Atoms have protons (mass 1, charge +1), neutrons (mass 1, charge 0) and electrons (mass 1/1836, charge -1). Atomic number Z = number of protons. Mass number A = protons + neutrons. Isotopes have the same Z but different A, so different numbers of neutrons.
Enthalpy change (ΔH) is heat energy transferred at constant pressure. Exothermic = ΔH negative, energy released to surroundings, temperature rises. Endothermic = ΔH positive, energy absorbed, temperature falls.
Know the standard definitions exactly: standard enthalpy of formation (ΔHf, one mole of compound from elements in standard states), standard enthalpy of combustion (ΔHc, one mole of substance burned completely in oxygen), and standard enthalpy of reaction (ΔHr). All measured at 298 K and 100 kPa.
Bond enthalpies are always positive (bond breaking absorbs energy, bond making releases it). Use mean bond enthalpies to estimate ΔH: ΔH = sum(bonds broken) minus sum(bonds formed). This only gives an estimate because mean values are averaged across many compounds, not exact for a specific molecule.
Hess's Law: the enthalpy change for a reaction is independent of the route taken, because enthalpy is a state function. Use this to build energy cycles from formation or combustion data.
Rate of reaction depends on collision frequency and the proportion of collisions with energy above the activation energy (Ea), in the correct orientation. Increasing temperature increases both collision frequency and, more importantly, the fraction of molecules with E greater than Ea (shown by a Maxwell-Boltzmann distribution shifting right and flattening).
Catalysts provide an alternative reaction pathway with lower Ea, increasing the proportion of successful collisions without being used up. They do not change ΔH or the position of equilibrium, only the rate of reaching it.
Rate equations are found experimentally, never from stoichiometry. Rate = k[A]^m[B]^n. Order with respect to a reactant is the power it's raised to; overall order is the sum. Units of k depend on overall order (mol dm-3 s-1 for zero order, s-1 for first order, mol-1 dm3 s-1 for second order).
Le Chatelier's Principle: a system at equilibrium shifts to oppose an imposed change. Increasing pressure shifts equilibrium towards the side with fewer gas moles; increasing temperature shifts towards the endothermic direction; a catalyst has no effect on position of equilibrium, only rate.
Kc and Kp are equilibrium constants: Kc uses concentrations (mol dm-3), Kp uses partial pressures (usually atm or Pa), only valid for gas-phase equilibria. Both only change with temperature, never with concentration, pressure or catalyst changes.
Redox reactions involve simultaneous oxidation and reduction. Oxidation is loss of electrons, reduction is gain of electrons - remember OIL RIG.
Oxidation numbers help track electron transfer. Free elements are 0, simple ions equal their charge, oxygen is usually -2 (except peroxides, -1, and OF2, +2), hydrogen is usually +1 (except metal hydrides, -1). The sum of oxidation numbers in a neutral compound is 0, and in an ion it equals the ion's charge.
You must be able to construct half-equations and combine them, balancing electrons so none are left over. For reactions in acidic solution, balance oxygen with water and hydrogen with H+ ions.
A standard electrode potential (E-standard, in volts) measures the tendency of a half-cell to be reduced, measured against the standard hydrogen electrode (SHE), which is defined as 0.00V. Standard conditions are 298K, 100 kPa, and 1 mol dm-3 solution concentration.
In the electrochemical series, the more positive the E-standard value, the greater the tendency to be reduced (gain electrons) - these are the stronger oxidising agents. The more negative, the greater the tendency to be oxidised - stronger reducing agents.
To work out cell EMF (Ecell): Ecell = E(positive electrode) minus E(negative electrode), using the more positive value as the cathode (reduction) and more negative as the anode (oxidation). A positive Ecell means the reaction is feasible (thermodynamically favourable), though feasibility does not guarantee the reaction happens quickly - kinetics can still make it very slow.
Learn the convention: anode (oxidation, negative electrode) on the left, cathode (reduction, positive electrode) on the right, with single vertical lines for phase boundaries and a double vertical line for the salt bridge. Write the more oxidised species furthest from the salt bridge on each side.
Hydrogen-oxygen fuel cells produce electricity from the reaction of hydrogen and oxygen, with water as the only product, making them cleaner than combustion engines, though hydrogen storage and production remain practical challenges.
