Comprehensive Cambridge O Level Chemistry Study Notes
Principles of Scientific Enquiry in Chemistry
The Nature of Investigation:
Science involves designing experiments to measure physical quantities like temperature and reaction rate.
Longer investigations examine relationships between variables, such as how the rate of reaction depends on temperature or concentration.
Safety and Risk Assessment:
No chemistry investigation should be performed without teacher approval and a thorough consideration of chemical hazards.
Risk assessments must be done before starting any practical work, potentially requiring the use of fume cupboards or specific ventilation.
Planning and Variables:
Independent Variable: The factor you deliberately change (e.g., temperature).
Dependent Variable: The factor that changes as a result (e.g., volume of gas collected).
Control Variables: Factors kept constant to ensure a fair test (e.g., concentration of reactants).
Data Collection and Reporting:
Accuracy: Using appropriate tools, like a burette instead of a measuring cylinder for higher precision.
Recording: Data should be recorded in tables with clear headings and correct units.
Significant Figures: Values must be recorded to the appropriate degree of precision of the equipment (e.g., a burette can be read to 2 decimal places, to the nearest 0.05 \text{ cm}^3).
Anomalous Results: Values that do not fit the trend should be identified and excluded from average calculations.
States of Matter and Kinetic Particle Theory
Defining the Three States:
Solid: Definite shape and volume at a given temperature. Particles vibrate about fixed positions in a regular lattice.
Liquid: Fixed volume but takes the shape of its container. Particles are close but move randomly and collide.
Gas: No definite shape or volume; spreads to fill the available space. Particles move randomly at high velocities.
Kinetic Particle Theory:
All matter is made of tiny moving particles (atoms, molecules, or ions).
The higher the temperature, the faster particles move on average.
Heavier particles move more slowly than lighter ones at the same temperature.
Changes of State:
Melting: Heating a solid gives particles energy to overcome attractive forces, breaking the regular structure. The temperature remains constant at the melting point.
Boiling/Evaporation: Heating a liquid allows particles to overcome surface forces to form a gas. Boiling occurs when gas bubbles form inside the liquid at the boiling point.
Condensation: Cooling a gas causes particles to move closer and attractive forces to become significant.
Intermolecular Forces: Substances with higher melting/boiling points have stronger forces of attraction between particles.
Expansion/Contraction: Heating increases particle vibration/movement, pushing them apart (expansion), while cooling draws them closer (contraction).
Pressure in Gases:
Caused by gas particles striking the walls of a container. Higher temperature increases particle energy and collision frequency, increasing pressure.
Compelling a gas into a smaller volume (increasing pressure) causes more frequent collisions and increases the temperature of the gas due to frictional forces and bond formation.
Diffusion of Matters
Mechanism: The haphazard and random spreading of gas or liquid particles to fill a space.
Rate of Diffusion:
Gases diffuse faster than liquids because gas particles move more rapidly.
Molecular Mass: Lighter particles (those with lower relative molecular mass) diffuse faster than heavier ones.
Example: Ammonia gas (Mr = 17) diffuses faster than hydrogen chloride gas (Mr = 36.5), forming a white cloud of ammonium chloride closer to the hydrochloric acid source.
Atoms, Elements, and Compounds
Definitions:
Element: A substance made of only one type of atom that cannot be broken down into simpler substances.
Atom: The smallest particle of an element, originally thought to be indivisible (Greek: atomos).
Molecule: A small group of atoms joined together (e.g., O2, P4, or H_2O).
Compound: A pure substance formed when two or more elements chemically combine in fixed proportions (Law of Constant Composition).
Mixture: Contains more than one substance not chemically joined; components retain their properties and can be separated physically.
Sub-atomic Particles:
Proton: Charge +1, Relative Mass 1, found in the nucleus.
Neutron: Charge 0, Relative Mass 1, found in the nucleus.
Electron: Charge -1, Relative Mass \frac{1}{1837}, orbits in shells.
Nuclear Notation:
Proton Number (Z): Number of protons in the nucleus (defines the element).
Mass Number (A): Total number of protons and neutrons.
