All of AQA CHEMISTRY Paper 2 in 25 minutes - GCSE Science Revision

GCSE AQA Chemistry Paper 2 Study Guide

1. Topics Covered

• Chemical reactions: Types of reactions, balancing equations.

• Energy changes in reactions: Exothermic and endothermic reactions, energy profile diagrams.

• Rates of reaction: Factors affecting the rate, calculations involving rates.

• Equilibria: Le Chatelier's principle, dynamic equilibrium.

• Acids and bases: pH scale, strong vs weak acids, neutralization reactions.

• Organic chemistry: Alkanes, alkenes, functional groups, alcohols, carboxylic acids.

• Electrolysis: Principles and processes of electrolysis, applications.

• The periodic table: Trends, groups, and periods, properties of metals and non-metals.

• Calculating concentrations: Molarity, dilutions, and concentration calculations.

2. Key Equations

• Rate of reaction = change in concentration / time.

• pH = -log[H+].

• Molarity (M) = moles of solute / volume of solution (L).

• Balanced chemical equation for combustion of hydrocarbons: CxHy + O2 → CO2 + H2O.

3. Important Definitions

• Exothermic reaction: A reaction that releases energy to the surroundings.

• Endothermic reaction: A reaction that absorbs energy from the surroundings.

• Dynamic equilibrium: A state in which the rates of the forward and backward reactions are equal.

4. Study Tips

• Practice past papers and mark schemes to familiarize yourself with question formats.

• Use flashcards for important terms and definitions.

• Engage in group study sessions to discuss and explain concepts with peers.

• Schedule regular revision, breaking down topics into manageable sections.

Detailed Notes on GCSE AQA Chemistry Paper 2

1. Chemical Reactions

Types of Reactions

• Combination: Two or more reactants combine to form a single product (e.g. A + B → AB).

• Decomposition: A single compound breaks down into two or more simpler products (e.g. AB → A + B).

• Displacement: An element displaces another in a compound (e.g. A + BC → AC + B).

• Redox: Involves the transfer of electrons; reduction and oxidation occur simultaneously.

Balancing Equations

• Ensure the number of atoms for each element is the same on both sides of the equation.

• Use coefficients to balance equations rather than changing subscripts in chemical formulas.

2. Energy Changes in Reactions

Exothermic Reactions

• Release energy, usually in the form of heat, causing the surroundings to become warmer.

• Example: Combustion reactions (e.g. burning of fuels).

Endothermic Reactions

• Absorb energy from the surroundings, causing the temperature to decrease.

• Example: Photosynthesis.

Energy Profile Diagrams

• Illustrate the energy changes during a reaction.

• Show the energy of reactants, products, and the activation energy needed for the reaction to proceed.

3. Rates of Reaction

Factors Affecting Rate

• Temperature: Higher temperatures increase kinetic energy, leading to more collisions.

• Concentration: Increased concentration of reactants typically leads to a higher rate.

• Surface Area: Larger surface areas allow more collisions and faster reactions (e.g. powdered vs chunks).

• Catalysts: Substances that speed up reactions without being consumed by lowering activation energy.

Calculations Involving Rates

• Rate = change in concentration / time.

4. Equilibria

Le Chatelier's Principle

• States that if a change is applied to a system at equilibrium, the system adjusts to counteract that change.

• Changes can include concentration, pressure, or temperature.

Dynamic Equilibrium

• Occurs when the rate of the forward reaction equals the rate of the backward reaction.

• Concentrations of reactants and products remain constant.

5. Acids and Bases

pH Scale

• Ranges from 0 (strong acids) to 14 (strong bases).

• pH 7 is neutral (pure water).

Strong vs Weak Acids

• Strong acids (e.g. HCl, H2SO4) dissociate completely in water.

• Weak acids (e.g. CH3COOH) do not fully dissociate in water.

Neutralization Reactions

• Reaction between an acid and a base to produce salt and water (e.g. HCl + NaOH → NaCl + H2O).

6. Organic Chemistry

Alkanes

• Saturated hydrocarbons with single bonds (e.g. methane, ethane).

• General formula: CnH2n+2.

Alkenes

• Unsaturated hydrocarbons with at least one double bond (e.g. ethene).

• General formula: CnH2n.

Functional Groups

• Groups of atoms responsible for the characteristic reactions of a compound (e.g. -OH is hydroxyl in alcohols).

Alcohols

• Contain the -OH functional group (e.g. ethanol).

• Used as solvents, fuels, and in alcoholic beverages.

Carboxylic Acids

• Contain the -COOH functional group (e.g. acetic acid).

• React with alcohols to produce esters.

7. Electrolysis

Principles and Processes

• Involves passing an electric current through a liquid or solution to cause a chemical change.

• Separates ionic compounds into their constituent elements.

Applications

• Used in metal extraction (e.g. aluminum), electroplating, and in the production of chlorine and hydrogen.

8. The Periodic Table

Trends

• Groups: Vertical columns, elements in the same group have similar chemical properties.

• Periods: Horizontal rows, properties change in a predictable way across a period (e.g. increasing atomic number).

Properties of Metals and Non-metals

• Metals: Good conductors of electricity and heat, malleable and ductile.

• Non-metals: Poor conductors, brittle in solid form.

9. Calculating Concentrations

Molarity

• Molarity (M) = moles of solute / volume of solution (L).

Dilutions

• To dilute a solution, add solvent while keeping the number of moles of solute constant.

• Use the formula C1V1 = C2V2 to calculate concentrations before and after dilution.