CAIE IGCSE Chemistry Comprehensive Study Guide: Redox, Physical Changes, Equilibrium, and Rates

Oxidation Numbers and Redox Reactions

Definition of Oxidation Number: The oxidation number of an element is the charge of an element or ion in a compound. This relates directly to the electrons gained or lost by the element or ion during the formation of the compound.
Representation of Oxidation Numbers: Oxidation numbers are represented by Roman numerals placed after the element name inside brackets.
- Example (Iron (II) oxide): The (II) indicates that the iron ion in the compound has an oxidation number of $+2$ .
- Example (Copper (I) chloride): The (I) indicates that the copper ion in the compound has an oxidation number of $+1$ .
Definition of Redox Reaction: Redox reactions are chemical processes where both oxidation and reduction occur simultaneously.
Oxidation and Reduction (General Definitions):
- Oxidation: The gain of oxygen by a substance.
- Reduction: The loss of oxygen from a substance.
- Specific Example: $H_2 + CuO \rightarrow Cu + H_2O$
  - Hydrogen has gained oxygen, so it has been oxidised.
  - Copper has lost oxygen, so it has been reduced.
Identifying Redox Reactions via Oxygen Transfer:
- Redox reactions must involve both the gain and loss of oxygen between reactants.
- Example: $H_2 + ZnO \rightarrow Zn + H_2O$
  - Hydrogen becomes water ( $H_2 \rightarrow H_2O$ ) by gaining oxygen (oxidised).
  - Zinc oxide becomes zinc ( $ZnO \rightarrow Zn$ ) by losing oxygen (reduced).
Further Identification Examples:
- Reaction: $Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2$
- Iron (III) oxide becomes Iron: This represents a loss of oxygen, meaning $Fe_2O_3$ has been reduced.
- Carbon monoxide becomes Carbon dioxide: This represents a gain of oxygen, meaning the $CO$ has been oxidised.

Oxidation and Reduction (Extended Content)

Electronic Definitions:
- Oxidation: The loss of electrons in a substance, which leads to an increase in its oxidation number.
- Reduction: The gain of electrons in a substance, which leads to a decrease in its oxidation number.
Identifying Redox via Electron Transfer: This is best achieved by splitting overall symbol equations into ionic half-equations to isolate the electron movement and remove spectator ions.
- Overall Equation: $Cl_2 + 2KI \rightarrow 2KCl + I_2$
- Half-equation 1 (Reduction): $Cl_2 + 2e^- \rightarrow 2Cl^-$ (Chlorine gains electrons).
- Half-equation 2 (Oxidation): $2I^- \rightarrow I_2 + 2e^-$ (Iodide ions lose electrons).
Oxidation Number Rules:
- (a) Elements in uncombined state: The oxidation number is always zero. (E.g., $Zn = 0$ , $Cl_2 = 0$ , $H_2 = 0$ ).
- (b) Monatomic ions: The oxidation number is equal to the charge of the ion. (E.g., $Zn^{2+} = +2$ , $Cl^- = -1$ , $H^+ = +1$ ).
- (c) Compounds: The sum of all oxidation numbers in a neutral compound must be zero.
  - Example (Zinc Chloride): In $ZnCl_2$ , $Zn$ is $+2$ and $Cl$ is $-1$ . Calculation: $(+2) + 2 \times (-1) = 0$ .
  - Example (Hydrogen Bromide): In $HBr$ , $H$ is $+1$ and $Br$ is $-1$ . Calculation: $(+1) + (-1) = 0$ .
- (d) Polyatomic ions: The sum of oxidation numbers is equal to the specific charge on the ion.
  - Example (Sulfate ion): In $SO_4^{2-}$ , the sum is $-2$ . If $S$ is $+6$ and each $O$ is $-2$ , the calculation is $(+6) + 4 \times (-2) = -2$ .
Tracking Electron Changes:
- If the oxidation number becomes more negative (decreases), electrons have been gained (Reduction).
- If the oxidation number becomes more positive (increases), electrons have been lost (Oxidation).
- Example Identification: $Cl_2 + 2KI \rightarrow 2KCl + I_2$
  - Chlorine is reduced from $0$ in $Cl_2$ to $-1$ in $KCl$ .
  - Iodide ions are oxidised from $-1$ in $KI$ to $0$ in $I_2$ .

