Stability constants, denoted as $ K_{stab} $, are fundamental parameters in coordination chemistry that quantify the stability of complexes formed between metal ions and ligands. Understanding stability constants is crucial for students of AS & A Level Chemistry (9701) as they provide insights into the formation, behavior, and properties of coordination compounds involving transition elements. This knowledge is essential for applications ranging from industrial processes to biological systems.

Key Concepts

Definition of Stability Constants

Stability constants, also known as formation constants, are equilibrium constants that describe the formation of a complex ion from its constituent metal ion and ligands. For a general complex formation reaction:

$$M^{n+} + xL \leftrightarrow ML_x^{(n-x)+}$$

where:

M$^{n+}$: Metal ion with charge $ n+ $
L: Ligand
ML$_x$$^{(n-x)+}$: Complex ion with $ x $ ligands

The stability constant ($ K_{stab} $) for this reaction is expressed as:

$$K_{stab} = \frac{[ML_x^{(n-x)+}]}{[M^{n+}][L]^x}$$

Here, [ML$_x$$^{(n-x)+}$], [M$^{n+}$], and [L] represent the equilibrium concentrations of the complex ion, free metal ion, and free ligand, respectively.

Types of Stability Constants

Stability constants can be categorized based on the number of ligands involved in the complex formation:

1:1 Complexes: Involving one metal ion and one ligand. Example:
1:2 Complexes: Involving one metal ion and two ligands. Example:

Factors Affecting Stability Constants

The magnitude of $ K_{stab} $ is influenced by several factors:

Charge of the Metal Ion: Higher positive charges on the metal ion generally increase the stability of the complex due to stronger electrostatic attractions between the metal ion and the ligands.
Nature of the Ligand: Ligands with higher donor atom electronegativity and those capable of forming multiple bonds with the metal ion tend to form more stable complexes.
Ionic Radius: Smaller metal ions can approach ligands more closely, leading to stronger bonding interactions and higher stability constants.
Coordination Number: The number of ligand donor atoms attached to the metal ion affects the overall stability; typically, intermediate coordination numbers lead to higher stability.
Solvent Effects: The solvent can stabilize or destabilize the complex through solvation of the metal ion and ligands, influencing the observed stability constant.

Calculating Stability Constants

To determine $ K_{stab} $, one must measure the equilibrium concentrations of the metal ion, ligand, and complex ion in solution. Consider the reaction:

$$M^{n+} + L \leftrightarrow ML^{(n)+}$$

If the initial concentrations are:

[M$^{n+}$]$_0$ = 1.0 M
[L]$_0$ = 1.0 M
[ML$_x$$^{(n-x)+}$]$_0$ = 0 M

At equilibrium:

[M$^{n+}$] = 1.0 - x
[L] = 1.0 - x
[ML$_x$$^{(n-x)+}$] = x

Thus, the stability constant is:

$$K_{stab} = \frac{x}{(1.0 - x)(1.0 - x)}$$

Applications of Stability Constants

Stability constants are utilized in various applications, including:

Analytical Chemistry: Determining the concentration of metal ions in solutions through complexation methods.
Biochemistry: Understanding metal ion interactions with biological molecules like proteins and enzymes.
Environmental Chemistry: Assessing the mobility and bioavailability of metal pollutants in ecosystems.

Advanced Concepts

Thermodynamic Basis of Stability Constants

The stability constant is fundamentally a reflection of the thermodynamic favorability of the complex formation. It is related to the change in Gibbs free energy ($ \Delta G $) of the reaction:

$$\Delta G = -RT \ln K_{stab}$$

Where:

R: Gas constant (8.314 J/mol.K)
T: Temperature in Kelvin
K_{stab}: Stability constant

A larger $ K_{stab} $ corresponds to a more negative $ \Delta G $, indicating a more spontaneous and stable complex formation.

Mathematical Derivation of Stepwise Stability Constants

For complexes that form in multiple steps, stepwise stability constants can be defined. Consider the sequential formation of a complex:

$$M^{n+} + L \leftrightarrow ML^{(n)+} \quad (K_1)$$ $$ML^{(n)+} + L \leftrightarrow ML_2^{(n-2)+} \quad (K_2)$$

The overall stability constant for the formation of ML$_2$ is the product of the stepwise constants:

$$K_{stab} = K_1 \times K_2$$

Chelate Effect and Its Impact on Stability Constants

The chelate effect refers to the enhanced stability of complexes that form ring structures through multi-dentate ligands compared to those with equivalent monodentate ligands. Multi-dentate ligands can form multiple bonds with the metal ion, increasing the overall stability constant:

$$L_n + M^{n+} \leftrightarrow ML_n^{(n-x)+}$$

Where $ L_n $ is a multi-dentate ligand with $ n $ donor sites. The formation of rings restricts the number of possible dissociation pathways, leading to higher $ K_{stab} $ values.

Interdisciplinary Connections: Stability Constants in Biochemistry

In biochemistry, stability constants play a pivotal role in understanding metal ion interactions with biomolecules. For example, the binding of iron to hemoglobin involves stability constants that dictate the affinity and release of oxygen molecules. Similarly, the stability of enzyme-substrate complexes is governed by these constants, influencing metabolic pathways and reaction rates.

