Cambridge IGCSE Chemistry Comprehensive Study Guide

THE NATURE OF MATTER AND STATES OF MATTER

• Matter covers all substances and materials of which the universe is composed. It has two common properties: it occupies space (volume) and it has mass.

• States of Matter: Any chemical substance can exist in three different forms depending on conditions: solid, liquid, or gas.

• Physical Properties Table (Differences in three states):

  • Solid: Has a fixed volume; high density; has a definite shape; does not flow.

  • Liquid: Has a fixed volume; moderate-to-high density; no definite shape (takes shape of container); generally flows easily.

  • Gas: No fixed volume (expands to fill the container); low density; no definite shape; flows easily.

• Fluids: Liquids and gases are fluids, meaning they can be poured or pumped from one container to another.

• Response to Pressure and Temperature:

  • All states show expansion (volume increase) when temperature increases and contraction (volume decrease) when temperature lowers. The effect is significantly larger for gases.

  • Compressibility: Gases are easily compressed. Liquids are only slightly compressible. The volume of a solid is unaffected by changing pressure.

CHANGES IN PHYSICAL STATE

• Changes of state can occurs by raising or lowering the temperature or changing the atmospheric pressure.

• Melting and Freezing:

  • Melting Point (m.p.): The temperature at which a substance turns to a liquid. Pure substances have sharp melting points.

  • Freezing Point (f.p.): The reverse process taking place sharply at the same temperature as the melting point (e.g., pure water melts and freezes at $0\,°C$).

• Sublimation: A direct change of state from solid to gas, or gas to solid, bypassing the liquid phase (e.g., solid carbon dioxide or 'dry ice' sublimes at $-78\,°C$).

• Evaporation, Boiling, and Condensation:

  • Evaporation: A process occurring at the surface of a liquid, involving the change from liquid to vapour at temperatures below the boiling point. Rate increases with larger surface area and higher temperatures.

  • Boiling: Occurs at a specific boiling point when gas bubbles form within the liquid. The boiling point is specific to a pure liquid when gas pressure equals atmospheric pressure ($100\,°C$ for water at standard pressure).

  • Volatile: Describes a liquid that evaporates easily due to weak intermolecular forces (e.g., ethanol boils at $78\,°C$).

  • Condensation: The change of a vapour or gas into a liquid, during which heat is given out.

• Purity and Impurities:

  • Pure Substance: One chemical element or compound with definite melting and boiling points.

  • Impurities: Contaminants change the value of melting or boiling points. Seawater (impure water) freezes below $0\,°C$ and boils above $100\,°C$. Impure substances often melt/boil over a range of temperatures.

KINETIC PARTICLE THEORY OF MATTER

• Modern understanding posits all matter consists of very small particles known as atoms.

• Key ideas of the theory:

  • All matter is made of tiny particles (atoms, molecules, or ions).

  • Particles are in constant motion; higher temperature correlates to higher average energy.

  • Intermolecular Space: Space between atoms or molecules (smallest in solids, largest in gases).

  • Intermolecular Forces: Weak attractive forces acting between molecules.

• State-Specific Particle Organization:

  • Solid: Particles packed close in a regular lattice; vibrate about fixed positions; cannot move freely.

  • Liquid: Particles closely packed but in an irregular arrangement; able to move past each other.

  • Gas: Particles arranged totally irregularly; spread very far apart; move randomly at high speeds.

• Energy Changes:

  • Endothermic Changes: Melting, evaporation, and boiling require energy input to overcome forces between particles. $\Delta H$ has a positive value.

  • Exothermic Changes: Condensation and freezing release energy as particles come closer and new forces form. $\Delta H$ has a negative value.

MIXTURES, SOLUTIONS, AND DIFFUSION

• Mixture: Two or more substances mixed but not chemically combined. They can be separated by physical means.

• Solution: Formed when a solute dissolves into a solvent.

  • Solute: The solid substance that dissolves.

  • Solvent: The liquid that dissolves the solute (water is common; others are organic solvents like ethanol or propanone).

  • Saturated Solution: Contains as much dissolved solute as possible at a specific temperature.

• Solubility: Measure of how much solute dissolves in a solvent at a given temperature. Generally increases with temperature for solids, but decreases for gases. Gas solubility increases with pressure.

