IAL AS Chemistry – Unit 2 Complete Revision Notes (Theory Only)

Specification: Edexcel International Advanced Level (IAL)
Unit: AS Unit 2
Core Practicals: NOT included

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1. Energetics

1.1 Enthalpy Change (ΔH)

Enthalpy change is the energy transferred as heat during a reaction at constant pressure.

• Exothermic reaction: ΔH < 0 (energy released)
• Endothermic reaction: ΔH > 0 (energy absorbed)

1.2 Standard Conditions

• Temperature: 298 K
• Pressure: 100 kPa
• Concentration: 1 mol dm⁻³

1.3 Types of Enthalpy Change

• ΔH_f° – standard enthalpy of formation
• ΔH_c° – standard enthalpy of combustion
• ΔH_n – enthalpy of neutralisation (≈ −57 kJ mol⁻¹)

1.4 Hess’s Law

The total enthalpy change of a reaction is independent of the route taken.

Calculation using formation enthalpies:

$$\mathrm{\Delta H}=\sum \mathrm{\Delta H}{f}_{\mathrm{products}}-\sum \mathrm{\Delta H}{f}_{\mathrm{reactants}}$$

Exam Tip

Always multiply ΔH values by the mole ratio before subtracting.

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2. Rates of Reaction

2.1 Definition

Rate of reaction is the change in concentration of a reactant or product per unit time.

2.2 Collision Theory

• Particles must collide
• Collisions must have energy ≥ activation energy
• Correct orientation required

2.3 Factors Affecting Rate

• Concentration
• Pressure (gases)
• Temperature
• Surface area
• Catalyst

2.4 Activation Energy (E_a)

Minimum energy required for a successful collision.

Catalysts lower E_a by providing an alternative reaction pathway.

2.5 Rate Equations

General form:

$$\mathrm{rate}=k{\mathrm{[A]}}^m{\mathrm{[B]}}^n$$

• Orders must be found experimentally
• Overall order = m + n

2.6 Units of Rate Constant

Depend on overall order (exam favourite).

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3. Chemical Equilibria

3.1 Dynamic Equilibrium

A reversible reaction where forward and backward reactions occur at equal rates.

3.2 Le Chatelier’s Principle

If a system at equilibrium is disturbed, it shifts to oppose the change.

3.3 Effects on Equilibrium

• Concentration: shifts position, K unchanged
• Pressure: affects gaseous equilibria
• Temperature: changes K

3.4 Equilibrium Constant (K_c)

$$K_c=\frac{{\mathrm{[C]}}^c{\mathrm{[D]}}^d}{{\mathrm{[A]}}^a{\mathrm{[B]}}^b}$$

Exam Notes

• Pure solids and liquids are excluded
• Large K → products favoured

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4. Redox Chemistry

4.1 Oxidation and Reduction

• Oxidation = loss of electrons
• Reduction = gain of electrons

4.2 Oxidation Numbers

• Elements = 0
• Group 1 metals = +1
• Oxygen usually = −2
• Hydrogen usually = +1

4.3 Redox Equations

Example:

$$\mathrm{Zn}+\mathrm{Cu}{\mathrm{2+}}^{}\to \mathrm{Zn}{\mathrm{2+}}^{}+\mathrm{Cu}$$

Zn is oxidised, Cu²⁺ is reduced.

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5. Common Exam Questions

Worked Calculation Example

Calculate ΔH using combustion data.

Method:

1. Write balanced equation
2. Multiply ΔH values
3. Apply Hess’s law

Explanation Question Tip

Always mention:

• Particle behaviour
• Energy changes
• Effect on rate or equilibrium

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6. Quick Formula Sheet

• q = mcΔT
• rate = Δconcentration / time
• ΔH = Σ(products) − Σ(reactants)

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7. Final Exam Advice

• Learn definitions word-for-word
• Show full working in calculations
• Always state units
• Do not mention experiments or practical steps

END OF UNIT 2 NOTES