Complex Problem-Solving Involving Stability Constants

Consider a solution containing 0.1 M of $ Cu^{2+} $ and 0.2 M of ammonia ($ NH_3 $). The formation of the complex $ [Cu(NH_3)_4]^{2+} $ has a stability constant $ K_{stab} = 1.0 \times 10^7 $. Determine the concentration of the complex formed at equilibrium.

Starting with the equilibrium expression:

$$K_{stab} = \frac{[Cu(NH_3)_4^{2+}]}{[Cu^{2+}][NH_3]^4}$$

Let $ x $ be the concentration of the complex formed. Then:

$$1.0 \times 10^7 = \frac{x}{(0.1 - x)(0.2 - 4x)^4}$$

Assuming $ x $ is much smaller than 0.1 and 0.2:

$$1.0 \times 10^7 \approx \frac{x}{(0.1)(0.2)^4}$$ $$x \approx 1.0 \times 10^7 \times 0.1 \times (0.2)^4$$ $$x \approx 1.0 \times 10^7 \times 0.1 \times 0.0016$$ $$x \approx 1600 \, M$$

However, this result is not feasible as concentrations cannot exceed the initial amounts. This indicates that the assumption $ x $ is much smaller is invalid. A more accurate approach requires solving the quartic equation numerically, which may involve iterative methods or approximation techniques.

Comparison Table

Aspect	Tropic	Stability Constants (Kstab)
Definition	Measures the strength of the bond in a complex.	Quantifies the equilibrium between metal ions and ligands in complex formation.
Dependence	Depends on metal ion and ligand properties.	Influenced by factors like charge, size, and ligand denticity.
Impact on Stability	Higher values indicate stronger, more stable complexes.	Higher Kstab signifies greater favorability of complex formation.
Applications	Used in crystallography and material science.	Essential in analytical chemistry, biochemistry, and environmental science.

Summary and Key Takeaways

Stability constants ($ K_{stab} $) quantify the strength of metal-ligand complexes.
Higher $ K_{stab} $ values indicate more stable and favorable complex formations.
Factors such as metal ion charge, ligand nature, and ionic radius significantly influence stability constants.
Advanced understanding involves thermodynamic principles, the chelate effect, and interdisciplinary applications.
Accurate calculation of $ K_{stab} $ is essential for predicting complex behavior in various chemical contexts.

Examiner Tip

Tips

Use the ICE Table Method: Organize initial concentrations, changes, and equilibrium concentrations in an ICE table to simplify complex calculations.

Memorize Key Stability Constants: Familiarize yourself with common stability constants of prevalent metal-ligand complexes to quickly reference during exams.

Understand the Chelate Effect: Remember that multi-dentate ligands form more stable complexes than mono-dentate ligands, which can help in predicting stability.

Practice Problem-Solving: Regularly solve numerical problems involving stability constants to build confidence and accuracy.

Did You Know

Stability constants play a crucial role in the design of pharmaceuticals. For instance, the effectiveness of certain chemotherapy drugs relies on their ability to form stable complexes with metal ions in the body, enhancing their delivery to cancer cells. Additionally, the Nobel Prize-winning work of Alfred Werner on coordination compounds laid the foundation for understanding stability constants, which revolutionized inorganic chemistry.

Common Mistakes

Incorrect Assumption of Completeness: Students often assume that all metal ions form complexes completely, leading to errors in calculating equilibrium concentrations.
Incorrect: Assuming [ML] = initial [L]
Correct: Account for the degree of complexation using equilibrium expressions.

Misinterpretation of Exponents: Misplacing exponents in the stability constant expression can lead to incorrect values.
Incorrect: $K_{stab} = \frac{[ML]}{[M][L]}$ without considering stoichiometry.
Correct: Include the correct stoichiometric coefficients, e.g., $K_{stab} = \frac{[ML_2]}{[M][L]^2}$.

Neglecting Activity Coefficients: Ignoring activity coefficients in solutions with high ionic strength can result in inaccurate stability constants.

FAQ

What is a stability constant?

A stability constant ($K_{stab}$) is an equilibrium constant that quantifies the stability of a complex ion formed between a metal ion and ligands in solution.

How is the stability constant calculated?

It is calculated using the concentrations of the complex ion, free metal ion, and free ligands at equilibrium, following the expression $K_{stab} = \frac{[ML_x]}{[M][L]^x}$.

Why are stability constants important in biochemistry?

They help in understanding how metal ions interact with biological molecules, influencing processes like oxygen transport and enzyme activity.

What factors increase the stability constant of a complex?

Higher metal ion charge, smaller ionic radius, stronger donor atoms in ligands, and the presence of multi-dentate ligands all contribute to higher stability constants.

Can stability constants vary with temperature?

Yes, stability constants can change with temperature as temperature affects the equilibrium position and the energetics of complex formation.

What is the difference between stepwise and overall stability constants?

Stepwise stability constants ($K_1$, $K_2$, etc.) refer to the formation of each ligand in successive steps, while the overall stability constant ($K_{stab}$) is the product of all stepwise constants for the complete complex formation.

