Atomic structure basics

• Atoms have protons (+1, mass 1) and neutrons (0, mass 1) in the nucleus, with electrons (-1, negligible mass) in shells around it.
• Atomic number (Z) = number of protons = number of electrons in a neutral atom.
• Mass number (A) = protons + neutrons.
• Isotopes have the same Z but different A (different neutron numbers) — same chemical properties, slightly different physical ones.
• Relative atomic mass (Ar) is a weighted mean based on isotopic abundance, measured relative to carbon-12.

Electron configuration

• Electrons fill orbitals in order of increasing energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p.
• Sub-shells hold a fixed maximum: s = 2 electrons, p = 6, d = 10, f = 14.
• Orbitals within a sub-shell are filled singly first (Hund's rule) before pairing up.
• Watch for the exam favourite: 4s fills before 3d, but empties first when forming transition metal ions (so Fe is [Ar]3d6 4s2 but Fe2+ is [Ar]3d6, not 3d4 4s2).
• Chromium and copper are exceptions: Cr is [Ar]3d5 4s1 and Cu is [Ar]3d10 4s1, because a half-full or full d sub-shell is extra stable.

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, since removing an electron from an increasingly positive ion needs more energy.
• Big jumps between successive ionisation energies show a change of shell — this is the classic evidence question for shell structure.
• Across a period, ionisation energy generally rises (more protons, same shell), but drops slightly at group 3 (start of p sub-shell, easier to remove) and group 6 (electron pairing repulsion).
• Down a group, ionisation energy falls because atomic radius increases and shielding increases, outweighing the extra protons.

Bonding

• Ionic bonding: electrostatic attraction between oppositely charged ions in a giant lattice; strong in all directions, high melting points, conducts only when molten or dissolved.
• Covalent bonding: shared pair of electrons between two atoms; can be simple molecular (low melting point, doesn't conduct) or giant covalent like diamond and graphite (very high melting point).
• Metallic bonding: delocalised electrons attract a lattice of positive metal ions; conducts as a solid and liquid, malleable because layers of ions can slide.
• Dative (coordinate) bonding: both shared electrons come from the same atom, shown with an arrow, e.g. in NH4+ or Al2Cl6.

Common mistakes

• Don't say 'electrons are lost to space' — always state the products form gaseous ions.
• Always give ionisation energy definitions in full with 'one mole' and 'gaseous' — examiners mark this strictly.
• Remember 4s empties before 3d when forming ions, even though it fills first.
• Distinguish shielding (inner electrons repelling outer ones) from nuclear charge (attracting them) — both affect radius and ionisation energy.

Enthalpy changes

Enthalpy change (ΔH) is heat energy transferred at constant pressure. Exothermic reactions release energy, ΔH is negative. Endothermic reactions absorb energy, ΔH is positive.

• Standard enthalpy of formation ΔHf°: forming 1 mole of a compound from its elements in standard states (100 kPa, 298 K).
• Standard enthalpy of combustion ΔHc°: burning 1 mole of a substance completely in oxygen under standard conditions.
• Bond enthalpies are always positive (breaking bonds absorbs energy); making bonds releases energy.
• Hess's Law: total enthalpy change is independent of the route taken. Use it to find ΔH indirectly via formation or combustion data.
• q = mcΔT is used in calorimetry, where c for water is 4.18 J g⁻¹ K⁻¹. Common mistake: forgetting to convert mass of solution properly or using the wrong sign.
• Mean bond enthalpy calculations: ΔH = (bonds broken) − (bonds made). These give estimates, not exact values, because mean bond enthalpies are averaged across many compounds.

Kinetics

Rate of reaction depends on collision theory: particles must collide with energy ≥ activation energy (Ea) and correct orientation.

• Increasing temperature increases rate because more particles have energy ≥ Ea (Maxwell-Boltzmann distribution shifts right and flattens) — this has a much bigger effect than the small increase in collision frequency.
• Catalysts provide an alternative route with lower Ea, increasing the proportion of successful collisions, without being used up.
• Rate equations: rate = k[A]^m[B]^n. Orders must be found experimentally, never from stoichiometry.
• Overall order = sum of individual orders. Units of k depend on overall order.
• Rate-determining step is the slowest step; only reactants appearing in the rate equation (up to that step) affect rate.
• The Arrhenius equation, k = Ae^(−Ea/RT), links rate constant to temperature and activation energy; a linear form (ln k vs 1/T) is used to find Ea from gradient = −Ea/R.

