell the electrochemical reaction is not reversible. Rechargeable cells involve redox reactions that can be reversed using electricity.

- A fuel cell can be used to convert chemical energy, contained in a fuel that is consumed, directly to electrical energy.

- Microbial fuel cells (MFCs) are a possible sustainable energy source using different carbohydrates or substrates present in waste waters as the fuel.

- The Nernst equation, E = E0- (RTnF) ln Q, can be used to calculate the potential of a half-cell in an electrochemical cell, under non-standard conditions.

- The electrodes in a concentration cell are the same but the concentration of the electrolyte solutions at the cathode and anode are different.

7. Nuclear fusion and nuclear fission

Understanding:

Nuclear fusion:

- The mass defect (∆m) is the difference between the mass of the nucleus and the sum of the masses of its individual nucleons.

- The nuclear binding energy (ΔE) is the energy required to separate a nucleus into protons and neutrons. Nuclear fission:

- The energy produced in a fission reaction can be calculated from the mass difference between the products and reactants using the Einstein mass–energy equivalence relationship

𝐸𝐸 = 𝑚𝑚𝑚𝑚

2.

- The different isotopes of uranium in uranium hexafluoride can be separated, using diffusion or centrifugation causing fuel enrichment.

- The effusion rate of a gas is inversely proportional to the square root of the molar mass (Graham’s Law).

- Radioactive decay is kinetically a first order process with the half-life related to the decay constant by the equation

𝜆𝜆 = ln 2 𝑡𝑡

12 .

- The dangers of nuclear energy are due to the ionizing nature of the radiation it produces which leads to the production of oxygen free radicals such as superoxide (O2-), and hydroxyl (HO·). These free radicals can initiate chain reactions that can damage DNA and enzymes in living cells.

8. Photovoltaic cells and dye-sensitized solar cells (DSSC)

Understanding:

Molecules with longer conjugated systems absorb light of longer wavelength.

- The electrical conductivity of a semiconductor increases with an increase in temperature whereas the conductivity of metals decreases.

- The conductivity of silicon can be increased by doping to produce n-type and p- type semiconductors.

- Solar energy can be converted to electricity in a photovoltaic cell.

- DSSCs imitate the way in which plants harness solar energy. Electrons are ""injected"" from an excited molecule directly into the TiO2 semiconductor.

- The use of nanoparticles coated with light-absorbing dye increases the effective surface area and allows more light over a wider range of the visible spectrum to be absorbed.

2 The cost of production and availability (reserves) of fossil fuels and their impact on the environment should be considered.

3 Students are not expected to recall specific fission reactions.

- The workings of a nuclear power plant are not required.

- Safety and risk issues include: health, problems associated with nuclear waste and core meltdown, and the possibility that nuclear fuels may be used in nuclear weapons.

- The equations,

𝑁𝑁 = 𝑁𝑁0𝑒𝑒𝜆𝜆𝜆𝜆 and 𝑡𝑡1 2 = ln2 𝜆𝜆

are given in section 1 of the data booklet.

4 Only a conjugated system with alternating double bonds needs to be covered.

5 Greenhouse gases to be considered are CH4, H2O and CO2.

6 A battery should be considered as a portable electrochemical source made up of one or more voltaic (galvanic) cells connected in series.

- The Nernst equation is given in the data booklet in section 1.

- Hydrogen and methanol should be considered as fuels for fuel cells. The operation of the cells under acid and alkaline conditions should be considered. Students should be familiar with proton-exchange membrane (PEM) fuel cells.

- The Geobacter species of bacteria, for example, can be used in some cells to oxidize the ethanoate ions (CH3COO-) under anaerobic conditions.

- The lead–acid storage battery, the nickel–cadmium (NiCad) battery and the lithium–ion battery should be considered.

- Students should be familiar with the anode and cathode half-equations and uses of the different cells.

7 Students are not expected to recall specific fission reactions.

- The workings of a nuclear power plant are not required.

- Safety and risk issues include: health, problems associated with nuclear waste, and the possibility that nuclear fuels may be used in nuclear weapons.

- Graham’s law of effusion is given in the data booklet in section 1.

- Decay relationships are given in the data booklet in section 1.

