Atomic structure

Atoms have a nucleus (protons + neutrons) surrounded by electrons in shells, sub-shells and orbitals.

• Relative mass: proton = 1, neutron = 1, electron = 1/1836 (nearly zero).
• Relative charge: proton = +1, neutron = 0, electron = -1.
• Atomic number (Z) = number of protons. Mass number (A) = protons + neutrons.
• Isotopes have the same number of protons but different numbers of neutrons, so same chemical properties but different physical properties (mass).

Electron configuration

Sub-shells fill in order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p...

• Sub-shell capacities: s = 2, p = 6, d = 10, f = 14 electrons.
• 4s fills before 3d, but 3d electrons are removed before 4s when transition metals form ions (a very common exam trap).
• Orbitals within a sub-shell each hold one electron before any pairs up, and paired electrons have opposite spins (Hund's rule and the Pauli exclusion principle).

Ionisation energy

First ionisation energy is the energy needed to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.

• Successive ionisation energies always increase, because removing an electron leaves the remaining ones more tightly held by the same nuclear charge.
• Big jumps in successive ionisation energies show a change of shell (electron shielding drops sharply).
• Across a period, ionisation energy generally rises (more protons, same shell, less shielding) but with dips at groups 2 to 3 (extra sub-shell, e.g. losing a p electron before s) and 5 to 6 (electron pairing repulsion).
• Down a group, ionisation energy falls: extra shells increase shielding and atomic radius, outweighing the bigger nuclear charge.

Bonding

Three main types:

• Ionic: electrostatic attraction between oppositely charged ions in a giant lattice; strong, high melting points, conducts only when molten or dissolved.
• Covalent: shared pair of electrons between atoms; can be simple molecular (low melting point) or giant covalent (very high melting point, e.g. diamond, silicon dioxide).
• Metallic: delocalised electrons attract a lattice of positive metal ions; conducts electricity, malleable.

Shapes of molecules

Electron pair repulsion theory: electron pairs (bonding and lone) arrange themselves to be as far apart as possible.

• Lone pairs repel more strongly than bonding pairs, so bond angles shrink by about 2.5 degrees per lone pair.
• Learn the key shapes and angles: linear 180 degrees, trigonal planar 120 degrees, tetrahedral 109.5 degrees, trigonal bipyramidal 90 and 120 degrees, octahedral 90 degrees.

Common mistakes

• Writing 3d before 4s in electron configurations (wrong order for filling, though 4s empties first on ionisation).
• Forgetting lone pairs when predicting shape and bond angle.
• Confusing electronegativity trends with ionisation energy trends.

Enthalpy basics

• Enthalpy change (ΔH) is measured in kJ mol⁻¹. Exothermic = ΔH negative (energy released), endothermic = ΔH positive (energy absorbed).
• Standard conditions for ΔH values: 100 kPa, a stated temperature (usually 298 K), and all substances in their standard states.
• Key definitions to quote exactly: standard enthalpy of formation (1 mole of compound from elements in standard states), standard enthalpy of combustion (1 mole of substance burned completely in oxygen).
• Bond enthalpies are always positive (breaking bonds needs energy); ΔH = (bonds broken) − (bonds formed).

Calculating enthalpy changes

• Hess's Law: the enthalpy change for a reaction is the same regardless of route, because energy is conserved. Draw the cycle and add arrows.
• For combustion cycles, ΔH = ΣΔHf(reactants) − ΣΔHf(products).
• For formation cycles, ΔH = ΣΔHc(reactants) − ΣΔHc(products) (note this one flips).
• Calorimetry: q = mcΔT, then divide by moles to get ΔH per mole. Common mistake: forgetting to convert mass of solution properly or using the wrong mass (use mass of solution, not just solute).

Kinetics

• Rate increases with temperature because more particles have energy ≥ activation energy (Ea) — the Maxwell-Boltzmann distribution shifts right and the area beyond Ea grows a lot for a small temperature rise.
• A catalyst provides an alternative pathway with lower Ea, increasing the proportion of successful collisions without being used up.
• Increasing concentration or pressure increases collision frequency, increasing rate.
• Rate = k[A]^m[B]^n — orders must be found experimentally, never assumed from the equation. Overall order = m + n.
• Rate constant k varies with temperature only (Arrhenius equation, k = Ae^(−Ea/RT)); catalysts and concentration do not change k.