1. Group 17

1.1 Physical Properties of the Group 17 Elements

1.1.1 Trends in Bond Strength of Halogen Molecules

1.1.2 Volatility Explained by Intermolecular Forces

1.1.3 Colours and Volatility of Chlorine, Bromine and Iodine

1.2 The Chemical Properties of the Halogen Elements and the Hydrogen Halides

1.2.1 Relative Reactivity of Halogen Elements as Oxidising Agents

1.2.2 Reactions of Halogens with Hydrogen

1.2.3 Thermal Stability of Hydrogen Halides

1.3 Some Reactions of the Halide Ions

1.3.1 Relative Reactivity of Halide Ions as Reducing Agents

1.3.2 Reactions of Halide Ions with Aqueous Silver Ions and Ammonia

1.3.3 Reactions of Halide Ions with Concentrated Sulfuric Acid

1.4 The Reactions of Chlorine

1.4.1 Reaction of Chlorine with Aqueous Sodium Hydroxide (Cold and Hot)

1.4.2 Use of Chlorine in Water Purification

2. Electrochemistry

2.1 Redox Processes: Electron Transfer and Changes in Oxidation Number

2.1.1 Calculating Oxidation Numbers

2.1.2 Using Oxidation Numbers to Balance Equations

2.1.3 Redox, Oxidation, Reduction, Disproportionation: Definitions

2.1.4 Oxidising and Reducing Agents: Definitions

2.1.5 Indicating Oxidation Numbers Using Roman Numerals

3. Chemical Energetics

3.1 Enthalpies of Solution and Hydration

3.1.1 Calculations Involving Solution and Hydration Energies

3.1.2 Effect of Ionic Charge and Radius on Hydration Enthalpy

3.1.3 Definition and Use of Enthalpy Change of Solution and Hydration

3.1.4 Construction and Use of Energy Cycles for Solution and Hydration

3.2 Entropy Change ΔS

3.2.1 Definition of Entropy

3.2.2 Prediction and Explanation of Entropy Changes

3.2.3 Calculation of Entropy Change for Reactions

3.3 Gibbs Free Energy Change ΔG

3.3.1 Gibbs Free Energy Equation

3.3.2 Calculations Using ΔG = ΔH – TΔS

3.3.3 Prediction of Reaction Feasibility

3.3.4 Effect of Temperature on Reaction Feasibility

3.4 Lattice Energy and Born–Haber Cycles

3.4.1 Definition and Use of Enthalpy Change of Atomisation and Lattice Energy

3.4.2 First Electron Affinity and Trends in Electron Affinity

3.4.3 Construction and Use of Born–Haber Cycles

3.4.4 Calculations Involving Born–Haber Cycles

3.4.5 Effect of Ionic Charge and Radius on Lattice Energy

4. Group 2

4.1 Magnesium to Barium, and Their Compounds

4.1.1 Variation in Solubility and Enthalpy Change of Solution of Hydroxides and Sulfates

4.1.2 Trend in Thermal Stability of Group 2 Nitrates and Carbonates

4.1.3 Effect of Ionic Radius on Polarisation of Large Anions

5. Atoms, Molecules and Stoichiometry

5.1 Formulas

5.1.1 Using Appropriate State Symbols

5.1.2 Empirical and Molecular Formulas: Definitions

5.1.3 Anhydrous, Hydrated and Water of Crystallisation

5.1.4 Calculation of Empirical and Molecular Formulas

5.1.5 Writing Ionic Compound Formulas from Ionic Charges and Oxidation Numbers

5.1.6 Predicting Ionic Charges from Periodic Table Positions

5.1.7 Common Ions: Names and Formulas

5.1.8 Writing and Balancing Chemical and Ionic Equations

5.2 Reacting Masses and Volumes (of Solutions and Gases)

5.2.1 Calculations Involving Reacting Masses and Percentage Yield

5.2.2 Calculations Involving Volumes of Gases

5.2.3 Calculations Involving Volumes and Concentrations of Solutions

5.2.4 Limiting Reagent and Excess Reagent Calculations

5.2.5 Deducing Stoichiometric Relationships from Calculations

5.3 Relative Masses of Atoms and Molecules

5.3.1 Unified Atomic Mass Unit

5.3.2 Relative Atomic Mass (Ar)

5.3.3 Relative Isotopic Mass

5.3.4 Relative Molecular Mass (Mr)

5.3.5 Relative Formula Mass

5.4 The Mole and the Avogadro Constant

5.4.1 Definition and Use of the Mole

5.4.2 Understanding the Avogadro Constant

6. Nitrogen Compounds

6.1 Primary and Secondary Amines

6.1.1 Production of Primary and Secondary Amines from Halogenoalkanes

6.1.2 Reduction of Amides and Nitriles to Amines

6.1.3 Reaction of Amines with Acyl Chlorides

6.1.4 Basicity of Aqueous Solutions of Amines

6.2 Phenylamine and Azo Compounds

6.2.1 Preparation of Phenylamine from Nitrobenzene

6.2.2 Reactions of Phenylamine with Bromine and Diazotisation

6.2.3 Relative Basicities of Ammonia, Ethylamine and Phenylamine

6.2.4 Formation and Use of Azo Compounds as Dyes

6.3 Amides

6.3.1 Production of Amides from Acyl Chlorides and Ammonia/Amines