- A binding energy curve is given in the data booklet in section 3

8 The relative conductivity of metals and semiconductors should be related to ionization energies.

- Only a simple treatment of the operation of the cells is needed. In p-type semiconductors, electron holes in the crystal are created by introducing a small percentage of a group 3 element. In n-type semiconductors inclusion of a group 5 element provides extra electrons.

- In a photovoltaic cell the light is absorbed and the charges separated in the silicon semiconductor. The processes of absorption and charge separation are separated in a dye-sensitized solar cell.

- Specific redox and electrode reactions in the newer Grätzel DSSC should be covered. An example is the reduction of I2/I3─ ions to I─

1 Societies are completely dependent on energy resources. The quantity of energy is conserved in any conversion but the quality is degraded.

2 The energy of fossil fuels originates from solar energy which has been stored by chemical processes over time. These abundant resources are non-renewable but provide large amounts of energy due to the nature of chemical bonds in hydrocarbons.

3 The fusion of hydrogen nuclei in the sun is the source of much of the energy needed for life on Earth. There are many technological challenges in replicating this process on Earth but it would offer a rich source of energy. Fission involves the splitting of a large unstable nucleus into smaller stable nuclei.

4 Visible light can be absorbed by molecules that have a conjugated structure with an extended system of alternating single and multiple bonds. Solar energy can be converted to chemical energy in photosynthesis.

5 Gases in the atmosphere that are produced by human activities are changing the climate as they are upsetting the balance between radiation entering and leaving the atmosphere.

6 Chemical energy from redox reactions can be used as a portable source of electrical energy.

7 Large quantities of energy can be obtained from small quantities of matter.

8 When solar energy is converted to electrical energy the light must be absorbed and charges must be separated. In a photovoltaic cell both of these processes occur in the silicon semiconductor, whereas these processes occur in separate locations in a dye-sensitized solar cell (DSSC).

The NCC 2023 curriculum for chemistry retains the core concepts and knowledge related to petrochemical production from FSC 2006. However, it places a greater emphasis on sustainability and renewable energy sources, with additional topics on energy sources, efficiency of energy transfer, and reducing carbon footprint. The specific changes and updates in the curriculum are aimed at providing a more comprehensive and up-to-date view of chemistry and addressing environmental concerns for the students.

 

Medicine
(Chemistry in Context)

 

 

1 Pharmaceutical products and drug action

Understandings:

In animal studies, the therapeutic index is the lethal dose of a drug for 50% of the population (LD50) divided by the minimum effective dose for 50% of the population (ED50).

- In humans, the therapeutic index is the toxic dose of a drug for 50% of the population (TD50) divided by the minimum effective dose for 50% of the population (ED50).

- The therapeutic window is the range of dosages between the minimum amounts of the drug that produce the desired effect and a medically unacceptable adverse effect.

- Dosage, tolerance, addiction and side effects are considerations of drug administration.

- Bioavailability is the fraction of the administered dosage that reaches the target part of the human body.

- The main steps in the development of synthetic drugs include identifying the need and structure, synthesis, yield and extraction.

- Drug–receptor interactions are based on the structure of the drug and the site of activity.

2 Aspirin and penicillin

Understandings:

Aspirin:

- Mild analgesics function by intercepting the pain stimulus at the source, often by interfering with the production of substances that cause pain, swelling or fever.

- Aspirin is prepared from salicylic acid.

- Aspirin can be used as an anticoagulant, in prevention of the recurrence of heart attacks and strokes and as a prophylactic. Penicillin:

- Penicillins are antibiotics produced by fungi.

- A beta-lactam ring is a part of the core structure of penicillins.

- Some antibiotics work by preventing cross-linking of the bacterial cell walls.

- Modifying the side-chain results in penicillins that are more resistant to the penicillinase enzyme.

3 Opiates

Understandings:

The ability of a drug to cross the blood–brain barrier depends on its chemical structure and solubility in water and lipids.

- Opiates are natural narcotic analgesics that are derived from the opium poppy.

- Morphine and codeine are used as strong analgesics. Strong analgesics work by temporarily bonding to receptor sites in the brain, preventing the transmission of pain impulses without depressing the central nervous system.

- Medical use and addictive properties of opiate compounds are related to the presence of opioid receptors in the brain.

4 pH regulation of the stomach

Understandings:

Non-specific reactions, such as the use of antacids, are those that work to reduce the excess stomach acid.

- Active metabolites are the active forms of a drug after it has been processed by the body.

