Comprehensive Study Guide for IGCSE and O Level Chemistry

States of Matter: Particles and Phase Changes

• The Kinetic Particle Theory: All matter consists of very tiny pieces called particles (atoms, molecules, or ions) that are in constant motion. The arrangement and movement of these particles define the state of a substance.
• Evidence for Particles:
  • Brownian Motion: First observed by Robert Brown in 1827 with pollen granules in water; explained by Albert Einstein in 1905 as the result of bombardment by invisible water particles.
  • Diffusion: The process by which particles mix and spread by colliding with others and bouncing in all directions, moving from areas of high concentration to low concentration. Observed in cooking smells, dust dancing in sunlight, or the spreading of potassium manganate(VII) in water.
• Three States of Matter:
  • Solids: Have a fixed shape and volume; particles are arranged in a regular lattice and only vibrate about fixed positions; transition to liquid at a fixed melting point.
  • Liquids: Flow to take the shape of their container; fixed volume; particles are close together but can slide past each other.
  • Gases: No fixed shape or volume; particles are far apart and move rapidly and randomly; exerting pressure by colliding with container walls.
• Phase Changes:
  • Melting: Solid to liquid. Energy is taken in to break the lattice structure.
  • Boiling: Liquid to gas throughout the liquid at a specific temperature.
  • Evaporation: Liquid to gas at the surface over a range of temperatures below the boiling point.
  • Condensing: Gas to liquid.
  • Freezing: Liquid to solid. The freezing point of water and melting point of ice are both $0\,^{\circ}C$.
• Heating and Cooling Curves: Graphs plotting temperature against time. Horizontal lines indicate a change of state where temperature remains constant as energy is used to break forces between particles (melting/boiling) or released as particles move closer (freezing/condensing).
• Gas Dynamics:
  • Effect of Temperature: Heating a gas increases its volume at constant pressure as particles move faster and hit walls with more force.
  • Effect of Pressure: Increasing pressure decreases gas volume by pushing particles closer together.
  • Rate of Diffusion: Depends on particle mass. Lighter particles (lower relative molecular mass, $M_r$) diffuse faster. For example, ammonia ($M_r = 17$) diffuses faster than hydrogen chloride ($M_r = 36.5$).

Atomic Structure and the Periodic Table

• Elements, Compounds, and Mixtures:
  • Element: Contains only one kind of atom (e.g., Sodium $\text{Na}$, Carbon $\text{C}$).
  • Compound: Different elements chemically bonded together (e.g., Water $\text{H}_2\text{O}$, Carbon Dioxide $\text{CO}_2$).
  • Mixture: Different substances physically mixed but not bonded; easily separated.
• Subatomic Particles:
  • Proton: Mass of $1\,\text{amu}$, positive charge ($1+$).
  • Neutron: Mass of $1\,\text{amu}$, no charge.
  • Electron: Negligible mass (taken as $0$), negative charge ($1-$).
• Atomic Definitions:
  • Proton Number (Atomic Number): Number of protons in the nucleus; identifies the element.
  • Nucleon Number (Mass Number): Total number of protons and neutrons ($\text{protons} + \text{neutrons} = \text{nucleons}$).
  • Isotopes: Different atoms of the same element with the same number of protons but different numbers of neutrons (e.g., $^{12}\text{C}$, $^{13}\text{C}$, $^{14}\text{C}$). Isotopes have identical chemical properties because they have the same number of outer shell electrons.
• Electron Arrangement: Electrons occupy shells (energy levels) around the nucleus:
  • First shell: max $2$ electrons.
  • Second shell: max $8$ electrons.
  • Third shell: max $8$ electrons (when it is the outer shell).
• Periodic Table Trends:
  • Groups: Vertical columns. Group number = number of outer shell electrons.
  • Periods: Horizontal rows. Period number = number of electron shells occupied.
  • Group VIII (Noble Gases): Have full outer shells; stable and unreactive (monatomic).