6.3.2 Hydrolysis of Amides with Acids or Alkalis

6.3.3 Reduction of Amides to Amines

6.3.4 Weak Basic Nature of Amides

6.4 Amino Acids

6.4.1 Acid–Base Properties and Zwitterions

6.4.2 Formation of Peptide Bonds Between Amino Acids

6.4.3 Electrophoresis of Amino Acids and Peptides

7. Halogen Compounds

7.1 Halogenoalkanes

7.1.1 Elimination Reactions of Halogenoalkanes

7.1.2 SN1 and SN2 Mechanisms of Nucleophilic Substitution

7.1.3 Reactivity of Halogenoalkanes and Bond Strength

7.1.4 Production of Halogenoalkanes: Free-Radical Substitution and Electrophilic Addition

7.1.5 Classification of Halogenoalkanes: Primary, Secondary, Tertiary

7.1.6 Nucleophilic Substitution Reactions of Halogenoalkanes

7.1.7 Identification of Halogens by Reaction with Silver Nitrate

8. Chemistry of Transition Elements

8.1 General Physical and Chemical Properties of the First Row of Transition Elements

8.1.1 Titanium to Copper: Definition of Transition Elements

8.1.2 Titanium to Copper: Sketching 3dxy and 3dz² Orbitals

8.1.3 Titanium to Copper: General Properties: Variable Oxidation States, Catalytic Behavior, Formation of

8.1.4 Titanium to Copper: Explanation of Variable Oxidation States and Catalysis

8.2 General Characteristic Chemical Properties of the First Set of Transition Elements

8.2.1 Formation of Complexes and Ligand Types

8.2.2 Shapes and Coordination Numbers of Complexes

8.2.3 Ligand Exchange Reactions

8.2.4 Redox Reactions Involving Transition Elements

8.2.5 Redox Calculations Using E° Values

8.3 Colour of Complexes

8.3.1 Degenerate and Non-Degenerate d Orbitals

8.3.2 Splitting of d Orbitals in Octahedral and Tetrahedral Complexes

8.3.3 Explanation of Colour Based on Light Absorption

8.3.4 Effect of Ligand Types on Colour

8.4 Stereoisomerism in Transition Element Complexes

8.4.1 Geometrical (cis-trans) and Optical Isomerism in Complexes

8.4.2 Deduction of Overall Polarity of Complexes

8.5 Stability Constants

8.5.1 Definition and Expression of Stability Constants (Kstab)

8.5.2 Ligand Exchange and Kstab Values

9. Hydroxy Compounds

9.1 Alcohols

9.1.1 Production of Alcohols: Steam Addition, Oxidation, Reduction and Hydrolysis

9.1.2 Reactions of Alcohols: Combustion, Substitution, Oxidation, Dehydration, Esterification

9.1.3 Classification of Alcohols: Primary, Secondary, Tertiary

9.1.4 Tri-iodomethane Test for CH₃CH(OH)– Group

9.1.5 Acidity of Alcohols Compared to Water

10. Equilibria

10.1 Chemical Equilibria: Reversible Reactions and Dynamic Equilibrium

10.1.1 Reversible Reactions and Dynamic Equilibrium

10.1.2 Le Chatelier’s Principle

10.1.3 Effect of Temperature, Concentration, Pressure and Catalysts on Equilibrium

10.1.4 Equilibrium Constant Expressions (Kc)

10.1.5 Mole Fraction and Partial Pressure

10.1.6 Equilibrium Constant Expressions (Kp)

10.1.7 Calculations Using Kc and Kp

10.1.8 Changes Affecting the Value of Equilibrium Constants

10.1.9 Industrial Applications: Haber Process and Contact Process

10.2 Brønsted–Lowry Theory of Acids and Bases

10.2.1 Common Acids and Alkalis: Names and Formulas

10.2.2 Brønsted–Lowry Theory: Acids and Bases

10.2.3 Strong and Weak Acids and Bases

10.2.4 pH Scale and Differences Between Strong and Weak Acids

10.2.5 Neutralisation and Salt Formation

10.2.6 Acid–Base Titration Curves

10.2.7 Choosing Suitable Indicators for Titrations

11. Nitrogen and Sulfur

11.1 Nitrogen and Sulfur

11.1.1 Lack of Reactivity of Nitrogen Explained by Bond Strength and Polarity

11.1.2 Basicity of Ammonia and Structure of the Ammonium Ion

11.1.3 Displacement of Ammonia from Ammonium Salts

11.1.4 Natural and Man-made Sources of Nitrogen Oxides

11.1.5 Role of Nitrogen Oxides in Photochemical Smog Formation

11.1.6 Role of Nitrogen Oxides in Acid Rain Formation

12. Carbonyl Compounds

12.1 Aldehydes and Ketones

12.1.1 Production of Aldehydes and Ketones by Oxidation of Alcohols

12.1.2 Reduction of Aldehydes and Ketones

12.1.3 Reaction of Aldehydes and Ketones with Hydrogen Cyanide

12.1.4 Mechanism of Nucleophilic Addition Reactions

12.1.5 Detection of Carbonyl Compounds Using 2, 4-DNPH

12.1.6 Distinguishing Aldehydes from Ketones Using Fehling’s and Tollens’ Tests

12.1.7 Tri-iodomethane Test for CH₃CO– Group

13. Chemical Bonding

13.1 Electronegativity and Bonding