5 Antiviral medications

Understandings:

Viruses lack a cell structure and so are more difficult to target with drugs than bacteria.

- Antiviral drugs may work by altering the cell’s genetic material so that the virus cannot use it to multiply. Alternatively, they may prevent the viruses from multiplying by blocking enzyme activity within the host cell.

6 Environmental impact of some medications

Understandings:

High-level waste (HLW) is waste that gives off large amounts of ionizing radiation for a long time.

- Low-level waste (LLW) is waste that gives off small amounts of ionizing radiation for a short time.

- Antibiotic resistance occurs when micro-organisms become resistant to antibacterials

Additional Higher Level Topics:

7 Taxol—a chiral auxiliary case study

Understandings:

Taxol is a drug that is commonly used to treat several different forms of cancer.

- Taxol naturally occurs in yew trees but is now commonly synthetically produced.

- A chiral auxiliary is an optically active substance that is temporarily incorporated into an organic synthesis so that it can be carried out asymmetrically with the selective formation of a single enantiomer.

8 Nuclear medicine

Understandings:

Alpha, beta, gamma, proton, neutron and positron emissions are all used for medical treatment.

- Magnetic resonance imaging (MRI) is an application of NMR technology.

- Radiotherapy can be internal and/or external.

- Targeted Alpha Therapy (TAT) and Boron Neutron Capture Therapy (BNCT) are two methods which are used in cancer treatment.

9 Drug detection and analysis

Understandings:

Organic structures can be analysed and identified through the use of infrared spectroscopy, mass spectroscopy and proton NMR.

- The presence of alcohol in a sample of breath can be detected through the use of either a redox reaction or a fuel cell type of breathalyser

1. Understand the concept of therapeutic index and therapeutic window in relation to drug administration and be able to calculate the same

2. Understand the mechanism of action and uses of Aspirin and Penicillin and explain the chemical structure of the same

3. Understand the mechanism of action of Opiates and the concept of opioid receptors in the brain

4. Understand the pH regulation of stomach and the concept of non-specific reactions and active metabolites

5. Understand the challenges in treating viral infections with drugs and the concept of Antiviral medications.

1 For ethical and economic reasons, animal and human tests of drugs (for LD50/ED50 and TD50/ED50 respectively) should be kept to a minimum.

2 Students should be aware of the ability of acidic (carboxylic) and basic (amino) groups to form ionic salts, for example soluble aspirin.

- Structures of aspirin and penicillin are available in the data booklet in section 37.

3 Structures of morphine, codeine and diamorphine can be found in the data booklet in section 37.

4 Antacid compounds should include calcium hydroxide, magnesium hydroxide, aluminium hydroxide, sodium carbonate and sodium bicarbonate.

- Structures for ranitidine and omeprazole can be found in the data booklet in section 37.

Understand the concept of therapeutic index and therapeutic window in relation to drug administration and the importance of considering factors such as dosage, tolerance, addiction and side effects. (Sources:

https://www.ncbi.nlm.nih.gov/books/NBK482457/ , https://www.sciencedirect.com/topics/medicine-and-dentistry/therapeutic-index

)

Understand the mechanism of action and uses of drugs such as aspirin and penicillin, including their chemical structures and effects on the body. (Sources: https://www.ncbi.nlm.nih.gov/books/NBK482458/,

https://www.ncbi.nlm.nih.gov/books/NBK11162/

)

Understand the mechanism of action and effects of opiates, including their interaction with opioid receptors in the brain and the potential for addiction. (Sources:

https://www.ncbi.nlm.nih.gov/books/NBK482459/ , https://www.ncbi.nlm.nih.gov/books/NBK310/

)

Understand the pH regulation of the stomach and the mechanism of action of antacids. (Sources: https://www.ncbi.nlm.nih.gov/books/NBK279/,

https://www.ncbi.nlm.nih.gov/books/NBK443365/

)

Understand the mechanism of action and challenges in the development of antiviral medications, as well as the potential environmental impact of some medications. (Sources: https://www.ncbi.nlm.nih.gov/books/NBK482460/,

https://www.ncbi.nlm.nih.gov/books/NBK5369/ )

1. What is the therapeutic index and how is it determined in humans and animals?

2. How do drug-receptor interactions work and what factors affect them?

3. What are the main steps in the development of synthetic drugs and what considerations are important in this process?

4. How do opiates work to relieve pain and what are the potential addictive properties of these drugs?

5. What are the challenges associated with antiviral medication and how do these drugs work to combat viral infections?

6. How do pH regulation and antacids work in the stomach and what are the potential side effects of these medications?

7. What are some of the medical applications of nuclear medicine and what are the risks associated with these treatments?