Equilibria

Dynamic equilibrium: forward and reverse rates are equal in a closed system, concentrations stay constant (not equal).

• Le Chatelier's Principle predicts the direction of shift when conditions change, but does NOT explain the rate or the equilibrium constant Kc/Kp itself.
• Kc is only affected by temperature; changing concentration, pressure or adding a catalyst does not change Kc, only the position of equilibrium (except catalysts, which change neither).
• Kc = [products]/[reactants], each raised to its stoichiometric power, only for homogeneous equilibria (all species same physical state); pure solids and liquids are omitted from Kc expressions.
• For an exothermic forward reaction, increasing temperature decreases Kc; for endothermic, increasing temperature increases Kc.
• Common mistake: confusing Kc with the rate constant k, or forgetting units on Kc, which depend on the equilibrium expression.
• pH and acid-base equilibria (Ka, Kw, pKa) link to this topic — remember Kw = 1.00 × 10⁻¹⁴ mol² dm⁻⁶ at 298 K.

Oxidation states

Oxidation state rules: uncombined elements = 0, simple ions = charge on the ion, oxygen is -2 (except -1 in peroxides, -1/2 in superoxides, and +2 in OF2), hydrogen is +1 (except -1 in metal hydrides like NaH), fluorine is always -1. All oxidation states in a neutral compound sum to zero; in an ion they sum to the ion's charge.

Oxidation and reduction

OILRIG: Oxidation Is Loss, Reduction Is Gain (of electrons). An oxidising agent gets reduced (gains electrons, its oxidation state falls); a reducing agent gets oxidised (loses electrons, its oxidation state rises). In a redox equation, the oxidation state changes must balance: electrons lost by one species equal electrons gained by the other.

Half-equations

Write half-equations separately for oxidation and reduction, balance atoms first, then charge using electrons, then combine so electrons cancel exactly. For reactions in acidic solution, balance oxygens with H2O and hydrogens with H+.

Electrode potentials

Standard electrode potential E-standard is measured against the standard hydrogen electrode (SHE, defined as 0.00V) at 298K, 100kPa, and 1.00 mol dm-3 solution concentration. A standard cell is two half-cells connected by a wire (electrons) and a salt bridge (ions, often KNO3-soaked filter paper) to complete the circuit without mixing solutions.

Cell EMF (E-cell) = E(positive electrode) - E(negative electrode), using the more positive minus the less positive (or the reduction potentials of the right-hand minus left-hand electrode as written). A more positive E-standard means a stronger oxidising agent (more likely to be reduced); a more negative E-standard means a stronger reducing agent.

Predicting feasibility

A reaction is feasible if E-cell is positive. Use E-cell to predict direction, but remember predictions from electrode potentials ignore kinetics: a reaction can be thermodynamically feasible yet too slow to observe (high activation energy), so real experiments do not always match the prediction.

Common mistakes

Mixing up which half-reaction is reversed when combining half-equations. Forgetting to scale electrons so they cancel before adding half-equations together. Getting E-cell the wrong way round (always more positive minus less positive, or right minus left as conventionally drawn). Confusing oxidation state change with actual charge on an ion. Forgetting that a negative E-cell means the reaction as written is not feasible, but the reverse reaction is.

Fuel cells and storage cells

Hydrogen fuel cells oxidise H2 at the negative electrode and reduce O2 at the positive electrode, producing only water; they are more efficient than burning fuel since they avoid heat-engine losses. Rechargeable cells (like Li-ion) can be recharged by reversing the cell reaction using an external supply; non-rechargeable cells cannot.

Periodicity: structure and trends

Period 3 (Na to Ar) shows how atomic structure controls physical properties.