Tests for Oxidising and Reducing Agents (Extended Content)

Acidified Aqueous Potassium Manganate (VII):
- Used as a test for reducing agents.
- The $Mn^{7+}$ ions in the solution are purple.
- When a reducing agent is present, $Mn^{7+}$ is reduced to $Mn^{2+}$ , causing a colour change from purple to colourless.
Potassium Iodide:
- Used as a test for oxidising agents.
- Contains $I^-$ ions.
- When an oxidising agent is present, $I^-$ ions are oxidised to $I_2$ .
- The solution changes from colourless to brown.
- Starch solution can be added as a further indicator, which turns the iodine solution blue-black.
Defining Agents:
- Oxidising Agent: A species that brings about oxidation by gaining electrons from other species. The agent itself is reduced. Typically non-metals or positive ions (e.g., $Br_2 + 2e^- \rightarrow 2Br^-$ ).
- Reducing Agent: A species that brings about reduction by losing/donating electrons to another species. The agent itself is oxidised. Typically metals or negative ions (e.g., $K \rightarrow K^+ + e^-$ ).
Agent Identification Example (Displacement): $Zn + CuSO_4 \rightarrow ZnSO_4 + Cu$
- Zinc is the reducing agent (Oxidised from $0$ to $+2$ , losing electrons).
- Copper ions ( $Cu^{2+}$ ) are the oxidising agent (Reduced from $+2$ to $0$ , gaining electrons).

Physical and Chemical Changes

Physical Changes:
- Characteristics: Can separate mixtures, involves changes of state (melting, freezing, etc.), and no new substances are formed.
- Reversibility: Generally reversible.
- Energy: No significant energy change.
- Appearance: Usually no colour change.
- Example: Fractional distillation of crude oil.
Chemical Changes:
- Characteristics: Requires a chemical reaction; products have different chemical compositions from reactants. Atoms are rearranged, but mass is conserved (Total Reactant Mass = Total Product Mass).
- Reversibility: Generally irreversible.
- Energy: Large energy changes; thermal energy is either absorbed (endothermic) or released (exothermic).
- Appearance: Colour changes can occur.
- Examples: Neutralisation reactions, digestion, burning, cooking.

Reversible Reactions and Equilibrium

Reversible Reactions: Some reactions allow products to react back into original reactants, denoted by the symbol $\rightleftharpoons$ . Their direction can be influenced by changing experimental conditions.
Hydrated vs. Anhydrous Compounds:
- Hydrated salts contain water molecules in their structure; anhydrous salts do not.
- Effect of Heat: Evaporates water from hydrated salts to leave anhydrous salt.
- Effect of Water: Adding water to anhydrous salt reforms the hydrated version.
- Copper (II) sulfate: Hydrated (blue crystals) $\rightleftharpoons$ Anhydrous (white crystals) + Water.
  - $CuSO_4.5H_2O(s) \rightleftharpoons CuSO_4(s) + 5H_2O(l)$ .
- Cobalt (II) chloride: Hydrated (pink crystals) $\rightleftharpoons$ Anhydrous (blue crystals) + Water.
  - $CoCl_2.6H_2O \rightleftharpoons CoCl_2 + 6H_2O$ .
Dynamic Equilibrium (Extended): Occurs when three conditions are met:
1. The reaction is in a closed system.
2. The rate of the forward reaction equals the rate of the reverse reaction.
3. The concentrations of reactants and products remain constant (no longer changing).
Le Chatelier’s Principle: If a system at equilibrium is changed, the system responds to counteract that change.
- Temperature:
  - Increase: Shifts toward the endothermic reaction direction.
  - Decrease: Shifts toward the exothermic reaction direction.
- Pressure (Gaseous reactions only):
  - Increase: Shifts toward the side with fewer gaseous molecules.
  - Decrease: Shifts toward the side with more gaseous molecules.
- Concentration:
  - Increase Reactant: Shifts to the right (more product made).
  - Decrease Product: Shifts to the right.
  - Decrease Reactant: Shifts to the left (more reactant made).
  - Increase Product: Shifts to the left.
- Catalyst: Has no effect on the position of equilibrium; it only allows the system to reach equilibrium faster.