8. How are drugs detected and analyzed in the body and what are the important considerations for this process?

Medicine was not separately discussed in 2006 curriculm and have been added here with the aim to understand its context with special emphasis on treating viral infections

 

Agriculture
(Chemistry in Context)

 

 

 

1. understand the chemical composition and function of different types of fertilizers, including their role in providing essential nutrients to crops and the impact of their application on soil health.

2. identify the different types of pesticides used in agriculture and describe their mode of action, including the potential benefits and risks associated with their use.

3. understand the chemical reactions that occur when acid rain falls on crops and soil and the effects it has on crop growth, including nutrient uptake and crop yield.

4. understand the basics of genetic engineering and how it is used in agriculture, including the development of genetically modified crops and the potential benefits and risks associated with their use.

5. understand how changes in temperature, precipitation, and extreme weather events can affect crop growth and yield, including the potential for crop failures and food shortages, as well as the potential to develop new crop varieties that are more resilient to changing climate conditions.

students should have a basic understanding of the chemical components of soil, including pH levels, nutrient content, and the presence of minerals and trace elements. They should also understand how these factors affect the growth and development of crops.

Teachers should guide students in learning about the different types of fertilizers available, including chemical fertilizers and organic options such as compost and manure. They should also teach students how to apply these fertilizers correctly and in appropriate amounts to ensure optimal plant growth.

students should be educated on the common pests and diseases that affect crops in Pakistan, as well as effective ways to control and prevent them. This includes understanding the life cycles of pests and the impact of environmental factors on their populations.

Teachers should provide students with information on various irrigation methods and the importance of water management in agriculture. They should also help students understand the environmental impact of irrigation methods and the potential for water conservation and sustainability.

students should be introduced to the latest technology used in agriculture, such as precision farming, GIS mapping, and remote sensing. This can help students understand how technology is used in modern agriculture and how it can improve crop yields and reduce environmental impact.

Following sources may be helpful

https://www.extension.purdue.edu/extmedia/AY/AY-342-W.pdf

https://www.nrcs.usda.gov/wps/portal/nrcs/detail/soils/edu/agronomy/fertilizer/?cid=nrcs142p2_054234

http://www.fao.org/3/ca2899en/CA2899EN.pdf

https://www.irrigation.org/

https://www.agriculture.com/machinery/technology

https://www.ipni.net/

1. How do chemical processes and compounds play a role in crop growth and development?

2. How can chemistry be used to improve soil fertility and crop yield?

3. How do agricultural chemicals, such as pesticides and fertilizers, impact the environment and human health?

 

 

Industry
(Chemistry in Context)

- Discuss the importance of the chemical industries in the economy of Pakistan.

- Describe the raw materials available in Pakistan for various chemical industries.

- Describe the chemical processes of addition and condensation polymerization.

- Interpret difference between petrochemical and chemicals derived from them.

- Describe the fractional distillation and refining of Petroleum List the various raw materials for Petrochemical industry.

- Identify the important fractions.

- Describe the basic building block processes in Petrochemical technology.

- Describe the Petrochemical process technology.(List some major petrochemicals.

- Discuss types and applications of hair dyes.

- Describe the formation and uses of PVC and Nylon.

- Describe preparation and applications of various cosmetics like nail varnish, nail polish remover and lipsticks.

- Describe types and applications of synthetic adhesives.

- List the safety measures and precautions in process industries.

- List various petrochemicals and their functions.

 

 

1. Understand the importance and significance of industrial chemistry in various industries such as manufacturing, energy, healthcare, and environmental protection.

2. Describe the chemical processes involved in industrial production, including addition and condensation polymerization, and the properties and uses of resulting materials.

3. Identify the raw materials and resources used in industrial chemistry, including those readily available in the context of Pakistan.

4. Explain the applications of industrial chemistry in industries such as petrochemical, cosmetics, cement, food production and more.

5. Elaborate on the safety measures and precautions necessary in industrial chemical processes and facilities.

 

1. How do chemical industries contribute to the economy of Pakistan and what are the raw materials available in the country for these industries?