Equilibria

• Dynamic equilibrium: forward and reverse rates are equal, concentrations stay constant, only in a closed system.
• Le Chatelier's Principle: system opposes any change imposed on it. Common mistakes: forgetting equilibrium shifts don't change Kc/Kp value (only temperature does), and confusing 'shift' with 'change in Kc'.
• Kc has units depending on the stoichiometry — always derive them, don't assume mol dm⁻³.
• Increasing temperature always increases the rate constant for both forward and reverse reactions, but favours the endothermic direction more, shifting the position of equilibrium.
• A catalyst speeds up both forward and reverse reactions equally, so it has no effect on the position of equilibrium or Kc — it only helps equilibrium get established faster.
• For gaseous equilibria, increasing pressure shifts the position towards the side with fewer gas moles.

Oxidation states

Oxidation state rules: uncombined elements = 0; simple ions = charge of the ion; oxygen = -2 (except peroxides, -1, and OF2, +2); hydrogen = +1 (except metal hydrides, -1); overall charge of a neutral compound = 0, of an ion = the ion's charge.

OIL RIG

Oxidation Is Loss of electrons, Reduction Is Gain of electrons. A reducing agent gets oxidised (loses electrons, donates them); an oxidising agent gets reduced (gains electrons, accepts them). Redox reactions can be split into two half-equations which must balance for charge and atoms, then combined by cancelling electrons.

Half-equations

Balance atoms other than O and H first, then balance O using H2O, then balance H using H+, then balance charge using electrons. When combining two half-equations, scale them so electrons lost equal electrons gained before adding together.

Electrode potentials

Standard electrode potential, E°, is measured against the standard hydrogen electrode (E° = 0.00 V exactly), under standard conditions: 298 K, 100 kPa, 1 mol dm⁻³ solutions. A standard cell uses a salt bridge (often containing KNO3 or KCl, never KCl with silver ions present as it would precipitate AgCl) to complete the circuit without mixing solutions.

Calculating cell EMF

EMF = E°(positive electrode) − E°(negative electrode), or E°(cathode) − E°(anode). The more positive E° value is the cathode (reduction happens there in the cell); the more negative is the anode (oxidation happens there). A more positive E° means a species is a better oxidising agent (more easily reduced).

Predicting feasibility

A reaction is thermodynamically feasible if the overall EMF is positive when combining the two half-equations correctly (oxidising agent's E° minus reducing agent's E°). Remember this only predicts feasibility, not whether the reaction actually happens at a measurable rate — kinetic barriers can prevent a feasible reaction occurring.

Common mistakes

Don't confuse oxidation state (a bookkeeping number) with actual ionic charge. Don't forget to check both half-equations balance for charge before combining. Don't assume a positive EMF guarantees a fast reaction — it only tells you about feasibility, not rate. When writing cell diagrams, the convention is anode (oxidation, left) | | cathode (reduction, right), with the more negative electrode on the left.

Storage and fuel cells

Rechargeable cells (like Li-ion) can be recharged because the cell reaction is reversible by applying an external EMF. Fuel cells (like hydrogen-oxygen) continuously convert chemical energy directly to electrical energy while fuel and oxidant are supplied, producing water as the only product in acidic conditions.

Periodicity: structure across Period 3

Across Period 3 (Na to Ar), atomic radius decreases and first ionisation energy generally increases, because nuclear charge increases while shielding stays roughly constant. There are small drops at Al (starts filling 3p, slightly higher energy sub-shell) and at S (electron pairing in 3p causes repulsion, lowering the energy needed).