Chemical Bonding and Structures

• Ionic Bonding: Formation of ions by transferring electrons from metals to non-metals to achieve full outer shells.
  • Ionic Bond: Strong electrostatic attraction between oppositely charged ions.
  • Giant Lattice: regular arrangement of alternating positive and negative ions.
  • Properties: High melting/boiling points; conduct electricity when molten or aqueous (ions are free to move); usually soluble in water.
• Covalent Bonding: Formed when a pair of electrons is shared between two non-metal atoms.
  • Simple Molecules: Small groups of atoms held by covalent bonds (e.g., $\text{H}_2$, $\text{Cl}_2$, $\text{H}_2\text{O}$, $\text{CH}_4$, $\text{CO}_2$).
  • Properties: Low melting/boiling points due to weak intermolecular forces; poor electrical conductivity.
• Giant Covalent Structures:
  • Diamond: Each carbon atom bonds to four others in a tetrahedral lattice; extremely hard; high melting point ($3550\,^{\circ}C$).
  • Graphite: Each carbon atom bonds to three others in layers; soft/slippery (layers slide); conducts electricity due to one delocalised electron per atom.
  • Silicon(IV) Oxide ($\text{SiO}_2$): Similar structure and properties to diamond.
• Metallic Bonding: Electrostatic attraction between positive metal ions in a giant lattice and a 'sea' of delocalised electrons.
  • Properties: Malleable/ductile (layers slide without breaking bonds); good conductors of heat and electricity.

Stoichiometry and Calculations

• Relative Masses:
  • Relative Atomic Mass ($A_r$): Average mass of an element's isotopes relative to $\frac{1}{12}$ the mass of a carbon-12 atom.
  • Relative Molecular/Formula Mass ($M_r$): Sum of all $A_r$ values in the formula.
• The Mole Concept:
  • Mole (mol): $6.02 \times 10^{23}$ particles (Avogadro constant).
  • $Amount\,(mol) = \frac{mass\,(g)}{molar\,mass\,(g/mol)}$.
• Chemical Equations:
  • State Symbols: $(s)$ solid, $(l)$ liquid, $(g)$ gas, $(aq)$ aqueous.
  • Conservation of Mass: Total mass of reactants = total mass of products.
• Gas and Solution Calculations:
  • Molar Gas Volume: $24\,dm^3$ at room temperature and pressure (r.t.p.).
  • Concentration ($mol/dm^3$): $Concentration = \frac{amount\,(mol)}{volume\,(dm^3)}$.
• Formulae and Ratios:
  • Empirical Formula: Simplest whole-number ratio of atoms in a compound.
  • Molecular Formula: Actual number of atoms of each element in one molecule.
  • Percentage Yield: $(\%\,yield) = \frac{actual\,mass}{theoretical\,mass} \times 100\%$.
  • Percentage Purity: $(\%\,purity) = \frac{pure\,mass}{total\,mass} \times 100\%$.

Electricity and Chemical Change

• Electrolysis: Decomposition of an ionic compound by an electric current.
  • Setup: Anode (positive), Cathode (negative), Electrolyte (molten or aqueous liquid).
  • Rules for Aqueous Solutions:
    • Cathode: $\text{H}^+$ or the least reactive metal ion is discharged (reduced).
    • Anode: Halide ions are discharged in concentrated solutions; otherwise, oxygen forms from $OH^-$ ions.
  • Samples:
    • Molten $PbBr_2$: lead at cathode, bromine at anode.
    • Brine (Conc. $NaCl$): hydrogen at cathode, chlorine at anode, sodium hydroxide remains.
• Refining Copper: Impure copper anode dissolves; pure copper builds up on cathode; sludge contains precious metals.
• Electroplating: Coating a metal object with another metal to improve appearance or corrosion resistance. Object is the cathode; metal to plate is the anode; electrolyte is a soluble salt of the plating metal.