Neutron Count: \text{Number of neutrons} = A - Z.
Isotopes:
Atoms of the same element with the same number of protons but different numbers of neutrons.
They have the same chemical properties (same electron configuration) but different physical properties like density.
Relative Atomic Mass (A_r): The average mass of an element's isotopes on a scale where {^{12}C} is exactly 12 units.
Formula: A_r = \frac{\text{Average mass of isotopes}}{\frac{1}{12} \text{ of the mass of a } {^{12}C} \text{ atom}}.
Electronic Configuration:
Electrons fill shells closest to the nucleus first.
Capacity: First shell (2), Second shell (8), Third shell (up to 18, but stable at 8).
Noble gases (Group VIII) have full outer shells, making them extremely unreactive.
Chemical Bonding and Structure
Ionic Bonding:
Occurs between metals and non-metals.
Involves the transfer of electrons from metal atoms (forming positive cations) to non-metal atoms (forming negative anions).
Ionic Bond: Strong electrostatic attraction between oppositely charged ions.
Structure: Giant ionic lattice (e.g., Sodium Chloride where each Na^+ is surrounded by 6 \text{ } Cl^-).
Properties: High melting/boiling points, soluble in water, conduct electricity only when molten or in solution.
Covalent Bonding:
Occurs between non-metal atoms.
Involves the sharing of electron pairs to achieve noble gas configurations.
Types: Single (one pair), Double (two pairs, e.g., O=O, CO_2), Triple (three pairs, e.g., N\equiv N).
Simple Molecular Structures: Small molecules with weak intermolecular forces (van der Waals). Low melting points (e.g., I2, CH4).
Giant Covalent Structures: Large networks of atoms held by strong covalent bonds. High melting points (e.g., Diamond, Graphite, SiO_2).
Allotropes of Carbon:
Diamond: Tetrahedral arrangement; every carbon bonded to four others. Hard, non-conductor.
Graphite: Layered structure; each carbon bonded to three others. Layers slide (lubricant). Contains delocalised electrons (conductor).
Metallic Bonding:
Lattice of positive metal ions in a "sea" of delocalised electrons.
Malleability/Ductility: Layers of ions can slide over each other without breaking the metallic bond.
Conduction: Delocalised electrons carry charge.
Redox Reactions and Oxidation Numbers
Oxidation: Gain of oxygen, loss of electrons, or an increase in oxidation number.
Reduction: Loss of oxygen, gain of electrons, or a decrease in oxidation number.
Oxidising Agent: A substance that oxidises another and is itself reduced.
Reducing Agent: A substance that reduces another and is itself oxidised.
Oxidation Number Rules:
Free elements are 0.
Monatomic ions equal their charge.
Sum of oxidation numbers in a neutral compound is 0.
Stoichiometry and Chemical Calculations
The Mole (mol): The amount of substance containing 6.02 \times 10^{23} particles (Avogadro’s constant).
Mass-Mole Calculations:
\text{Mass (g)} = \text{moles} \times \text{molar mass (g/mol)}.
\text{Number of moles} = \frac{\text{Mass}}{\text{Molar mass}}.
Gases: One mole of any gas occupies 24 \text{ dm}^3 at room temperature and pressure (r.t.p.).
\text{Moles} = \frac{\text{Volume in } dm^3}{24}.
Solutions:
\text{Concentration (mol/dm}^3) = \frac{\text{Moles}}{\text{Volume in } dm^3}.
\text{Concentration (g/dm}^3) = \text{Concentration (mol/dm}^3) \times M_r.
Empirical and Molecular Formulae:
Empirical: Simplest whole-number ratio of atoms.
Molecular: Actual number of atoms in a molecule (\text{Molecular mass} / \text{Empirical mass} = n).
Yield and Purity:
\% \text{ Yield} = \frac{\text{Actual yield}}{\text{Theoretical yield}} \times 100.
\% \text{ Purity} = \frac{\text{Mass of pure product}}{\text{Total mass of impure sample}} \times 100.
Electrochemistry
Electrolysis: The decomposition of an ionic compound (molten or in solution) by an electric current.