2. The process of addition and condensation polymerization and how it is used in the production of various materials?

3. How does the petrochemical industry in Pakistan operate, including the process of fractional distillation and refining of petroleum, and the various raw materials and major petrochemicals used?

4. How is cement and food production in Pakistan impacted by the use of chemical processes and what are the safety measures and precautions taken in these industries?

Previously, the content was considerably subjected onto petrochemical and cosmetic industries. Now, the students get holistic understanding of significance of various industries, exploit their resources, chemical process involved,properties and usage of finished products in addition to the preventive measure adopted within Pakistan.

 

 

Grades 11-12 Experimentation SLOs

 

 

2006 National Curiculum

CAIE A Levels Curriculum 2025-2027

IB DP Curriculum 2016

NCC 2023

Guidance on SLOs

(Eleboration on the extent and depth of study and assessment expectations)

Rationale

 

 

 

 

Successful collection of data and observations

Students should be able to:

• set up apparatus

• follow instructions given in the form of written instructions or diagrams

• use apparatus to collect an appropriate quantity of data

• make observations, including subtle differences in colour, solubility or quantity of materials

• make measurements using pipettes, burettes, measuring cylinders, thermometers and other common

laboratory apparatus; Students should record burette readings to the nearest 0.05 cm

and, when using a

thermometer calibrated at 1 °C intervals, temperature readings should be recorded to the nearest 0.5 °C.

Quality of measurements or observations

Students should be able to:

• make accurate and consistent measurements and observations, including achieving concordant titres (two

titres within 0.10 cm3 of each other) and precise colour descriptions.

Decisions relating to measurements or observations

Students should be able to:

• decide how many tests or observations to perform

• make measurements that span a range and have a distribution appropriate to the experiment

• identify where repeated readings or observations are appropriate

• replicate readings or observations as necessary, including where an anomaly is suspected

• identify where confirmatory tests are appropriate and the nature of such tests

• select reagents to distinguish between given ions.

Recording data and observations

Students should be able to:

• present numerical data, values or observations in a single table of results with headings and units that conform

to accepted scientific conventions, i.e. volume / cm3, volume (cm3) or volume in cm3

• record raw readings of a quantity to the same degree of precision, e.g. if one measurement of mass in a

collection of raw data is given as 0.06 g, then all the masses in that collection should be given to the nearest

0.01 g. The degree of precision recorded should be compatible with the measuring instrument used, e.g. a

measuring cylinder calibrated at 1.0 cm3 should be read to the nearest 0.5 cm3

• record observations to the same level of detail, e.g. observations of qualitative variables such as colour should

be recorded in simple language such as ‘blue’ or ‘yellow’. Where fine discrimination is required, terms such as ‘pale’ or ‘dark’ should be used, and comparisons made such as ‘darker brown than at three minutes’ or ‘paler

green than with 0.2 mol dm−3'.

Display of calculation and reasoning

Students should be able to:

• show working in calculations and key steps in reasoning: where calculations are carried out, all the stages in the

calculation should be recorded, so that credit can be given for correctly displayed working

• use the correct number of significant figures for calculated quantities (this should be the same as or one

more than the smallest number of significant figures in the provided or experimentally determined data).

For example, if titre volume is measured to four significant figures, e.g. 23.45 cm3, then the calculated molar concentrations from this should be given to four significant figures, e.g. 1.305 mol dm−3 or 0.9876 mol dm. However, if the concentration of one of the reactants is given to three significant figures, then the calculated

concentration could be given to three or four significant figures. For example, if the concentration of alkali

in an acid-base titration is given as 0.100 mol dm−3, then the concentration of the acid may be shown as 0.1305 mol dm−3 or 0.131 mol dm

Data Layout

Students should be able to:

• present data in a single table of results−3

• draw an appropriate table in advance of taking readings or making observations, so that they do not have to

copy their results

• record all data in the table

• use the appropriate presentation method to produce a clear presentation of the data, e.g. graph lines and graph

points should be drawn using a sharp pencil

• plot appropriate variables on appropriate, clearly labelled x- and y-axes (the same convention for axis labels

should be used as for table headings)

• choose scales for graph axes that allow the graph to be read easily, such as 1, 2 or 5 units to a 20 mm square;

the data points should occupy at least half of the graph grid in both x- and y-directions

• plot all points using a cross × or circled dot ʘ to an appropriate accuracy

• draw straight lines or smooth curves of best fit to show the trend of a graph; a line of best fit should show

an even distribution of points on either side of the line along its entire length, anomalous points should be

identified.

