Chemical Reactions

1. Physical and Chemical Changes

Physical Change

• No new substance formed.
• Often reversible (e.g., melting ice).
• Change in state or shape.

Chemical Change

• New substance(s) formed.
• Often irreversible.
• Accompanied by energy change, colour change, or gas evolution.

2. Rates of Reaction

Collision Theory

• For a reaction to occur, particles must collide with:
  1. Sufficient Energy: Energy $\ge$ Activation Energy ($E_a$).
  2. Correct Orientation.
• Effect of Factors:
  • Temp/Concentration/Pressure: Increase collision frequency and/or proportion of particles with energy $\ge E_a$.
  • Catalysts: Provide an alternative pathway with a lower activation energy ($E_a$).

Factors Affecting Rate

• Concentration: Higher concentration $\rightarrow$ more particles per unit volume $\rightarrow$ higher rate.
• Pressure (Gases): Higher pressure $\rightarrow$ particles closer together $\rightarrow$ higher rate.
• Surface Area: Smaller particles (powder) $\rightarrow$ more exposed surface $\rightarrow$ higher rate.
• Temperature: Higher temperature $\rightarrow$ particles move faster $\rightarrow$ higher rate.
• Catalyst: Substance that increases rate without being consumed. Enzymes are biological catalysts.

Investigating Rates

• Methods:
  • Measuring mass loss (if gas escapes).
  • Measuring volume of gas produced (using gas syringe).
  • Measuring time for a colour change or precipitate to form.
• Evaluation: Consider accuracy of apparatus (e.g., gas syringe vs. measuring cylinder) and precision of timing.

3. Reversible Reactions and Equilibrium

Reversible Reactions

• Reactions that can proceed in both forward and reverse directions.
• Symbol: $\rightleftharpoons$
• Hydrated vs Anhydrous:
  • $\text{CuSO}_4\cdot 5\text{H}_2\text{O}$ (Blue) $\rightleftharpoons \text{CuSO}_4$ (White) + $5\text{H}_2\text{O}$
  • $\text{CoCl}_2\cdot 6\text{H}_2\text{O}$ (Pink) $\rightleftharpoons \text{CoCl}_2$ (Blue) + $6\text{H}_2\text{O}$

Dynamic Equilibrium

• Occurs in a closed system.
• Condition: Rate of forward reaction = Rate of reverse reaction.
• Observation: Concentrations of reactants and products remain constant.

Equilibrium Shifts

• Le Chatelier’s Principle: If a system at equilibrium is stressed, it shifts to oppose the change.
• Temperature: Increase in temp shifts equilibrium in the endothermic direction.
• Pressure: Increase in pressure shifts equilibrium towards the side with fewer gas molecules.
• Concentration: Increasing a reactant shifts equilibrium towards the products.

4. Industrial Processes

Haber Process (Ammonia Synthesis)

• Equation: $\text{N}_2(\text{g}) + 3\text{H}_2(\text{g}) \rightleftharpoons 2\text{NH}_3(\text{g})$
• Sources: $\text{N}_2$ from air, $\text{H}_2$ from methane (natural gas).
• Conditions: 450 °C, 20,000 kPa, Iron catalyst.
• Optimization: Balance between rate (high temp) and yield (low temp for exothermic reaction).

Contact Process (Sulfuric Acid Synthesis)

• Equation: $2\text{SO}_2(\text{g}) + \text{O}_2(\text{g}) \rightleftharpoons 2\text{SO}_3(\text{g})$
• Sources: $\text{SO}_2$ from sulfur burning/roasting, $\text{O}_2$ from air.
• Conditions: 450 °C, 200 kPa, Vanadium(V) oxide catalyst.

5. Redox Reactions

Basic Definitions

• Redox: A reaction where oxidation and reduction happen simultaneously.
• Oxidation: Gain of oxygen.
• Reduction: Loss of oxygen.
• Identification: Look for gain/loss of oxygen in equations.

Advanced Redox

• Oxidation: Loss of electrons or increase in oxidation number.
• Reduction: Gain of electrons or decrease in oxidation number.
• Oxidation Numbers: Use Roman numerals (e.g., $\text{Fe(III)} = \text{Fe}^{3+}$).
• Oxidising Agent: Substance that oxidises another; it is itself reduced.
• Reducing Agent: Substance that reduces another; it is itself oxidised.

Identifying Redox via Colour Changes

• Acidified $\text{KMnO}_4$: Purple $\rightarrow$ Colourless (indicates reduction/oxidising agent acting).
• Aqueous $\text{KI}$: Colourless $\rightarrow$ Brown (indicates oxidation/reducing agent acting).