13.1.1 Definition of Electronegativity

13.1.2 Factors Affecting Electronegativity

13.1.3 Trends in Electronegativity Across a Period and Down a Group

13.1.4 Predicting Bond Type Using Electronegativity Differences

13.2 Ionic Bonding

13.2.1 Definition of Ionic Bonding

13.2.2 Ionic Bonding Examples: Sodium Chloride, Magnesium Oxide, Calcium Fluoride

13.3 Metallic Bonding

13.3.1 Definition of Metallic Bonding

13.4 Covalent Bonding and Coordinate (Dative Covalent) Bonding

13.4.1 Definition of Covalent Bonding

13.4.2 Covalent Bonding in Common Molecules (H₂, O₂, N₂, Cl₂, HCl, CO₂, NH₃, CH₄, C₂H₆, C₂H₄)

13.4.3 Expanded Octet in Period 3 Elements (SO₂, PCl₅, SF₆)

13.4.4 Definition of Coordinate (Dative Covalent) Bonding

13.4.5 Coordinate Bonding in NH₄⁺ and Al₂Cl₆

13.4.6 Sigma (σ) and Pi (π) Bonds: Orbital Overlap

13.4.7 Hybridisation: sp, sp² and sp³ Orbitals

13.4.8 Bond Energy and Bond Length Concepts

13.5 Shapes of Molecules

13.5.1 Shapes and Bond Angles Using VSEPR Theory (Examples: BF₃, CO₂, CH₄, NH₃, H₂O, SF₆, PF₅)

13.5.2 Predicting Shapes and Bond Angles of Similar Molecules and Ions

13.6 Intermolecular Forces, Electronegativity and Bond Properties

13.6.1 Hydrogen Bonding: Examples in Water and Ammonia

13.6.2 Anomalous Properties of Water Due to Hydrogen Bonding

13.6.3 Bond Polarity and Dipole Moments

13.6.4 Van der Waals' Forces: Instantaneous Dipole–Induced Dipole and Permanent Dipole–Dipole Forces

13.6.5 Comparison of Bond Strengths: Ionic, Covalent, Metallic vs. Intermolecular Forces

13.7 Dot-and-Cross Diagrams

13.7.1 Using Dot-and-Cross Diagrams to Represent Ionic, Covalent and Coordinate Bonding

14. Electrochemistry

14.1 Electrolysis

14.1.1 Prediction of Products During Electrolysis

14.1.2 Faraday Constant and Relationship with Avogadro Constant and Electron Charge

14.1.3 Calculations Involving Charge, Mass and Volume in Electrolysis

14.1.4 Determination of Avogadro Constant by Electrolysis

14.2 Standard Electrode Potentials and the Nernst Equation

14.2.1 Definition of Standard Electrode Potential and Standard Cell Potential

14.2.2 Description of the Standard Hydrogen Electrode

14.2.3 Measurement of Standard Electrode Potentials

14.2.4 Calculation of Standard Cell Potentials

14.2.5 Deduction of Electron Flow and Reaction Feasibility Using Electrode Potentials

14.2.6 Reactivity Predictions Using E° Values

14.2.7 Construction of Redox Equations from Half-Equations

14.2.8 Variation of Electrode Potential with Concentration

14.2.9 Using the Nernst Equation for Electrode Potentials

14.2.10 Relationship Between ΔG° and E°cell

15. Introduction to Organic Chemistry

15.1 Formulas, Functional Groups and the Naming of Organic Compounds

15.1.1 Definition of Hydrocarbons

15.2 Formulas

15.2.1 Functional Groups and Physical/Chemical Properties

15.2.2 Interpretation of General, Structural, Displayed and Skeletal Formulas

15.2.3 Systematic Nomenclature of Simple Aliphatic Organic Compounds

15.2.4 Deducing Molecular and Empirical Formulas

15.3 Characteristic Organic Reactions

15.3.1 Terminology of Organic Reactions: Homologous Series, Saturated/Unsaturated, Fission, Radicals, Nucle

15.3.2 Types of Organic Reactions: Addition, Substitution, Elimination, Hydrolysis, Condensation, Oxidation

15.3.3 Mechanisms of Organic Reactions: Free Radical Substitution, Electrophilic Addition, Nucleophilic Sub

15.4 Shapes of Organic Molecules; σ and π Bonds

15.4.1 Shapes and Bond Angles in Organic Molecules

15.4.2 Sigma (σ) and Pi (π) Bond Arrangements

15.4.3 Planarity in Organic Molecules

15.5 Isomerism: Structural Isomerism and Stereoisomerism

15.5.1 Types of Structural Isomerism: Chain, Positional, Functional Group

15.5.2 Types of Stereoisomerism: Geometrical (cis-trans) and Optical

15.5.3 Geometrical Isomerism in Alkenes

15.5.4 Chirality and Optical Isomerism

15.5.5 Identifying Chiral Centres and Isomer Types

16. Genetic technology

16.1 Genetically modified organisms in agriculture

16.1.1 GM crops and animals and their implications

16.2 Primary Amines

16.2.1 Production of Primary Amines by Reaction of Halogenoalkanes with Ammonia

16.3 Nitriles and Hydroxynitriles

16.3.1 Production of Nitriles by Reaction of Halogenoalkanes with KCN

16.3.2 Production of Hydroxynitriles by Reaction of Aldehydes and Ketones with HCN