Industrial Processes (Extended Content)

The Haber Process (Ammonia Production):
- Equation: $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$ .
- Sources: Nitrogen extracted from air via fractional distillation. Hydrogen produced by reacting methane with steam.
- Conditions: Temperature of $450^{\circ}C$ , Pressure of $20\,000\,kPa$ ( $200\,atm$ ), and an Iron catalyst.
- Logic: Forward reaction is exothermic. High temperature reduces yield but is necessary for rate (compromise). High pressure increases yield (fewer moles on right) but is expensive and carries explosion risks. Unreacted gases are recycled to reduce costs.
The Contact Process (Sulfuric Acid Production):
- Stage 2 Equation: $2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g)$ .
- Sources: Sulfur dioxide from burning sulfur in air or roasting sulfide ores. Oxygen from the air.
- Conditions: Temperature of $450^{\circ}C$ , Pressure of $200\,kPa$ ( $2\,atm$ ), and a Vanadium (V) oxide catalyst.
- Logic: Forward reaction is exothermic (compromise temperature used). Increased pressure would increase yield, but with $SO_3$ being an acidic gas, the risk of explosion is too high at extreme pressures; hence only $2\,atm$ is used.

Rate of Reaction

Definition: The time taken for a reactant to be used up or a product to be formed.
Factors Affecting Rate:
- Concentration: Increase leads to faster rate.
- Pressure (Gases): Increase leads to faster rate.
- Surface Area (Solids): Increase (e.g., using powder instead of lumps) leads to faster rate.
- Temperature: Increase leads to faster rate.
- Catalysts/Enzymes: Adding a catalyst increases rate; it remains unchanged at the end and does not affect the type of products formed.
Practical Measurement:
- $Rate = \frac{\text{mass of reactant used}}{\text{time}}$
- $Rate = \frac{\text{mass of product formed}}{\text{time}}$
- $Rate = \frac{\text{volume of gas formed}}{\text{time}}$
- Units: $g/s$ or $cm^3/s$ .
Variables and Apparatus:
- Independent Variable: Factor changed (e.g., temperature intervals).
- Dependent Variable: Factor measured (e.g., mass produced).
- Controlled Variables: Factors kept constant (e.g., concentration, surface area).
- Equipment: Balance (mass), Stopwatch (time), Gas Syringe (volume of gas), Thermometer (temperature), Beaker with cross (precipitate visibility test).
Graph Interpretation:
- Steeper gradient indicates a faster rate.
- Horizontal line means the reaction is complete.
- Tangent Calculation: To find the rate at a specific time, draw a tangent and calculate gradient: $Rate = \frac{\Delta y}{\Delta x}$ .
- Initial rate is found by drawing the tangent at $t = 0\,s$ .

Collision Theory (Extended Content)

Requirements for Reaction:
1. Reactant particles must collide successfully.
2. Particles must possess energy equal to or greater than the Activation Energy ( $E_a$ ).
Explanations of Rate Factors:
- Concentration & Pressure: More particles per unit volume results in more frequent collisions.
- Surface Area: A fine powder provides a greater surface-area-to-volume ratio, exposing more particles and resulting in more frequent collisions.
- Temperature: Increases kinetic energy of particles, causing them to move faster. This lead to more frequent collisions and a higher proportion of collisions having the required $E_a$ .
- Catalysts: Increases rate by lowering the Activation Energy ( $E_a$ ) required for the reaction. This allows more particles to have sufficient energy for successful collisions.
Reaction Pathway Diagrams: Shows how a catalyst provides an alternative pathway with a lower $E_a$ barrier.