Understanding mole fraction and partial pressure is fundamental in the study of chemical equilibria, especially within the context of reversible reactions and dynamic equilibrium. These concepts are pivotal for AS & A Level students pursuing Chemistry - 9701, as they form the basis for analyzing gaseous mixtures and predicting the behavior of systems in equilibrium.

Key Concepts

Mole Fraction: Definition and Calculation

The mole fraction, denoted as $ X $, is a dimensionless quantity that represents the ratio of the number of moles of a particular component to the total number of moles in the mixture. It is a way to express the concentration of a component in a mixture without referring to its volume or mass. $$ X_i = \frac{n_i}{n_{total}} $$ where $ X_i $ is the mole fraction of component $ i $, $ n_i $ is the number of moles of component $ i $, and $ n_{total} $ is the total number of moles of all components in the mixture. **Example:** Consider a gas mixture containing 2 moles of nitrogen ($ N_2 $) and 3 moles of oxygen ($ O_2 $). The mole fractions are calculated as follows: $$ X_{N_2} = \frac{2}{2+3} = \frac{2}{5} = 0.4 $$ $$ X_{O_2} = \frac{3}{2+3} = \frac{3}{5} = 0.6 $$

Partial Pressure: Definition and Dalton’s Law

Partial pressure refers to the pressure exerted by a single component of a gaseous mixture. According to Dalton's Law of Partial Pressures, the total pressure of a gas mixture is equal to the sum of the partial pressures of each individual component. $$ P_{total} = P_1 + P_2 + P_3 + \ldots + P_n $$ where $ P_{total} $ is the total pressure and $ P_1, P_2, P_3, \ldots, P_n $ are the partial pressures of the components. Dalton's Law can also be expressed using mole fractions: $$ P_i = X_i \cdot P_{total} $$ where $ P_i $ is the partial pressure of component $ i $, $ X_i $ is its mole fraction, and $ P_{total} $ is the total pressure of the gas mixture. **Example:** Using the previous gas mixture example with a total pressure of 1 atm, $$ P_{N_2} = X_{N_2} \cdot P_{total} = 0.4 \cdot 1 \, atm = 0.4 \, atm $$ $$ P_{O_2} = X_{O_2} \cdot P_{total} = 0.6 \cdot 1 \, atm = 0.6 \, atm $$

Relationship Between Mole Fraction and Partial Pressure

The mole fraction directly relates to the partial pressure in a gaseous mixture. By understanding one, the other can be easily determined using Dalton's Law. This relationship is crucial in applications such as gas stoichiometry and equilibrium calculations.

Ideal Gas Law and Its Application

The Ideal Gas Law integrates the concepts of mole fraction and partial pressure in the context of gas behavior: $$ PV = nRT $$ where $ P $ is pressure, $ V $ is volume, $ n $ is the number of moles, $ R $ is the universal gas constant, and $ T $ is temperature. For a mixture of gases, the Ideal Gas Law can be applied to each component individually using partial pressures.

Calculating Mole Fraction in Multicomponent Systems

In systems with more than two components, the mole fraction of each component is calculated similarly by dividing the individual moles by the total moles. The sum of all mole fractions in a mixture equals 1. $$ \sum_{i=1}^{n} X_i = 1 $$ **Example:** In a mixture containing 1 mole of helium ($ He $), 2 moles of neon ($ Ne $), and 3 moles of argon ($ Ar $): $$ X_{He} = \frac{1}{1+2+3} = \frac{1}{6} \approx 0.167 $$ $$ X_{Ne} = \frac{2}{6} \approx 0.333 $$ $$ X_{Ar} = \frac{3}{6} = 0.5 $$

Applications of Mole Fraction and Partial Pressure

These concepts are not only fundamental in theoretical chemistry but also have practical applications in various fields such as: - **Chemical Engineering:** Designing reactors and separation processes. - **Environmental Science:** Modeling atmospheric gas compositions. - **Medicine:** Understanding gas mixtures in respiratory therapies. - **Industrial Processes:** Managing gas mixtures in manufacturing settings.

Temperature and Its Effect on Partial Pressure

Temperature influences the behavior of gas mixtures. According to the Ideal Gas Law, as temperature increases, the pressure of a gas increases if volume is held constant. This relationship affects partial pressures and, consequently, the mole fraction distribution in dynamic equilibria.

Real Gases vs. Ideal Gases

While the Ideal Gas Law provides a good approximation, real gases deviate from ideality under high pressure or low temperature. Understanding these deviations is essential for accurate calculations in practical scenarios.