• Diffusion: The process by which different fluids mix due to the random motion of particles.

  • Particles move from high to low concentration until evenly spread.

  • Rate is slower in liquids than gases; does not happen in solids.

  • Molecular Mass Effect: Lighter gas molecules (lower relative molecular mass $M_r$) move faster and diffuse more quickly than heavier ones. Example: Ammonia ($M_r = 17$) diffuses faster than Hydrochloric acid ($M_r = 36.5$).

ATOMIC STRUCTURE AND SUBATOMIC PARTICLES

• Atoms: The smallest particles of an element that participate in chemical reactions.

• Elements: Substances made of only one type of atom (118 known; 94 occur naturally).

• Subatomic Particles Table:

  • Proton: Relative mass = $1$; Relative charge = $+1$; Location = Nucleus.

  • Neutron: Relative mass = $1$; Relative charge = $0$; Location = Nucleus.

  • Electron: Relative mass = $1/1840$ (negligible); Relative charge = $-1$; Location = Shells outside nucleus.

• Atomic Symbols:

  • Proton Number (Z) (Atomic Number): The number of protons in the nucleus. Defines the element.

  • Mass Number (A) (Nucleon Number): Total number of protons and neutrons.

  • Representation format: ${}^{A}_{Z}X$.

  • Relationship: Number of neutrons = $A - Z$.

  • Atoms are electrically neutral because Number of Protons = Number of Electrons.

ISOTOPES AND RELATIVE ATOMIC MASS

• Isotopes: Atoms of the same element with the same proton number but different nucleon numbers (different numbers of neutrons).

• Properties: Isotopes have the same chemical properties (same electron configuration) but different physical properties (e.g., density, diffusion rate).

• Radioisotopes: Unstable isotopes that break up spontaneously emitting radiation (e.g., Tritium or carbon-$14$).

• Relative Atomic Mass ($A_r$): Average mass of naturally occurring atoms of an element on a scale where carbon-$12$ is exactly $12$ units.

• Formula for $A_r$ Calculation:   ((\text{Mass of Isotope 1} \times \text{Abundance %}) + (\text{Mass of Isotope 2} \times \text{Abundance %})) / 100

  • Example for Chlorine: $(35 \times 75) + (37 \times 25) = 3550 / 100 = 35.5$.

ELECTRONIC CONFIGURATION

• Electron Shells (Energy Levels): Orbitals around the nucleus filled in specific orders:

  • First level: holds a maximum of $2$ electrons.

  • Second level: holds a maximum of $8$ electrons.

  • Third level: often holds $8$ to achieve stable 'noble gas' configuration.

• Periodic Table Relationship:

  • Group Number: Number of electrons in the outer shell.

  • Period Number: Number of occupied electron shells.

• Noble Gases (Group VIII): Very unreactive because they have a full outer shell (usually $8$ electrons; helium has $2$).

CHEMICAL BONDING TYPES

• Compounds: Two or more elements chemically combined in fixed proportions.

• Covalent Bonding: Formed by sharing pairs of electrons between two atoms (usually non-metals), leading to noble gas configurations.

  • Simple Molecular Elements: Diatomic molecules ($H_2$, $Cl_2$, $O_2$, $N_2$).

  • Covalent Properties: Low melting/boiling points (weak intermolecular forces); poor electrical conductivity (no free ions/electrons).

• Ionic Bonding: Strong electrostatic force of attraction between oppositely charged ions (metals and non-metals).

  • Cation: Positive ion (atom loses electrons).

  • Anion: Negative ion (atom gains electrons).

  • Ionic Properties: High melting/boiling points (strong lattice); soluble in water; conduct electricity when molten or aqueous (free ions).

• Giant Structures:

  • Giant Ionic Lattice: Alternating cations and anions ($NaCl$).

  • Giant Covalent: Diamond (tetrahedral, all valence electrons used, hard, no conductivity), Graphite (layered hexagons, three bonds per carbon, slippery, conducts electricity via free electrons), Silicon(IV) oxide ($SiO_2$, similar to diamond).

  • Metallic Bonding: Electrostatic attraction between positive ions in a regular lattice and a 'sea' of delocalised electrons. Explains electrical/thermal conductivity, malleability, and ductility.

CHEMICAL FORMULAE AND EQUATIONS

• Valency: Combining power based on group number.