−3

Analysis, conclusions and evaluation

Interpretation of data or observations

Students should be able to:

• describe the patterns and trends shown by data in tables and graphs

• describe and summarise the key points of a set of observations

• calculate quantities from data, or calculate the mean from repeated values, or make other appropriate

calculations

• find an unknown value by using coordinates, a point of intersection or intercepts on a graph

• determine the gradient of a straight-line graph, using two points that are more than half of the length of the

axes apart

• extrapolate the line of a graph.

Calculations may involve mean, percentage, percentage gain or loss, rate of reaction, concentration, molar mass

and volume of gases or other appropriate calculations.

Drawing conclusions

Students should be able to:

• draw conclusions from an experiment, giving an outline description of the main features of the data,

considering whether experimental data support a given hypothesis, and making further predictions

• draw conclusions from interpretations of observations, data and calculated values

• make scientific explanations of data, observations and conclusions that they have described.

Students may be required to prove or disprove given hypotheses, using deductions from the data, observations

or calculated values. Simple scientific explanations form a part of such conclusions and therefore Students will

be expected to refer to knowledge and understanding gained in the theory part of the course in order to provide

explanations of their practical conclusions.

Identifying sources of error and suggesting improvements

Students should be able to:

• evaluate the effectiveness of control variables

• comment on errors intrinsic in measuring devices, e.g. a thermometer consistently reading 1 °C above actual

temperature, or in experiments where limitations of the method introduce errors, e.g. heat loss when trying to

assess enthalpy change

• show an understanding of the distinction between systematic errors, e.g. zero error in a balance hence the

reason for using the same balance for all weighings within an experiment, and random errors, e.g. change

in room temperature when investigating the effect of changing concentration on the rate of a reaction.

(A statement of ‘human errors’ is not acceptable; though there are occasionally errors arising in the observer’s

ability to observe, e.g. in the disappearing cross experiment, which would be a random error.)

• identify the most significant sources of error in an experiment

• state the uncertainty in a quantitative measurement and express such uncertainty in a measurement as

an actual or percentage error. For the purpose of this syllabus, the maximum uncertainty in a quantitative

measurement is half the difference between the closest calibrations, e.g. for a thermometer calibrated at 1 °C

the maximum uncertainty is ± 0.5 °C, therefore the maximum percentage error in a temperature change of

14.0 °C = ((2 × 0.5)/14.0) × 100 = 7.14%

• suggest realistic modifications to an experimental arrangement that will improve the accuracy of the

experiment or the observations that can be made

• suggest ways in which to extend the investigation to answer a new question.

Some of the possible tasks include:

• a hands-on laboratory investigation

• using a spreadsheet for analysis and modelling

• extracting data from a database and analysing it graphically

• producing a hybrid of spreadsheet/database work with a traditional hands-on investigation

• using a simulation provided it is interactive and open-ended.

Some tasks may consist of relevant and appropriate qualitative work combined with quantitative work.

The practical programme should be flexible enough to allow a wide variety of practical activities to be carried out.

These could include:

• short labs or projects extending over several weeks

• computer simulations

• using databases for secondary data

• developing and using models

• data-gathering exercises such as questionnaires, user trials and surveys

• data-analysis exercises

• fieldwork.

The 2006 National Curriculum largely requires students to reproduce experiments they have already done in their classes; this does not help with learning higher order thinking skills.

The experimental SLOs have taken inspiration from the IB Diploma and CAIE O/A level curriculum. In Grade 11 the emphasis is put on building sophistication of students' practical data collection and analysis skills. In assessment, students will be tasks to do that involve experimentally testing a hypothesis, and then critiquing the procedure after having collected and analysed the data In Grade 12, the emphasis is put on developing the ability to design from scatch experimental procedures and for analysis; the assessment will be based on students being able to articulate and design well thought out designs for testing hypotheses.

Titration experiments

 

 

 

Students are expected to understand how to correctly set up a burette in order to carry out titrations.

Students are expected to carry out a rough titration first.

Students are expected to carry out titrations until concordant results are obtained.

