National Year 13 Chemistry

Chapters: 3

1. 3.1 Physical chemistry

3.1.8 Thermodynamics · 3.1.9 Rate equations · 3.1.10 Equilibrium constant Kp for homogeneous systems · 3.1.11 Electrode potentials and electrochemical cells · 3.1.12 Acids and bases

Spontaneity, Entropy and Gibbs Energy – A spontaneous process happens by itself, without outside help (it may be fast or slow). ΔH alone cannot predict it: ice melts on its own although it takes in heat. Entropy S measures how spread out energy and matter are; ΔS = q_rev/T. Second law: for any spontaneous change, ΔS_total = ΔS_system + ΔS_surroundings > 0; at equilibrium it is zero. Third law: a perfect crystal at 0 K has zero entropy. Gibbs energy G = H − TS combines both: at constant T and p, ΔG = ΔH − TΔS, and ΔG < 0 means spontaneous, ΔG = 0 means equilibrium. It links to the equilibrium constant by ΔG° = −RT ln K = −2.303 RT log K.
Integrated Rate Equations: Zero and First Order – Integrated rate equations link concentration with time. Zero order: [R] = [R]₀ − kt (straight line), t½ = [R]₀/2k. First order: k = (2.303/t) log([R]₀/[R]), ln[R] falls in a straight line with slope −k, and t½ = 0.693/k, which does not depend on the starting amount.
Chemical Equilibrium: Kc, Kp, Q and Gibbs Energy – In a closed container a reversible reaction goes both ways. After some time the forward and backward rates become equal, so amounts stop changing, but the reaction does not stop. This is dynamic equilibrium. At equilibrium the ratio of products to reactants (each raised to its coefficient) is a fixed number, the equilibrium constant K. Kc uses concentrations, Kp uses partial pressures, and Kp = Kc(RT)^Δn. Pure solids and liquids are left out of K. The reaction quotient Q tells the direction: Q < K goes forward, Q > K goes backward, Q = K is equilibrium. K and Gibbs energy are linked: ΔG = ΔG° + RT ln Q and ΔG° = −RT ln K.
Galvanic Cells and the Nernst Equation – A galvanic cell turns the energy of a redox reaction into electricity. Oxidation happens at the anode (−) and reduction at the cathode (+). Each electrode has a potential measured against the standard hydrogen electrode (0 V). E°cell = E°cathode − E°anode. The Nernst equation, E = E° − (0.059/n) log Q at 298 K, gives the cell voltage at any concentration, and ΔG = −nFE links voltage to energy.
Ionic Equilibrium: Acids, Bases, Ionisation and pH – Acids and bases can be defined in three ways: Arrhenius (give H⁺ or OH⁻ in water), Brønsted–Lowry (proton donor or acceptor) and Lewis (electron-pair acceptor or donor). Strong acids and bases ionise fully; weak ones ionise only a little and set up an equilibrium with constant Ka or Kb. Water itself ionises: Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C. pH = −log[H⁺] and pH + pOH = 14. For a weak electrolyte, α ≈ √(K/C), so dilution increases α (Ostwald's dilution law). For a conjugate pair, Ka × Kb = Kw. Salts of weak acids or weak bases react with water (hydrolysis) and give non-neutral solutions.

2. 3.2 Inorganic chemistry

3.2.4 Properties of Period 3 elements and their oxides · 3.2.5 Transition metals · 3.2.6 Reactions of ions in aqueous solution

Period 3 Elements and Their Oxides – Across Period 3 (Na, Mg, Al, Si, P, S) the elements burn in oxygen to make oxides. The oxides change from ionic (Na₂O, MgO, Al₂O₃) to giant covalent (SiO₂) to small molecules (P₄O₁₀, SO₂, SO₃). That structure decides their melting point and how they behave with water, acids and bases: basic on the left, amphoteric Al₂O₃ in the middle, acidic on the right.
Transition Elements (d-Block) – Transition elements are the d-block metals of groups 3 to 12. Their last electron goes into an inner (n−1)d orbital. Because the d-orbitals are only partly filled, they show many oxidation states, coloured ions, magnetism (μ = √(n(n+2)) BM), good catalytic power and easy alloy formation.
Metal-Aqua Ions: Reactions in Aqueous Solution – In water, a metal ion is held by six water ligands: [M(H₂O)₆]²⁺ or [M(H₂O)₆]³⁺. A 3+ ion pulls hard on the O–H bonds, so it releases H⁺ and its solution is more acidic. Adding OH⁻, NH₃ or CO₃²⁻ removes H⁺ step by step until a neutral hydroxide precipitates. Al(OH)₃ dissolves again in excess OH⁻ (amphoteric); Cu(OH)₂ dissolves in excess ammonia by ligand substitution. 2+ ions give carbonates with CO₃²⁻, but 3+ ions give hydroxides and CO₂ gas.

3. 3.3 Organic chemistry

3.3.7 Optical isomerism · 3.3.8 Aldehydes and ketones · 3.3.9 Carboxylic acids and derivatives · 3.3.10 Aromatic chemistry · 3.3.11 Amines · 3.3.12 Polymers · 3.3.13 Amino acids, proteins and DNA · 3.3.14 Organic synthesis · 3.3.15 Nuclear magnetic resonance spectroscopy · 3.3.16 Chromatography