Melting points across Period 3

• Na, Mg, Al: giant metallic structures, melting point rises Na < Mg < Al because of increasing charge on the ion and more delocalised electrons per atom, giving stronger metallic bonding.
• Si: giant covalent (macromolecular), very high melting point (1410C) - loads of strong covalent bonds to break.
• P4, S8, Cl2, Ar: simple molecular or monatomic, melting points are low and depend on the strength of London (van der Waals) forces. S8 has more electrons than P4 so a higher melting point than expected from the trend; Ar is lowest as a single atom with weak dispersion forces only.
• Common mistake: students say 'ionic bonds break' for Period 3 melting - only Na, Mg, Al are metallic, not ionic. Nothing in Period 3 itself is ionic bonded (that's between elements, e.g. NaCl).

Reactions of Period 3 elements with water and oxygen

• Na reacts vigorously with cold water: 2Na + 2H2O -> 2NaOH + H2. Mg reacts very slowly with cold water but readily with steam: Mg + H2O(g) -> MgO + H2.
• Al does not react with water (protective oxide layer, Al2O3, is inert and adherent).
• Oxides: Na2O and MgO are basic (react with acids). Al2O3 is amphoteric (reacts with both acids and bases). SiO2 is acidic but weakly so (only reacts with hot concentrated NaOH). P4O10, SO2/SO3 are acidic, forming acidic solutions in water (e.g. SO3 + H2O -> H2SO4).

Group 2: reactivity and solubility trends

• Reactivity with water increases down the group (Mg barely reacts with cold water; Ba reacts vigorously) because ionisation energy decreases as atomic radius increases and shielding increases, so electrons are lost more easily.
• Solubility of hydroxides increases down the group (Mg(OH)2 is only slightly soluble, giving a weakly alkaline solution used as milk of magnesia; Ba(OH)2 is much more soluble).
• Solubility of sulfates decreases down the group (MgSO4 is soluble; BaSO4 is essentially insoluble - this is the basis of the test for sulfate ions using barium chloride/nitrate, forming a white precipitate).

Group 7: halogens and halides

• Reactivity (oxidising power) decreases down the group as atomic radius increases and the incoming electron is more shielded, so attraction to the nucleus weakens. Displacement reactions confirm this: chlorine displaces bromide and iodide from solution; bromine displaces iodide but not chloride.
• Reducing power of halide ions increases down the group (I- is the strongest reductant, F- the weakest) - shown by reactions with concentrated sulfuric acid: NaCl gives only steamy fumes of HCl; NaBr also produces some SO2 (Br- reduces H2SO4); NaI produces H2S and a black solid (I2/S), since I- is oxidised all the way and reduces sulfur from +6 to -2.
• Test for halide ions: add dilute nitric acid then silver nitrate solution. AgCl = white precipitate (dissolves in dilute ammonia), AgBr = cream precipitate (dissolves in concentrated ammonia only), AgI = yellow precipitate (insoluble even in concentrated ammonia).

Naming and formulae

Organic chemistry runs on IUPAC nomenclature. Find the longest carbon chain through the functional group, number it to give the lowest locants, and name substituents alphabetically as prefixes.

• General formulae: alkanes CnH2n+2, alkenes CnH2n, and know how to derive homologous series formulae for alcohols, haloalkanes etc.
• Empirical formula is the simplest whole-number ratio; molecular formula is the actual number of atoms, found using molar mass.
• Structural isomers have the same molecular formula but different structures (chain, position, or functional group isomerism).

Types of formula

• General, empirical, molecular, structural, displayed and skeletal formulae are all examinable.
• Skeletal formulae show only the carbon skeleton as lines, with hydrogens on carbon omitted.

Reaction mechanisms

You must be able to draw curly arrow mechanisms accurately.

• Free-radical substitution (alkanes with halogens, UV light): initiation (homolytic fission), propagation (two steps), termination.
• Electrophilic addition (alkenes with Br2, HBr, H2SO4): curly arrow from the pi bond to the electrophile, forming a carbocation intermediate.
• Markovnikov's rule: with unsymmetrical alkenes and HX, the major product forms via the more stable (more substituted) carbocation.