A knowledge of the following titrations will be expected:

• acid-alkali titration (this could be weak or strong acid and weak or strong alkali) and the use of indicators listed

on page 61

• potassium manganate(VII) titration with hydrogen peroxide, iron(II) ions, nitrite ions or ethanedioic acid or its

salts

• sodium thiosulfate and iodine titrations.

This list is not exhaustive, and simple titrations involving other reagents may also be set and additional information

provided where necessary.

 

Rates experiments

 

 

 

Students are expected to be able to follow instructions to mix reagents and record the time for an observation

to occur

an example of such an experiment is the time taken on mixing solutions of sodium thiosulfate and an acid for the print on a piece of paper to be obscured by the precipitate produced.

 

Gravimetric experiments

 

 

 

Students are expected to be able to heat a solid in a crucible on a pipe-clay triangle and record any mass change

an example of such an experiment is the determination of the water of hydration of a hydrated salt by evaporation of the water and calculation of the change in mass.

 

Thermometric experiments

 

 

 

Students are expected to be able to accurately use and take readings from thermometers

an example of such

an experiment is the determination of the enthalpy change of reaction by recording of temperature changes and

subsequent calculation of enthalpy changes and use of Hess’s law.

 

Gas volume experiments

 

 

 

Students are expected to be able to set up apparatus for a gas collection over water method

an example of such an experiment is the determination of the composition of a solid from the volume of carbon dioxide produced on reaction of a carbonate with an acid.

 

Qualitative analysis

 

 

 

Students should understand the appropriate methods to be used when carrying out qualitative analysis tests:

• to treat all unknown materials with caution

• to use an appropriate quantity of the material under test

• to add only the specified amount

• to work safely, e.g. to use a test-tube holder when heating a solid in a hard-glass test-tube

• to record all observations, even when this is ‘no change’ or ‘remains a colourless solution’

• to use excess alkali where a precipitate is produced on addition of NaOH(aq) or NH3 (aq) to determine its solubility

• to identify a gas whose formation is shown by effervescence.

Students may also be required to carry out the following organic analysis tests and/or interpret the positive test

result to identify the functional group present:

• the production of an orange/red precipitate with Fehling’s reagent to indicate the presence of the aldehyde functional group

• the production of a silver mirror/black precipitate with Tollens’ reagent to indicate the presence of the aldehyde functional grou

• the production of a yellow precipitate with alkaline aqueous iodine to indicate the presence of the CH3CO or CH3CH(OH) group

• the change in colour of acidified potassium manganate(VII) from purple to colourless to indicate the presence of a compound that can be oxidised.

Students should be familiar with carrying out qualitative analysis reactions for all elements, compounds and

ions listed in the Qualitative analysis notes, and with using these notes to make conclusions about the unknown

substance being tested. However, the substances to be investigated may contain ions not included in these

notes. In such cases Students will not be expected to identify the ions but only to record observations and draw

conclusions of a general nature where possible.

 

Recommendations

 

 

 

It is recommended that following practicals be conducted

Formula of magnesium oxide

Determining the Mr of an unknown gas

Acid-base titrations

A green acid-base practical

Analysis of aspirin tablets

CaCO3 in egg shells

Enthalpy changes

Reaction rates

Rate-dependent factors

Determining Ea for a reaction

Iodination of propanone

Titrations with a pH meter

Redox titration with KMnO4

Voltaic cells

3-D molecular modelling

It is suggested that following topics be included in practicals to be conducted

Common chemical reactions (mixing these combinations; acid-alkali, acid-metal, acid-metal carbonate, precipitation, concentrated acid-salt)

Determining trends and properties Elements & oxides of the third period

Determining trends and properties The halogens

Boiling points of mixtures

Polarity of molecules

Le Chatelier's principle

Determining Kc for an esterification reaction

Redox reactions of vanadium

Chlorine in swimming pools

Analysis of Cu(II) ions in solution

Percentage of copper in brass

Electrolytic cells

Reactions of organic compounds

Hydrolysis of halogenoalkanes

Preparation of 1,3-dinitrobenzene

Determination of an organic structure

Preparation of nylon 6,6

Determination of vitamin C content

Hydrolysis of starch

Preparation & purification of aspirin

Where possible, usage of ICT resources such as data loggers, graphing software (Desmos, Geogebra), simulations (pHet), spreadsheets for data processing and manipulation should be used, introduced and encouraged.