Chemistry

3. Physical Chemistry
3.1 Energetics>3.2 Rates of Reaction>3.3 Reversible Reactions and Equilibria>

1. Principles of Chemistry

1.2 Elements, Compounds & Mixtures >1.3 Atomic Structure >1.4 The Periodic Table >1.5 Chemical Formulae, Equations and Calculations >1.6 Ionic Bonding >1.7 Covalent Bonding >1.8 Metallic Bonding >1.9 Electrolysis >

2. Inorganic Chemistry

2.1 Group 1 (Alkali Metals) – Lithium, Sodium and Potassium >2.2 Group 7 (Halogens) – Chlorine, Bromine and Iodine >2.3 Gases in the Atmosphere >2.4 Reactivity Series >2.5 Extraction and Uses of Metals >2.6 Acids, Alkalis and Titrations >2.7 Acids, Bases and Salt Preparations >2.8 Chemical Tests >

4. Organic Chemistry

4.1 Introduction >4.2 Crude Oil >4.3 Alkanes >4.4 Alkenes >4.5 Alcohols >4.6 Carboxylic Acids >4.7 Esters >4.8 Synthetic Polymers >

3.3 Reversible Reactions and Equilibria

3.3.1 Know that some reactions are reversible and this is indicated by the symbol ⇌ in equations

COMPLETE REACTION

Some reactions go to completion
Meaning that: when all reactants are used up to form products, the reaction stops
Indicated by the “→” in the chemical equation
For example: 2H2O2 → 2H2O + O2

REVERSIBLE REACTION

Some reactions are reversible
Meaning that: products can react or decompose and form reactants again
Reaction occurs in both directions
- Forward reaction (reactant → product)
- Reverse reaction (product → reactant)
Indicated by the “⇌” in the chemical equation
For example: CuSO4・5H2O ⇌ CuSO4 + 5H2O

3.3.2 Describe reversible reactions such as the dehydration of hydrated copper(II) sulfate and the effect of heat on ammonium chloride

Examples of Reversible Reaction

Dehydration of hydrated copper(II) sulphate
Thermal decomposition of ammonium chloride

Example 1 Dehydration of Hydrated Copper(II) Sulphate

When hydrated CuSO₄・5H₂O crystals are heated in a test tube
- Water is removed
- Blue crystals turn into white powder
When anhydrous CuSO4 is added to water
- White powder turns blue
- Heat is given off (reaction is exothermic)

CuSO4・5H2O (s) ⇌ CuSO4 (s) + 5H2O (l)

EXAMPLE 2 THERMAL DECOMPOSITION OF AMMONIUM CHLORIDE

Heating ammonium chloride produces ammonia and hydrogen chloride gases

NH₄Cl (s) → NH₃ (g) + HCl (g)

As the hot gases cool down they recombine to form solid ammonium chloride

NH₃ (g) + HCl (g) → NH₄Cl (s)

ammonium chloride ⇌ ammonia + hydrogen chloride

NH4Cl (s) ⇌ NH3 (g) + HCl (g)

3.3.3C Know that a reversible reaction can reach dynamic equilibrium in a sealed container

DYNAMIC EQUILIBRIUM

Dynamic equilibrium can be reached when a reversible reaction is taken place in a closed system
None of the participating chemical substances are able to leave the container

3.3.4C Know that the characteristics of a reaction at dynamic equilibrium are:

The forward and reverse reactions occur at the same rate
The concentrations of reactants and products remain constant.

When a reaction is at dynamic equilibrium:

Both forward and reverse reactions are occuring at the same rate
Therefore, the net change of the system is constant, and the concentrations of reactants and products respectively remain constant
Dynamic equilibrium does not mean the concentrations of reactants and products are the same
Dynamic equilibrium does not mean the reaction stops, both forward and reverse reactions are still happening.

3.3.5C Understand why a catalyst does not affect the position of equilibrium in a reversible reaction

Presence of catalyst does not affect the position of equilibrium

Catalyst increases the rate of reaction
- In this case, it increases the rate of both forward and reverse reactions
It does not change the concentration of reactants and products at equilibrium
Position of equilibrium is not affected

3.3.6C Know the effect of changing either temperature or pressure on the position of equilibrium in a reversible reaction:

An increase (or decrease) in temperature shifts the position of equilibrium in the direction of the endothermic (or exothermic) reaction
An increase (or decrease) in pressure shifts the position of equilibrium in the direction that produces fewer (or more) moles of gas

References to Le Chatelier’s principle are not required

When the condition of a system at equilibrium is changed:

The system would respond to oppose the change
This is known as Le Chatelier’s principle (not required)
Factors to be discussed:
- Temperature
- Pressure

Effect of Temperature

Example

ICl + Cl₂ ⇌ ICl₃

Dark brown Yellow

When the equilibrium mixture is heated, it becomes dark brown in colour.

HOW DO WE KNOW WHETHER THE BACKWARD REACTION IS EXOTHERMIC OR ENDOTHERMIC?

Equilibrium has shifted to the left as the colour dark brown means that more of ICI is produced
Increasing temperature moves the equilibrium in the endothermic direction
So the backward reaction is endothermic

Effect of Pressure

Example

2NO₂ ⇌ N₂O₄

Brown gas Colourless gas

WHAT IS THE EFFECT OF AN INCREASE IN PRESSURE ON THE POSITION OF EQUILIBRIUM?

Number of gas molecules on the left = 2
Number of gas molecules on the right = 1
An increase in pressure will cause equilibrium to shift in the direction that produces the smaller number of molecules of gas (2 → 1)
So equilibrium shifts to the right