16.3.3 Hydrolysis of Nitriles to Produce Carboxylic Acids

17. Atomic Structure

17.1 Particles in the Atom and Atomic Radius

17.1.1 Structure of the Atom: Nucleus and Electron Shells

17.1.2 Subatomic Particles: Protons, Neutrons, Electrons

17.1.3 Atomic Number, Proton Number, Mass Number, Nucleon Number

17.1.4 Mass and Charge Distribution in Atoms

17.1.5 Deflection of Charged Particles in Electric Fields

17.1.6 Calculations Involving Protons, Neutrons, Electrons

17.1.7 Trends in Atomic and Ionic Radii Across Periods and Down Groups

17.2 Isotopes

17.2.1 Definition and Notation of Isotopes

17.2.2 Chemical Properties of Isotopes

17.2.3 Physical Properties of Isotopes: Mass and Density Differences

17.3 Electrons

17.3.1 Energy Levels, and Atomic Orbitals: Shells, Sub-shells and Orbitals

17.3.2 Energy Levels,and Atomic Orbitals: Principal Quantum Number (n) and Ground State Configurations

17.3.3 Energy Levels and Atomic Orbitals: Order of Sub-shell Energy Levels (s, p, d)

17.3.4 Energy Levels and Atomic Orbitals: Full and Shorthand Electronic Configurations

17.3.5 Energy Levels and Atomic Orbitals: Electron Box Diagrams

17.3.6 Energy Levels and Atomic Orbitals: Shapes of s and p Orbitals

17.3.7 Energy Levels and Atomic Orbitals: Free Radicals and Unpaired Electrons

17.4 Ionisation Energy

17.4.1 First Ionisation Energy: Definition and Equations

17.4.2 Trends in Ionisation Energy Across Periods and Down Groups

17.4.3 Successive Ionisation Energies and Electron Shells

17.4.4 Factors Affecting Ionisation Energy: Nuclear Charge, Radius, Shielding

17.4.5 Deducing Electronic Configurations Using Ionisation Energy Data

17.4.6 Position of Elements Using Ionisation Energy Patterns

18. Polymerisation (Condensation Polymers)

18.1 Condensation Polymerisation

18.1.1 Formation of Polyesters from Diols and Dicarboxylic Acids or Dioyl Chlorides

18.1.2 Formation of Polyamides from Diamines and Dicarboxylic Acids or Dioyl Chlorides

18.1.3 Formation of Polymers from Amino Acids

18.2 Predicting the Type of Polymerisation

18.2.1 Identifying the Type of Polymerisation from Monomers or Polymer Structure

18.3 Degradable Polymers

18.3.1 Chemical Inertness and Disposal of Poly(alkenes)

18.3.2 Biodegradability of Polyesters and Polyamides

18.3.3 Degradation by Light

19. Carboxylic Acids and Derivatives

19.1 Carboxylic Acids

19.1.1 Production of Carboxylic Acids: Oxidation, Hydrolysis of Nitriles and Esters

19.1.2 Reactions of Carboxylic Acids with Metals, Alkalis and Carbonates

19.1.3 Esterification Reactions with Alcohols

19.1.4 Reduction of Carboxylic Acids

19.1.5 Relative Acidity of Carboxylic Acids, Phenols and Alcohols

19.1.6 Acidity of Chlorine-Substituted Carboxylic Acids

19.2 Esters

19.2.1 Production of Esters by Condensation of Carboxylic Acids and Alcohols

19.2.2 Hydrolysis of Esters with Acid or Alkali

20. Introduction to A Level Organic Chemistry

20.1 Formulas, Functional Groups and the Naming of Organic Compounds

20.1.1 Functional Groups in Aromatic and Extended Organic Molecules

20.1.2 Use of Structural, Displayed and Skeletal Formulas

20.1.3 Systematic Nomenclature for Aliphatic and Aromatic Molecules

20.2 Characteristic Organic Reactions

20.2.1 Electrophilic Substitution and Addition–Elimination Mechanisms

20.3 Shapes of Aromatic Organic Molecules; σ and π Bonds

20.3.1 Shape and Bonding in Benzene: sp² Hybridisation and Delocalised π System

20.4 Isomerism: Optical

20.4.1 Physical and Biological Properties of Optical Isomers

20.4.2 Concepts: Optically Active Compounds, Racemic Mixtures, Chiral Catalysts

21. Reaction Kinetics

21.1 Rate of Reaction

21.1.1 Definition of Rate of Reaction

21.1.2 Collision Theory: Effective and Non-Effective Collisions

21.1.3 Effect of Concentration and Pressure on Reaction Rate

21.1.4 Calculating Reaction Rate from Experimental Data

21.2 Effect of Temperature on Reaction Rates and the Concept of Activation Energy

21.2.1 Definition of Activation Energy

21.2.2 Boltzmann Distribution and Activation Energy

21.2.3 Effect of Temperature on Reaction Rate and Collision Frequency

21.3 Homogeneous and Heterogeneous Catalysts

21.3.1 Definition and Mechanism of Catalysis

21.3.2 Catalysts Lowering Activation Energy