Alkanes

• Alkanes are saturated hydrocarbons, relatively unreactive due to strong, non-polar C-C and C-H bonds.
• Combustion: complete combustion gives CO2 and H2O; incomplete gives CO or C (soot) - a key exam point on air quality.
• Cracking (thermal or catalytic) breaks long-chain alkanes into shorter alkanes and alkenes to match market demand for fuels.

Alkenes

• Alkenes are unsaturated, containing a C=C double bond, which is a region of high electron density - this is why they undergo electrophilic addition, not substitution.
• Test for unsaturation: shake with bromine water - orange to colourless decolourisation confirms a C=C bond.
• E/Z isomerism arises from restricted rotation around the C=C bond; needs two different groups on each carbon. Use CIP priority rules to assign E or Z.

Common mistakes

• Forgetting hydrogens on skeletal formulae, or miscounting the chain length when numbering.
• Drawing curly arrows from atoms instead of bonds or lone pairs.
• Confusing E/Z with cis/trans when priority groups differ from position groups.
• Not stating 'major product' when applying Markovnikov's rule.

Optical isomerism

Optical isomers (enantiomers) occur when a carbon has four different groups attached - a chiral centre.

• They are non-superimposable mirror images.
• Enantiomers rotate plane-polarised light in opposite directions.
• A racemic mixture (racemate) has equal amounts of both enantiomers and shows no net optical rotation.
• Common mistake: forgetting that nucleophilic addition to a carbonyl (eg HCN to an aldehyde/ketone) produces a racemic mixture because the nucleophile can attack from either face.

Aldehydes and ketones

Both contain the carbonyl group C=O, made by oxidising alcohols.

• Aldehydes are oxidised by Fehling's solution (blue Cu2+ to brick-red Cu2O precipitate) and Tollens' reagent (silver mirror forms). Ketones give no reaction with either.
• Both react with HCN (via a nucleophilic addition mechanism, using KCN plus dilute acid) to form hydroxynitriles.
• 2,4-DNP (Brady's reagent) forms an orange precipitate with any carbonyl, confirming C=O but not distinguishing aldehyde from ketone. The precipitate's melting point identifies the specific carbonyl compound.

Carboxylic acids and esters

Carboxylic acids react with carbonates to release CO2 (fizzing) - a good test.

• Esterification: acid + alcohol, warmed with concentrated H2SO4 catalyst, forms an ester plus water (equilibrium reaction).
• Esters are hydrolysed by dilute acid (reversible) or by NaOH (irreversible, forms carboxylate salt plus alcohol) - alkaline hydrolysis is used to make soap from fats (saponification).

Aromatic chemistry (benzene)

Benzene's actual structure is a delocalised ring of six electrons above and below the plane, not alternating double bonds (Kekule model is wrong).

• Evidence: benzene's real enthalpy of hydrogenation is less exothermic than predicted for three C=C bonds, and all C-C bond lengths are equal (between single and double).
• Benzene undergoes electrophilic substitution, not addition, because the delocalised ring is very stable.
• Nitration: benzene plus concentrated HNO3/H2SO4 at 50 degrees C forms nitrobenzene; the electrophile is NO2+.
• Friedel-Crafts acylation: benzene plus acyl chloride with AlCl3 catalyst, under anhydrous reflux, forms a phenyl ketone.

Amines, amides, amino acids and polymers

Amines are bases (lone pair on N); primary amines are made by reducing nitriles or nitrobenzene.

• Amino acids are zwitterions at their isoelectric point, are amphoteric, and form peptide bonds by condensation.
• Condensation polymers (polyesters, polyamides) link monomers via ester or amide bonds, losing a small molecule (water or HCl) each time.

Organic analysis

• IR spectroscopy identifies bonds via absorption: O-H (broad, 2500-3300 for acids, 3200-3550 for alcohols), C=O (1680-1750), C-H (2850-3100).
• Mass spectrometry: the molecular ion peak (M+) gives Mr; fragmentation patterns and loss of specific masses (eg 15 for CH3, 45 for COOH) help identify structure.
• NMR: chemical shift, splitting patterns (n+1 rule) and integration ratios build up the full structure - always check the AQA data sheet for shift values in the exam.