• Atomic radius decreases left to right across a period. Increasing nuclear charge pulls the same-shell electrons in tighter, and shielding stays roughly constant.
• First ionisation energy generally increases across a period, but with two dips: Mg to Al (Al's outer electron is in a 3p orbital, higher energy and slightly shielded by the 3s pair, so easier to remove) and P to S (S has a paired 3p electron; the repulsion between the paired electrons makes it easier to lose than expected).
• Melting points rise from Na to Al (giant metallic structures, more delocalised electrons per atom increase metallic bonding strength), peak at Si (giant covalent, strong covalent bonds throughout the lattice), then crash for P4, S8, Cl2 and Ar (simple molecular or monatomic, held only by weak van der Waals forces, so melting point depends on molecule size: S8 > P4 > Cl2 > Ar).

Reactions of Period 3 elements and oxides with water

• Na and Mg react with water; Na reacts vigorously (2Na + 2H2O to 2NaOH + H2), Mg reacts very slowly with cold water but readily with steam to give MgO and H2.
• Oxides across the period show a clear acid-base trend: Na2O and MgO are basic (form alkaline solutions or react with acids); Al2O3 is amphoteric (reacts with both acids and alkalis); SiO2 is acidic but insoluble; P4O10, SO2/SO3 are acidic, dissolving in water to give acidic solutions (e.g. SO3 + H2O to H2SO4).

Group 2: the alkaline earth metals

• Reactivity increases down the group (Mg to Ba) because ionisation energy falls as atomic radius increases and shielding increases, so the outer electrons are lost more easily.
• Reaction with water gets more vigorous down the group: Mg barely reacts with cold water, Ca reacts steadily, Ba reacts vigorously.
• Solubility trends matter for exam calculations: hydroxides become MORE soluble down the group (Mg(OH)2 is only slightly soluble, giving milk of magnesia; Ba(OH)2 is quite soluble), while sulfates become LESS soluble down the group (MgSO4 is soluble, BaSO4 is essentially insoluble, used as the classic test for sulfate ions with acidified BaCl2).

Group 7: the halogens

• Reactivity (oxidising power) decreases down the group as atomic radius increases and shielding increases, making it harder to attract an extra electron.
• Displacement reactions demonstrate this: chlorine displaces bromine and iodine from their halide solutions; bromine displaces iodine only.
• Test for halide ions: add acidified silver nitrate. Cl- gives a white precipitate (soluble in dilute NH3), Br- gives a cream precipitate (soluble in concentrated NH3 only), I- gives a yellow precipitate (insoluble in concentrated NH3).

Common mistakes

• Writing 'ionisation energy decreases across the period' — it generally increases, only two specific dips exist.
• Forgetting that halogen reactivity is about gaining electrons, so it decreases down the group, the opposite trend to Group 2 metal reactivity.
• Mixing up sulfate (decreasing solubility down Group 2) with hydroxide (increasing solubility down Group 2).

Nomenclature and formulae

Organic chemistry runs on IUPAC naming. Learn the stem for chain length (meth-1, eth-2, prop-3, but-4, pent-5, hex-6) then the suffix for the functional group: -ane (alkane), -ene (alkene), -ol (alcohol), -al (aldehyde), -one (ketone), -oic acid (carboxylic acid), -oate (ester).

• General formulae matter: alkanes CnH2n+2, alkenes CnH2n.
• Empirical formula is the simplest whole-number ratio; molecular formula is the actual number of atoms; you get molecular from empirical using the Mr.
• Displayed, structural, skeletal and general formulae are all examinable — practise converting between them.

Isomerism

Structural isomers share a molecular formula but differ in structure: chain, position and functional group isomers.

Stereoisomers share the same structural formula but atoms are arranged differently in space.

• E/Z isomerism arises from restricted rotation around a C=C double bond and needs two different groups on each carbon of the double bond.
• Use CIP (Cahn-Ingold-Prelog) priority rules based on atomic number to assign E (opposite sides, higher priorities) or Z (same side).
• Common mistake: forgetting that cis-trans naming only works when each carbon carries an H; E/Z is the general system examiners want.

Reaction mechanisms

Alkanes undergo free-radical substitution with halogens in UV light: initiation (homolytic fission), propagation (two steps), termination. Learn to draw curly arrows showing single-barbed (radical) movement.

Alkenes undergo electrophilic addition, for example with bromine water (decolourises, test for unsaturation) or hydrogen halides. With unsymmetrical alkenes and HX, Markovnikov's rule applies: the major product has the halogen on the carbon that gives the more stable (more substituted) carbocation intermediate.