21.3.3 Reaction Pathway Diagrams With and Without Catalysts

22. Organic Synthesis

22.1 Organic Synthesis

22.1.1 Identification of Functional Groups in Organic Molecules

22.1.2 Prediction of Properties and Reactions of Organic Molecules

22.1.3 Designing Multi-Step Synthetic Routes

22.1.4 Analysis of Synthetic Routes and Possible By-products

23. Hydrocarbons (Arenes)

23.1 Arenes

23.1.1 Substitution Reactions of Benzene and Methylbenzene

23.1.2 Mechanism of Electrophilic Substitution in Arenes

23.1.3 Predicting Halogenation: Side-Chain vs. Aromatic Ring

23.1.4 Directing Effects of Substituents in Electrophilic Substitution

24. Equilibria

24.1 Acids and Bases

24.1.1 Conjugate Acids and Bases

24.1.2 pH, Ka, pKa and Kw: Definitions and Calculations

24.1.3 Calculations of pH for Strong Acids, Strong Alkalis and Weak Acids

24.1.4 Definition and Formation of Buffer Solutions

24.1.5 Buffer Solutions in pH Control and Blood Chemistry

24.1.6 Calculations Involving Buffer Solutions

24.1.7 Solubility Product (Ksp) and Calculations

24.1.8 Common Ion Effect and Solubility Calculations

24.2 Partition Coefficients

24.2.1 Definition and Calculation of Partition Coefficient (Kpc)

24.2.2 Factors Affecting Partition Coefficients Based on Solute and Solvent Polarity

25. Analytical Techniques

25.1 Thin-Layer Chromatography

25.1.1 Terms and Interpretation: Stationary Phase, Mobile Phase, Rf Value, Solvent Front, Baseline

25.1.2 Explanation of Differences in Rf Values

25.2 Gas / Liquid Chromatography

25.2.1 Terms and Interpretation: Stationary Phase, Mobile Phase, Retention Time

25.2.2 Composition Analysis and Retention Time Explanation

25.3 Carbon-13 NMR Spectroscopy

25.3.1 Interpretation of C-13 NMR Spectra

25.3.2 Prediction of Number and Type of C Environments

25.4 Proton (¹H) NMR Spectroscopy

25.4.1 Interpretation of ¹H NMR Spectra: Chemical Shifts, Peak Areas, Splitting Patterns (n+1 Rule)

25.4.2 Prediction of Spectral Features

25.4.3 Use of TMS and Deuterated Solvents

25.4.4 Identification of O–H and N–H Protons by D₂O Exchange

26. Halogen Compounds (Arenes)

26.1 Halogen Compounds

26.1.1 Production of Halogenoarenes by Substitution

26.1.2 Reactivity Comparison: Halogenoalkanes vs. Halogenoarenes

27. Hydroxy Compounds (Phenol)

27.1 Alcohols

27.1.1 Reaction of Alcohols with Acyl Chlorides to Form Esters

27.2 Phenol

27.2.1 Directing Effects of Hydroxyl Group in Phenol Reactions

27.2.2 Reactions of Phenol with Bases, Sodium and Diazonium Salts

27.2.3 Nitration and Bromination of Phenol

27.2.4 Acidity of Phenol Compared to Water and Ethanol

27.2.5 Different Reaction Conditions for Phenol vs. Benzene

27.2.6 Production of Phenol from Diazonium Salts

28. Polymerisation

28.1 Addition Polymerisation

28.1.1 Description of Addition Polymerisation

28.1.2 Deduction of Repeat Units from Monomers

28.1.3 Identification of Monomers from Polymer Structures

28.1.4 Challenges of Disposal: Non-biodegradability and Harmful Combustion Products

29. Hydrocarbons

29.1 Alkanes

29.1.1 Production of Alkanes: Hydrogenation and Cracking

29.1.2 Combustion of Alkanes: Complete and Incomplete

29.1.3 Free-Radical Substitution of Alkanes

29.1.4 Mechanism of Free-Radical Substitution: Initiation, Propagation, Termination

29.1.5 Cracking for Useful Alkanes and Alkenes

29.1.6 Environmental Impact of Alkane Combustion

29.2 Alkenes

29.2.1 Production of Alkenes: Elimination and Dehydration Reactions

29.2.2 Electrophilic Addition Reactions of Alkenes

29.2.3 Oxidation of Alkenes: Cold Dilute and Hot Concentrated KMnO₄

29.2.4 Test for C=C Bond Using Aqueous Bromine

29.2.5 Mechanism of Electrophilic Addition in Alkenes

29.2.6 Stability of Carbocations and Markovnikov’s Rule

30. States of Matter

30.1 The Gaseous State: Ideal and Real Gases and pV = nRT

30.1.1 Origin of Pressure in Gases

30.1.2 Ideal Gas Assumptions

30.1.3 Ideal Gas Equation: pV = nRT

30.1.4 Applications of the Ideal Gas Equation

30.2 Bonding and Structure

30.2.1 Lattice Structures of Crystalline Solids

30.2.2 Structures and Bonding: Effect on Physical Properties

30.2.3 Deducing Structure and Bonding from Data

31. The Periodic Table: Chemical Periodicity

31.1 Periodicity of Physical Properties of the Elements in Period 3

31.1.1 Trends in Atomic Radius, Ionic Radius, Melting Point and Electrical Conductivity