Chemical Energetics and Redox

• Energy Changes:
  • Exothermic: Transfers thermal energy to surroundings (temp increases). $\Delta H$ is negative.
  • Endothermic: Takes in thermal energy (temp decreases). $\Delta H$ is positive.
  • Activation Energy ($E_a$): Minimum collision energy required for reaction.
  • Bond Energies: Breaking bonds is endothermic; making bonds is exothermic. $\Delta H = \text{Energy in (breaking)} - \text{Energy out (making)}$.
• Redox Reactions:
  • Oxidation: Gain of oxygen or loss of electrons (OIL).
  • Reduction: Loss of oxygen or gain of electrons (RIG).
  • Oxidation Number: Shows how many electrons an atom has lost, gained, or shared. Group I = $+I$, Group II = $+II$, Oxygen = $-II$, Chlorine = $-I$. Rise in ox. no. = oxidation; Fall = reduction.
  • Agents: Oxidising agents oxidise others and are reduced themselves (e.g., Acidified $KMnO_4$ turns purple to colourless). Reducing agents reduce others and are oxidised (e.g., $KI$ turns colourless to red-brown).

Acids, Bases, and Salts

• Acids: Proton ($H^+$) donors. Strong acids (e.g., $HCl$) dissociate completely; weak acids (e.g., $CH_3COOH$) dissociate partially.
  • Reactions: Acid + metal $\rightarrow$ salt + $H_2$; Acid + base $\rightarrow$ salt + $H_2O$; Acid + carbonate $\rightarrow$ salt + $H_2O + CO_2$.
• Bases and Alkalis: Metal oxides/hydroxides are bases. Soluble bases are alkalis ($OH^-$ sources). Bases are proton acceptors.
• Indicators:
  • Litmus: Red in acid, blue in alkali.
  • Methyl Orange: Red in acid, yellow in alkali.
  • Thymolphthalein: Colourless in acid, blue in alkali.
• pH Scale: $0-6$ (acidic), $7$ (neutral), $8-14$ (alkaline).
• Salt Preparation:
  • Soluble Salts: React acid with excess metal/base/carbonate then filter and crystallise; or titration with alkali.
  • Insoluble Salts: Precipitation by mixing two soluble salt solutions.

Rates of Reaction and Equilibrium

• Collision Theory: For particles to react, they must collide with energy $\ge E_a$. Rate increases with concentration, gas pressure, temperature, and surface area (powdered solid).
• Catalysts: Speed up reactions by lowering $E_a$. Enzymes are biological catalysts (proteins).
• Reversible Reactions: Some reactions reach dynamic equilibrium in a closed system where forward/reverse rates are equal.
  • Le Chatelier's Principle: System opposes external changes.
    • Temp increase: Favours endothermic side.
    • Pressure increase: Favours side with fewer gas molecules.
  • Haber Process (Ammonia): $N_2 + 3H_2 \rightleftharpoons 2NH_3$. Conditions: $450\,^{\circ}C$, $200\,atm$, Iron catalyst.
  • Contact Process (Sulfuric Acid): $2SO_2 + O_2 \rightleftharpoons 2SO_3$. Conditions: $450\,^{\circ}C$, $2\,atm$, Vanadium(V) oxide catalyst.

Organic Chemistry

• Petroleum Fractions: Refinery gas (heating), Gasoline (petrol), Naphtha (chemicals), Kerosene (jet fuel), Diesel/Gas oil (trucks), Fuel oil (ships), Lubricating oil, Bitumen (roads).
• Alkanes: Saturated hydrocarbons ($C_nH_{2n+2}$). Substitution reaction with chlorine in UV light.
• Alkenes: Unsaturated hydrocarbons ($C_nH_{2n}$) with $C=C$. Test: Bromine water turns orange to colourless. Addition reactions with steam (alcohols), hydrogen (alkanes), and bromine.
• Alcohols: $C_nH_{2n+1}OH$. Ethanol made by fermentation ($30\,^{\circ}C$, yeast, no air) or addition of steam to ethene ($300\,^{\circ}C$, $60\,atm$, phosphoric acid catalyst).
• Carboxylic Acids: $C_nH_{2n+1}COOH$. Ethanoic acid made by oxidation of ethanol. React with alcohols to form Esters (ester linkage $-COO-$).
• Polymers:
  • Addition: Thousands of alkene monomers join (e.g., Polyethene, PVC, PTFE).
  • Condensation: Two different monomers join with elimination of small molecules (e.g., Nylon-polyamide, PET-polyester).
  • Proteins: Natural polyamides built from amino acid monomers ($NH_2$ and $COOH$ groups).