Practical skills and hazards

• Alcohols are classified primary, secondary or tertiary by the carbon bonded to -OH; this determines oxidation behaviour (primary to aldehyde then acid, secondary to ketone only, tertiary resists oxidation) using acidified potassium dichromate(VI), which turns from orange to green.
• Combustion and cracking (thermal or catalytic) of alkanes are key industrial reactions - cracking produces shorter alkanes and alkenes to meet fuel demand.
• Common mistakes: mixing up homolytic and heterolytic fission, drawing curly arrows from atoms instead of bonds or lone pairs, and confusing empirical with molecular formula in calculations.

Aldehydes and ketones

Both contain the carbonyl group C=O. Aldehydes have it at the end of the chain (CHO) and are made by oxidising primary alcohols; ketones have it mid-chain (C=O flanked by two carbons) and come from secondary alcohols.

• Test to tell them apart: Tollens' reagent (ammoniacal silver nitrate) gives a silver mirror with aldehydes only, as the aldehyde is oxidised to a carboxylic acid.
• Fehling's solution turns from blue Cu2+ to a brick-red Cu2O precipitate with aldehydes; ketones give no change.
• Both react with 2,4-dinitrophenylhydrazine (2,4-DNP) to give an orange precipitate, confirming a carbonyl is present (but not which type).
• Reduction: NaBH4 reduces aldehydes to primary alcohols and ketones to secondary alcohols.
• Nucleophilic addition with HCN (via KCN + dilute acid) forms hydroxynitriles; the mechanism involves the CN- nucleophile attacking the delta-positive carbonyl carbon.

Carboxylic acids and esters

Carboxylic acids (COOH) are weak acids, reacting with carbonates to fizz CO2, with metals to give H2, and with alcohols (conc. H2SO4 catalyst, reflux) to form esters plus water in condensation reactions. Esters are hydrolysed by acid (reversible, gives acid + alcohol) or base (irreversible, gives carboxylate salt + alcohol - this is saponification).

Aromatic chemistry (benzene)

Benzene, C6H6, has a delocalised ring of six pi electrons above and below the plane, making all six C-C bonds equal length (139 pm, between single 154 pm and double 134 pm). This delocalisation makes benzene more stable than the Kekule model predicts by about 150 kJ/mol (from enthalpy of hydrogenation data), and explains why benzene undergoes electrophilic substitution, not addition - substitution keeps the stable ring intact.

• Nitration: benzene + conc. HNO3/conc. H2SO4 catalyst, 50 degC, gives nitrobenzene; the electrophile is NO2+.
• Friedel-Crafts acylation: benzene + acyl chloride, AlCl3 catalyst, anhydrous reflux, gives a phenyl ketone; used to add a carbon side chain since benzene cannot do simple addition reactions.

Amines, amides, amino acids and polymers

Amines are bases (lone pair on N). Amino acids are amphoteric, existing as zwitterions at neutral pH. Condensation polymers: polyesters (diol + dicarboxylic acid, ester links) and polyamides (diamine + diacid or diacyl chloride, amide links, e.g. nylon).

Organic analysis

• Mass spectrometry: molecular ion peak (M+) gives Mr; fragmentation pattern helps identify structure.
• IR spectroscopy: identifies functional groups from absorption at characteristic wavenumbers, e.g. C=O around 1680-1750 cm-1, O-H (acid) broad 2500-3300 cm-1, O-H (alcohol) broad 3200-3550 cm-1, N-H 3300-3500 cm-1.
• NMR: chemical shift (ppm, relative to TMS at 0), integration trace gives relative proton ratios, splitting patterns follow the n+1 rule (a peak is split into n+1 lines by n adjacent, non-equivalent protons).

Common mistakes

• Confusing Tollens' (silver mirror) with Fehling's (brick-red precipitate) - know both.
• Forgetting 2,4-DNP only shows a carbonyl is present, not aldehyde vs ketone.
• Writing benzene with alternating double bonds instead of a delocalised ring in mechanisms.
• Forgetting acid hydrolysis of esters is reversible but base hydrolysis is not.