Optical Isomerism: Chiral Centres and Enantiomers – Optical isomers are stereoisomers that are mirror images of each other but cannot be placed on top of each other. They happen when a carbon atom has four different groups: a chiral centre. The two forms, called enantiomers, have the same physical and chemical properties except that they rotate plane-polarised light by the same angle in opposite directions, and they can act differently in the body. A 50 : 50 mix (racemic mixture) has no overall effect on the light.
Aldehydes and Ketones: Structure, Preparation and Reactions – Aldehydes and ketones both have a carbonyl group, C=O. In an aldehyde the carbonyl carbon has at least one H on it (−CHO). In a ketone it has two carbon groups (−CO−). The C=O bond is polar: carbon is slightly positive, oxygen slightly negative. So electron-rich particles (nucleophiles) attack the carbon. This one idea explains most of their reactions: addition of HCN, NaHSO₃, alcohols and ammonia derivatives. Aldehydes are easy to oxidise (Tollens and Fehling tests), ketones are not. A hydrogen on the carbon next to C=O (α-H) is slightly acidic, which gives aldol condensation. Aldehydes with no α-H give the Cannizzaro reaction instead.
Carboxylic Acids: Carboxyl Group, Acidity and Reactions – A carboxylic acid has the carboxyl group, −COOH: a C=O and an O−H on the same carbon. It can give away H⁺, and the ion left behind (carboxylate, −COO⁻) is stable because its negative charge is shared equally by two oxygen atoms. That is why these acids are stronger than phenols and alcohols (but weaker than mineral acids like HCl). Groups that pull electrons (Cl, F, NO₂) make the acid stronger; groups that push electrons (alkyl) make it weaker. We make carboxylic acids by oxidising alcohols, aldehydes or alkylbenzenes, by hydrolysing nitriles, esters and amides, or from Grignard reagents and CO₂. Their key reactions: salt formation with NaHCO₃ (CO₂ fizz), esterification, acid chloride and anhydride formation, reduction to alcohols, decarboxylation and the HVZ reaction.
Benzene and Aromaticity – Benzene (C₆H₆) is a flat ring of six sp² carbons. Its six π electrons are spread over the whole ring, so all C–C bonds are equal and the ring is extra stable. A ring like this is called aromatic when it is cyclic, planar, fully conjugated and has 4n + 2 π electrons (Hückel rule). Benzene prefers electrophilic substitution (nitration, halogenation, sulphonation, Friedel–Crafts) over addition, because substitution keeps the stable ring. A group already on the ring decides where the next group goes: ortho/para or meta. Benzene and fused-ring hydrocarbons are toxic and can cause cancer.
Amines: Structure, Naming, Preparation, Basicity and Reactions – An amine is ammonia with one, two or three H atoms swapped for carbon groups. The N keeps a lone pair, so amines are bases and nucleophiles. Count the C atoms on N to get 1°, 2° or 3°. Make them by reducing nitro compounds, nitriles or amides, from alkyl halides + NH₃, by Gabriel (1° only) or Hofmann bromamide (one C fewer). Alkyl groups push electrons and raise basicity; the ring in aniline pulls the pair away and lowers it. Key reactions: salts with acids, acylation, carbylamine test (1°), Hinsberg test (1°/2°/3°), nitrous acid and ring substitution in aniline.
Polymers – A polymer is a giant molecule made by joining many small molecules (monomers) into a long chain. In addition polymerisation, monomers with a C=C double bond open up and link with nothing lost (ethene → poly(ethene)). In condensation polymerisation, two kinds of monomer with reactive groups at both ends join and give off a small molecule such as water at every link (nylon, polyester). Nature makes polymers too: starch, cellulose, proteins, DNA, rubber. Separate chains give thermoplastics that melt and can be recycled; cross-linked chains give thermosets that never melt. Most plastics do not rot, so we must reduce, reuse and recycle them.
Proteins and Enzymes: Amino Acids, Peptide Bond, Structure and Denaturation – Proteins are long chains of amino acids. Each amino acid has an amino group (–NH₂), an acid group (–COOH), a hydrogen and a side group R on one carbon. The –COOH of one amino acid joins the –NH₂ of the next, losing water, to make a peptide bond (–CO–NH–). The order of amino acids is the primary structure; coils and sheets are the secondary structure; the full 3D fold is the tertiary structure; several chains together make the quaternary structure. Heat, acid or alcohol can undo the folds (denaturation). Enzymes are mostly proteins that speed up reactions in the body.
Organic Synthesis – Organic synthesis means making a target molecule from simpler starting materials through a planned series of reactions. Each step changes one functional group and needs the right reagent and conditions (temperature, catalyst, solvent, reflux or distillation). Families of compounds are linked (alkane ↔ alkene → haloalkane → alcohol → aldehyde/ketone → carboxylic acid → ester; haloalkane → amine or nitrile), so we can move between them like stations on a map. Chemists plan backwards from the target (retrosynthesis), prefer short routes with high yield and atom economy, and then carry out the reaction, separate, purify and check the product.
NMR Spectroscopy: Carbon-13 and Proton NMR – Nuclear magnetic resonance (NMR) uses a strong magnet and radio waves to see the different environments of ¹H or ¹³C atoms in a molecule. Each environment gives a peak at its own chemical shift δ (ppm), measured from TMS at δ = 0. In ¹³C NMR, the number of peaks = number of carbon environments. In ¹H NMR, the peak area (integration) gives the ratio of H atoms, and the n + 1 rule tells how many H are on the neighbouring carbon.
Chromatography: Separating a Mixture by How Much It Sticks – Chromatography separates the parts of a mixture. It always has two phases. The stationary phase stays still (paper, a silica layer, a packed column). The mobile phase moves over it (a liquid solvent or a gas). Each substance in the mixture sticks to the stationary phase by a different amount. The one that sticks less moves faster and goes further. In paper chromatography and thin-layer chromatography (TLC) we measure the Rf value: distance moved by the spot divided by distance moved by the solvent front. In column chromatography the substance that sticks least comes out of the bottom first. In gas chromatography (GC) each substance leaves the column at its own retention time, and the area of its peak tells how much is there. We use chromatography to test purity, identify substances and separate them.