Henry’s Law and Partial Pressure

Henry's Law relates the solubility of a gas in a liquid to its partial pressure above the liquid. It is expressed as: $$ C = k_H \cdot P $$ where $ C $ is the concentration of the dissolved gas, $ k_H $ is Henry's law constant, and $ P $ is the partial pressure. This law is particularly important in understanding gas solubility in biological systems and industrial applications.

Partial Pressure in Respiratory Physiology

In human physiology, the concept of partial pressure is crucial for understanding gas exchange in the lungs. Oxygen and carbon dioxide partial pressures drive the diffusion of these gases across the alveolar membrane into and out of the blood.

Calculating Partial Pressures in Chemical Equilibria

In reversible reactions, partial pressures determine the position of equilibrium. By applying the principles of mole fraction and Dalton’s Law, one can predict how changes in pressure or composition affect the equilibrium state.

Le Chatelier’s Principle and Partial Pressure

Le Chatelier’s Principle states that a system at equilibrium will adjust to counteract any imposed changes. Changes in partial pressure of reactants or products can shift the equilibrium position, which can be quantitatively analyzed using mole fraction and partial pressure calculations.

Advanced Concepts

Mathematical Derivation of Dalton’s Law

Dalton’s Law can be derived from the kinetic molecular theory, which assumes that gases consist of a large number of small particles in random motion. Each gas in a mixture behaves independently, contributing to the total pressure without interacting with other gases. Let’s consider a mixture of two ideal gases, Gas A and Gas B: - Let $ P_A $ and $ P_B $ be the partial pressures. - The total pressure $ P_{total} $ is: $$ P_{total} = P_A + P_B $$ Each partial pressure is proportional to its mole fraction: $$ P_A = X_A \cdot P_{total} $$ $$ P_B = X_B \cdot P_{total} $$ This derivation assumes ideal behavior, where interactions between gas molecules are negligible.

Raoult’s Law and its Relation to Partial Pressure

Raoult’s Law describes the partial pressure of a solvent in an ideal solution: $$ P_A = X_A^0 \cdot P_{A}^0 $$ where $ X_A^0 $ is the mole fraction of solvent in the solution and $ P_{A}^0 $ is the vapor pressure of the pure solvent. While Dalton’s Law applies to gas mixtures, Raoult’s Law is applicable to solutions, highlighting the interplay between mole fraction and vapor pressure in different states of matter.

Graham’s Law of Effusion and Molecular Mass

Graham’s Law relates the rate of effusion of a gas to its molecular mass: $$ \frac{Rate_1}{Rate_2} = \sqrt{\frac{M_2}{M_1}} $$ where $ Rate_1 $ and $ Rate_2 $ are the effusion rates of gases 1 and 2, and $ M_1 $ and $ M_2 $ are their molar masses. This law illustrates the relationship between kinetic properties of gases and their molecular characteristics, expanding the understanding of partial pressures in dynamic systems.

Vapor-Liquid Equilibrium (VLE)

In systems where liquids and vapors coexist, vapor-liquid equilibrium is achieved when the partial pressures of each component in the vapor phase are proportional to their mole fractions in the liquid phase, as per Raoult’s Law. Analyzing VLE is crucial in distillation processes, where separation of components is driven by differences in partial pressures and mole fractions.

Non-Ideal Gas Behavior and Partial Pressure

Real gases exhibit deviations from ideal behavior due to intermolecular forces and finite molecular sizes. The Van der Waals equation modifies the Ideal Gas Law to account for these factors: $$ \left( P + \frac{a n^2}{V^2} \right) (V - nb) = nRT $$ where $ a $ and $ b $ are Van der Waals constants specific to each gas. These adjustments affect partial pressure calculations in non-ideal gas mixtures.

Henry’s Law Constants and Partial Pressure

Henry’s Law constants vary with temperature and provide a quantitative measure of gas solubility. Understanding these constants allows for the calculation of partial pressures in solution-based equilibria, bridging the gap between gaseous and dissolved states.

Partial Pressure in Electrochemistry

In electrochemistry, partial pressures of gaseous reactants influence electrode potentials and reaction rates. The Nernst equation incorporates partial pressures to predict cell potentials under non-standard conditions. $$ E = E^\circ - \frac{RT}{nF} \ln Q $$ where $ Q $ includes activities that can depend on partial pressures in gaseous systems.

Partial Pressure and Fugacity

Fugacity accounts for non-ideal behavior by substituting partial pressure with a corrected value that better represents the escaping tendency of a gas. This concept is essential for accurate thermodynamic calculations in real gas systems.

Intermolecular Forces and Partial Pressure

Intermolecular forces influence both physical properties and chemical behavior. Variations in partial pressures can alter the extent of these forces in a mixture, affecting properties like boiling point and solubility.