31.1.2 Explanation of Melting Point and Conductivity Based on Structure and Bonding

31.2 Periodicity of Chemical Properties of the Elements in Period 3

31.2.1 Reactions of Elements with Oxygen, Chlorine and Water

31.2.2 Oxidation Numbers of Oxides and Chlorides

31.2.3 Reactions of Oxides with Water and pH of Solutions

31.2.4 Acid-Base Behaviour of Oxides and Hydroxides

31.2.5 Reactions of Chlorides with Water and pH of Solutions

31.2.6 Trends Explained by Bonding and Electronegativity

31.2.7 Bonding Types from Chemical and Physical Properties

31.3 Chemical Periodicity of Other Elements

31.3.1 Predicting Properties Based on Group Trends

31.3.2 Deducing Element Identity from Physical and Chemical Data

32. Organic Synthesis

32.1 Organic Synthesis

32.1.1 Identification of Organic Functional Groups

32.1.2 Predicting Properties and Reactions

32.1.3 Designing Multi-Step Synthetic Routes

32.1.4 Analyzing Synthetic Routes and Identifying By-products

33. Reaction Kinetics

33.1 Simple Rate Equations, Orders of Reaction and Rate Constants

33.1.1 Terms: Rate Equation, Order of Reaction, Overall Order, Rate Constant, Half-life, Rate-determining S

33.1.2 Deducing Orders from Graphs or Experimental Data

33.1.3 Interpretation of Concentration–Time and Rate–Concentration Graphs

33.1.4 Calculating Initial Rate from Data

33.1.5 Half-life of First-Order Reactions and Calculations

33.1.6 Calculating Rate Constants Using Different Methods

33.1.7 Predicting Reaction Mechanisms and Rate-Determining Steps

33.1.8 Identifying Catalysts and Intermediates from Mechanisms

33.1.9 Effect of Temperature on Rate Constants

33.2 Homogeneous and Heterogeneous Catalysts

33.2.1 Mode of Action of Heterogeneous Catalysts (Adsorption and Desorption)

33.2.2 Examples of Industrial Heterogeneous Catalysis

33.2.3 Mode of Action of Homogeneous Catalysts

33.2.4 Examples of Homogeneous Catalysis in Atmospheric Chemistry

34. Analytical Techniques

34.1 Infrared Spectroscopy

34.1.1 Analysis of Infrared Spectra to Identify Functional Groups

34.2 Mass Spectrometry

34.2.1 Interpretation of Mass Spectra: m/e Values and Isotopic Abundances

34.2.2 Calculation of Relative Atomic Mass and Molecular Mass

34.2.3 Identification of Molecules by Fragmentation Patterns

34.2.4 Determination of Carbon Atom Number Using [M+1]+ Peak

34.2.5 Detection of Bromine and Chlorine Atoms Using [M+2]+ Peak

35. Chemical Energetics

35.1 Enthalpy Change ΔH

35.1.1 Exothermic and Endothermic Reactions

35.1.2 Reaction Pathway Diagrams: Enthalpy Change and Activation Energy

35.1.3 Standard Conditions and Standard Enthalpy Changes

35.1.4 Enthalpy Changes: Reaction, Formation, Combustion, Neutralisation

35.1.5 Energy Transfer in Breaking and Making Bonds

35.1.6 Calculations Using Bond Energies

35.1.7 Calculations Using Experimental Data: q = mcΔT and ΔH = –mcΔT/n

35.2 Hess’s Law

35.2.1 Application of Hess’s Law to Construct Energy Cycles

35.2.2 Calculations Using Hess’s Law and Bond Energy Data

36. Group 2

36.1 Magnesium to Barium, and their Compounds

36.1.1 Reactions of Group 2 Elements with Oxygen, Water and Acids

36.1.2 Reactions of Group 2 Oxides, Hydroxides and Carbonates with Water and Acids

36.1.3 Thermal Decomposition of Group 2 Nitrates and Carbonates

36.1.4 Trends in Physical and Chemical Properties

36.1.5 Variation in Solubility of Hydroxides and Sulfates

37. Carboxylic Acids and Derivatives (Aromatic)

37.1 Esters

37.1.1 Production of Esters from Alcohols and Acyl Chlorides

37.2 Acyl Chlorides

37.2.1 Comparison of Hydrolysis of Acyl Chlorides, Alkyl Chlorides and Aryl Chlorides

37.2.2 Production of Acyl Chlorides from Carboxylic Acids

37.2.3 Addition–Elimination Mechanism of Acyl Chlorides

37.2.4 Reactions of Acyl Chlorides with Water, Alcohols, Phenols, Ammonia and Amines

37.3 Carboxylic Acids

37.3.1 Further Oxidation of Methanoic Acid and Ethanedioic Acid

37.3.2 Reactions of Carboxylic Acids with PCl₃, PCl₅ and SOCl₂

37.3.3 Effect of Chlorine Substitution on Acid Strength

37.3.4 Relative Acidities of Carboxylic Acids, Phenols and Alcohols

37.3.5 Production of Benzoic Acid from Alkylbenzene

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Definition and Expression of Stability Constants (Kstab)

Introduction