• Formulae:

  • Molecular Formula: Actual number of atoms in a molecule.

  • Empirical Formula: Simplest whole number ratio of atoms in a compound.

• Balanced Equations: Total mass of reactants equals total mass of products (Law of Conservation of Mass).

• State Symbols: $(s)$ solid, $(l)$ liquid, $(g)$ gas, $(aq)$ aqueous solution.

• Ionic Equations: Simplified equations showing only particles that take part in the reaction (excluding spectator ions).

• Relative Masses:

  • Relative Molecular Mass ($M_r$): Sum of $A_r$ of all atoms in a molecule.

  • Relative Formula Mass ($M_r$): Used for ionic compounds (sum of $A_r$ in the formula unit).

THE MOLE AND STOICHIOMETRY

• The Mole (mol): Unit of amount containing the Avogadro constant ($6.02 \times 10^{23}$) of particles.

• Molar Mass: mass of $1$ mole of a substance (in $g/mol$).

• Fundamental Equations:

  • $n = \text{mass}(g) / \text{Molar Mass}(g/mol)$.

  • For Gases: $n = \text{volume}(dm^3) / 24\,dm^3/mol$ at r.t.p.

  • For Solutions: $n = (\text{concentration} \times \text{volume}) / 1000$ (volume in $cm^3$).

• Reaction Metrics:

  • Percentage Yield: $(\text{Actual Yield} / \text{Predicted Yield}) \times 100$.

  • Percentage Purity: $(\text{Mass of Pure product} / \text{Mass of Impure product}) \times 100$.

  • Limiting Reactant: The reactant that is completely consumed first, determining the maximum product yield.

ELECTROCHEMISTRY

• Electrolysis: Breakdown of an ionic compound, molten or aqueous, by electricity.

• The Cell:

  • Cathode: Negative electrode (attracts cations, reduction takes place).

  • Anode: Positive electrode (attracts anions, oxidation takes place).

• Electrode Products:

  • Molten binary salts: Metal at cathode, non-metal at anode.

  • Aqueous solutions:

  • At cathode: Metal or Hydrogen ($H_2$ is produced if the metal is more reactive than hydrogen).

  • At anode: Non-metal other than hydrogen (Halogen preferred over Oxygen if concentrated; otherwise Oxygen).

• Half-Equations Example (molten $PbBr_2$):

  • Cathode: $Pb^{2+} + 2e \rightarrow Pb$.

  • Anode: $2Br^- \rightarrow Br_2 + 2e$.

• Electroplating: Using electrolysis to coat an object with metal. Object = cathode; Anode = plating metal; Electrolyte = salt of plating metal.

• Hydrogen-Oxygen Fuel Cells: Generates electricity via reaction: $2H_2(g) + O_2(g) \rightarrow 2H_2O(l)$. Efficient and clean but hydrogen storage is difficult.

CHEMICAL ENERGETICS

• Exothermic Reactions: Release thermal energy ($\Delta H$ is negative). Surroundings get warmer (e.g., combustion).

• Endothermic Reactions: Absorb thermal energy ($\Delta H$ is positive). Surroundings get colder (e.g., photosynthesis).

• Bond Energies:

  • Bond breaking = Endothermic.

  • Bond making = Exothermic.

  • $\Delta H = (\text{Energy needed to break bonds}) - (\text{Energy released when bonds made})$.

• Activation Energy ($E_a$): Minimum energy required for a reaction to proceed.

RATES OF REACTION

• Factors Affecting Rate:

  • Surface Area: Larger area (powdered solids) increases collision frequency.

  • Concentration/Pressure: More particles in a unit volume increase collision frequency.

  • Temperature: Increases kinetic energy of particles, leading to more frequent and successful (energy > E_a) collisions.

  • Catalyst: Lowers activation energy by providing an alternative pathway. Not used up.

• Enzymes: Specific protein catalysts in biological organisms.

REVERSIBLE REACTIONS AND EQUILIBRIUM

• Reversible Reactions: Marked by $\rightleftharpoons$. Can go forwards or backwards.

• Dynamic Equilibrium: Rates of forward and reverse reactions are equal in a closed system; concentrations remain constant.

• Le Chatelier’s Principle: System shifts to oppose changes.

  • Temp increase: Favours endothermic dir.