Anode (+): Non-metals produced; oxidation occurs.
Cathode (-): Metals or hydrogen produced; reduction occurs.
Extraction of Aluminium:
Ore: Bauxite (Al2O3).
Alumina is dissolved in molten cryolite (Na3AlF6) to lower the melting point from 2017^\circ C to approx. 900^\circ C.
Anodes (Graphite) must be replaced regularly as they react with oxygen to form CO_2.
Hydrogen-Oxygen Fuel Cells:
Produce electricity using H2 and O2; only product is water.
Advantage: Pollution-free. Disadvantage: Hydrogen is difficult to store.
Chemical Energetics and Rates
Exothermic Reactions: Transfer thermal energy to surroundings (\Delta H is negative). Examples: Combustion, Neutralisation.
Endothermic Reactions: Take in thermal energy from surroundings (\Delta H is positive). Examples: Photosynthesis, Thermal decomposition.
Bond Energies: \Delta H = \text{Energy to break bonds} - \text{Energy released making bonds}.
Factors Affecting Rate:
Concentration/Pressure: Increases particles per unit volume, leading to more frequent collisions.
Surface Area: Increasing surface area of solids exposes more particles for collision.
Temperature: Increases kinetic energy; more particles have energy > E_a; collisions are more frequent and successful.
Catalyst: Provides alternative route with lower activation energy (E_a).
Reversible Reactions and Equilibrium
Equilibrium: In a closed system, the rate of the forward reaction equals the rate of the reverse reaction.
Le Chatelier's Principle:
Increasing pressure favours the side with fewer gas molecules.
Increasing temperature favours the endothermic reaction.
Increasing concentration of a reactant favours the forward reaction.
Haber Process (Ammonia): N2 + 3H2 \rightleftharpoons 2NH_3. Conditions: 450^\circ C, 20,000 \text{ kPa}, Iron catalyst.
Contact Process (Sulfuric Acid): 2SO2 + O2 \rightleftharpoons 2SO_3. Conditions: 450^\circ C, 200 \text{ kPa}, Vanadium(V) oxide catalyst.
Acids, Bases, and Salts
Theories:
Acids: Proton donors (H^+). Strong acids dissociate completely (e.g., HCl); weak acids partially dissociate (e.g., ethanoic acid).
Bases: Proton acceptors. Alkalis are soluble bases (OH^-).
Oxides:
Acidic: Non-metal oxides (e.g., SO2, CO2).
Basic: Metal oxides (e.g., CuO, CaO).
Amphoteric: React with both acids and bases (e.g., Al2O3, ZnO).
Salt Preparation:
Acid + Excess Metal: For MAZIT metals.
Acid + Insoluble Base/Carbonate: Heating and filtration required.
Titration: For soluble bases (Group I).
Precipitation: Mixing two soluble salts to make an insoluble one (e.g., BaCl2 + Na2SO4 \rightarrow BaSO4(s)).
Periodic Table Trends
Group I (Alkali Metals): Reactivity increases down the group as outer electron is further from the nucleus.
Group VII (Halogens): Reactivity decreases down the group. Darker color and higher density down the group.
Transition Elements: High density, high melting points, variable oxidation states, form colored compounds, act as catalysts.
Group VIII (Noble Gases): Chemically inert due to full outer shells.
Organic Chemistry
Alkanes (CnH{2n+2}): Saturated hydrocarbons. Undergo substitution with chlorine in UV light.
Alkenes (CnH{2n}): Unsaturated (C=C double bond). Undergo addition with bromine (decolourises), hydrogen, and steam (to make alcohols).
Alcohols (R-OH): Made by fermentation or catalytic hydration of ethene. Used as fuels and solvents.
Carboxylic Acids (R-COOH): Weak acids; react with alcohols to form Esters (sweet-smelling compounds).
Polymers:
Addition: Monomers with double bonds join (e.g., poly(ethene)).
Condensation: Join with the loss of a small molecule like water (e.g., Nylon (polyamide), PET (polyester), Proteins (natural polyamides)).