Partial Pressure and Reaction Kinetics

The rate of a reaction involving gaseous reactants depends on their partial pressures. Higher partial pressures typically increase the reaction rate by increasing the frequency of effective collisions.

Equilibrium Constants and Partial Pressure

In gas-phase equilibria, equilibrium constants are expressed in terms of partial pressures: $$ K_p = \frac{P_{C}^c P_{D}^d}{P_{A}^a P_{B}^b} $$ where $ P_{C} $, $ P_{D} $, $ P_{A} $, and $ P_{B} $ are partial pressures of the reactants and products, and $ a $, $ b $, $ c $, and $ d $ are their stoichiometric coefficients.

Colligative Properties and Mole Fraction

Colligative properties depend on the mole fraction of solute particles. While primarily discussed in solutions, understanding mole fraction’s influence on boiling point elevation and freezing point depression provides a broader perspective on its impact in various systems.

Comparison Table

Aspect	Mole Fraction	Partial Pressure
Definition	Ratio of moles of a component to total moles in the mixture	Pressure exerted by a single component in a gaseous mixture
Expression	$X_i = \frac{n_i}{n_{total}}$	$P_i = X_i \cdot P_{total}$
Unit	Dimensionless	Pressure units (atm, Pa, etc.)
Application	Concentration in mixtures, solutions	Behavior of gas mixtures, equilibria calculations
Dependence on Total Quantity	Depends on total moles	Depends on total pressure
Use in Laws	Raoult’s Law, mole fraction dependent properties	Dalton’s Law, partial pressure dependent properties

Summary and Key Takeaways

Mole fraction represents the ratio of a component’s moles to the total moles in a mixture.
Partial pressure is the pressure contributed by an individual gas in a mixture, as per Dalton’s Law.
Dalton’s and Raoult’s Laws link mole fraction and partial pressure, essential for equilibrium calculations.
Advanced applications include real gas behavior, VLE, and interdisciplinary connections with physiology and engineering.
Understanding these concepts is crucial for solving complex chemical equilibrium problems in AS & A Level Chemistry.

Examiner Tip

Tips

To easily remember the relationship between mole fraction and partial pressure, use the mnemonic "Mole Presses Total": Mole fraction multiplied by total pressure gives the partial pressure ($P_i = X_i \cdot P_{total}$). When dealing with multiple components, always verify that the sum of mole fractions equals one. Practice converting between mole fraction and percentage composition to strengthen your understanding for exams.

Did You Know

Did you know that the concept of mole fraction is not only essential in chemistry but also plays a critical role in environmental science? For instance, understanding the mole fractions of greenhouse gases like carbon dioxide and methane helps scientists assess their impact on global warming. Additionally, mole fractions are pivotal in the formulation of breathing gas mixtures used by divers, ensuring safety under varying pressure conditions.

Common Mistakes

One common mistake students make is confusing mole fraction with mass fraction. For example, they might incorrectly use mass instead of moles when calculating mole fractions. Another error is neglecting to ensure that the sum of all mole fractions equals one, leading to inaccurate partial pressure calculations. Additionally, students often misapply Dalton’s Law by not considering the total pressure correctly when determining individual partial pressures.

FAQ

What is the mole fraction of a component in a mixture?

The mole fraction is the ratio of the number of moles of a specific component to the total number of moles in the mixture, expressed as $X_i = \frac{n_i}{n_{total}}$.

How does Dalton’s Law relate to partial pressure?

Dalton’s Law states that the total pressure of a gaseous mixture is equal to the sum of the partial pressures of each individual component. Mathematically, $P_{total} = P_1 + P_2 + \ldots + P_n$.

Can mole fraction be greater than one?

No, mole fractions are always between 0 and 1 because they represent a ratio of a part to the whole, and the sum of all mole fractions in a mixture equals one.

How do temperature changes affect partial pressures in a gas mixture?

According to the Ideal Gas Law, an increase in temperature, while keeping volume constant, will increase the pressure of each gas component in the mixture, thereby increasing their partial pressures.

What is the difference between Dalton’s Law and Raoult’s Law?

Dalton’s Law applies to the partial pressures of gases in a mixture, while Raoult’s Law applies to the vapor pressure of components in an ideal liquid solution. Dalton deals with gas-phase mixtures, whereas Raoult’s relates to liquid solutions.

How is mole fraction used in calculating equilibrium constants?

In gas-phase equilibria, mole fractions are used to determine the partial pressures of reactants and products. These partial pressures are then used to express the equilibrium constant $K_p$, which relates the concentrations of products to reactants at equilibrium.

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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Mole Fraction and Partial Pressure

Introduction