  • Pressure increase: Favours side with fewer gas molecules.

• Industrial Applications:

  • Haber Process: $N_2 + 3H_2 \rightleftharpoons 2NH_3$. Conditions: $450\,°C$, $20000\,kPa$, iron catalyst.

  • Contact Process: $2SO_2 + O_2 \rightleftharpoons 2SO_3$. Conditions: $450\,°C$, $200\,kPa$, vanadium(V) oxide catalyst.

• Fertilisers: NPK providing Nitrogen, Phosphorus, Potassium.

ACIDS AND BASES

• Definitions:

  • Acid: Proton ($H^+$) donor. pH < 7.

  • Base: Proton ($H^+$) acceptor. pH > 7. Alkalis are soluble bases.

• Acid Reactions:

  • Metal + Acid $\rightarrow$ Salt + Hydrogen.

  • Base + Acid $\rightarrow$ Salt + Water (Neutralisation).

  • Carbonate + Acid $\rightarrow$ Salt + Water + Carbon Dioxide.

• Indicators:

  • Litmus: Red in acid, Blue in alkali.

  • Thymolphthalein: Colourless in acid, Blue in alkali.

  • Methyl Orange: Red in acid, Yellow in alkali.

• Oxides:

  • Basic oxides: metal oxides ($MgO, CuO$).

  • Acidic oxides: non-metal oxides ($SO_2, CO_2$).

  • Amphoteric oxides: react with acids and bases ($Al_2O_3, ZnO$).

  • Neutral oxides: ($H_2O, CO$).

CHEMICAL ANALYSIS TESTS

• Cations:

  • Li: Red flame.

  • Na: Yellow flame.

  • K: Lilac flame.

  • Ca: Orange-red flame.

  • Ba: Green flame.

  • Cu: Blue-green flame.

  • Aqueous ($+ NaOH$/$NH_3$): $Fe^{2+}$ (green ppt), $Fe^{3+}$ (red-brown ppt), $Cu^{2+}$ (blue ppt).

• Anions:

  • Carbonate: add acid $\rightarrow$ $CO_2$ (test with limewater).

  • Halides: add $AgNO_3$ $\rightarrow$ $Cl^-$ (white), $Br^-$ (cream), $I^-$ (yellow).

  • Sulfate: add $Ba(NO_3)_2$ $\rightarrow$ White ppt ($BaSO_4$).

• Gases:

  • $H_2$: lighted splint (pop).

  • $O_2$: glowing splint (relights).

  • $Cl_2$: damp litmus (bleached).

  • $NH_3$: damp red litmus (blue).

  • $SO_2$: acidified $KMnO_4$ (purple to colourless).

ORGANIC CHEMISTRY INTRODUCTION

• Homologous Series: Same functional group, general formula, and similar properties.

• Alkanes ($C_nH_{2n+2}$): Saturated hydrocarbons. Methane to Butane are gases ($n=1..4$). Substitutes with chlorine in UV light.

• Alkenes ($C_nH_{2n}$): Unsaturated (C=C). Undergo addition reactions. Discolourise bromine water.

• Alcohols ($C_nH_{2n+1}OH$): Ethanol manufactured by fermentation or addition of steam to ethene.

• Carboxylic Acids ($C_nH_{2n+1}COOH$): Weak acids; e.g., ethanoic acid in vinegar.

• Esters: Formed from alcohol and acid ($RCOOR'$) with loss of water. Used for flavorings.

• Polymers:

  • Addition polymerisation: opening C=C bonds (Poly(ethene)).

  • Condensation polymerisation: loss of small molecule ($H_2O$). Polyamides (nylon), polyesters (PET), and proteins (natural polyamides formed from 20 amino acids).

ENVIRONMENTAL CHEMISTRY

• Air Pollution:

  • $SO_2$ (fossil fuels) and NOx (car engines) cause acid rain.

  • $CO$ (incomplete combustion) is toxic.

  • Particulates irritate lungs.

• Greenhouse Gases: $CO_2$ and Methane ($CH_4$) trap infrared radiation causing global warming and climate change.

• Water Treatment: Screening, sedimentation, filtration, and chlorination for safety.

• Pollutants in water: Microplastics, sewage, nitrates/phosphates (cause algae blooms), and toxic heavy metals ($Pb, Hg$).