Across Period 3 (Na to Ar), atomic radius decreases because nuclear charge increases while shielding stays roughly constant, pulling electrons in tighter. First ionisation energy generally rises across the period, with small drops at Al (starts filling 3p, lower energy than 3s) and S (first paired electron in 3p, extra repulsion). Melting points rise from Na to Al (metallic bonding gets stronger as more delocalised electrons and smaller ionic radius increase attraction), peak hugely at Si (giant covalent, strong directional bonds), then drop sharply from P to Ar (simple molecular structures, only weak van der Waals forces between molecules, so little energy needed to melt them).
Na2O and MgO are basic oxides, they form alkaline solutions with water (NaOH pH 13-14, Mg(OH)2 is only slightly soluble so pH is lower, around 9-10). Al2O3 is amphoteric, it does not react with water but reacts with both acids and bases. SiO2 does not react with water (giant covalent, very stable). P4O10, SO2 and SO3 are acidic oxides, they react with water to give acidic solutions (phosphoric, sulfurous and sulfuric acid), pH roughly 0-2 for the sulfur and phosphorus oxides.
Going down Group 2 (Be to Ba), reactivity with water increases, Mg reacts very slowly with cold water but readily with steam, while Ca, Sr and Ba react increasingly vigorously with cold water. This is because ionisation energy decreases down the group (atomic radius increases, more shielding, easier to lose the outer electrons), even though this is only a partial explanation examiners want you to state clearly. Solubility trends matter for exams: Group 2 hydroxides become MORE soluble down the group (Mg(OH)2 nearly insoluble, Ba(OH)2 fairly soluble), but Group 2 sulfates become LESS soluble down the group (MgSO4 soluble, BaSO4 essentially insoluble, this is the basis of the test for sulfate ions using BaCl2).
Reactivity as oxidising agents decreases down Group 7 (F2 most reactive, I2 least) because atomic radius increases and shielding increases, so incoming electrons are attracted less strongly. Displacement reactions: a more reactive (higher) halogen will displace a less reactive halide ion from solution, e.g. chlorine water added to potassium bromide solution turns the solution orange as Br2 forms.
Organic chemistry starts with naming. Use IUPAC rules: find the longest carbon chain containing the principal functional group, number it to give the lowest locants to that group, and name substituents alphabetically with locants. Know the difference between empirical, molecular, general, structural, displayed and skeletal formulae, and be able to convert between them.
Structural isomers have the same molecular formula but different structures: chain, position and functional group isomers. Stereoisomers have the same structural formula but different spatial arrangement. E/Z isomerism arises from restricted rotation around a C=C double bond and needs the CIP priority rules (higher atomic number = higher priority on each carbon) when the groups are not simply cis/trans.
Alkanes are saturated, general formula CnH2n+2, fairly unreactive due to strong, non-polar C-H and C-C bonds. Free-radical substitution with halogens (e.g. CH4 + Cl2 in UV light) proceeds via initiation (homolytic fission), propagation and termination steps. A common mistake is forgetting termination gives multiple possible products, not just the 'expected' one.
Alkenes are unsaturated, general formula CnH2n, and the C=C double bond (one sigma, one pi bond) is the reactive site, undergoing electrophilic addition. Learn the mechanism for addition of hydrogen halides and halogens, including curly arrows starting from the pi bond. Markovnikov's rule: with unsymmetrical alkenes and HX, the major product has H added to the carbon with more hydrogens already attached, because the more stable (more substituted) carbocation intermediate forms preferentially.
Carbon-halogen bond polarity and bond enthalpy control reactivity: C-I is weakest and most reactive, C-F is strongest and least reactive, despite C-F being most polar. This trips students up constantly - reactivity in nucleophilic substitution depends on bond enthalpy, not polarity. Nucleophilic substitution with OH-, CN- and NH3 produces alcohols, nitriles and amines respectively. Elimination with hot ethanolic OH- produces alkenes.
Classified as primary, secondary or tertiary by the carbon bonded to -OH. Oxidation with acidified potassium dichromate(VI) differs by class: primary alcohols oxidise to aldehydes then carboxylic acids (distil off the aldehyde to stop at that stage; reflux for full oxidation), secondary alcohols oxidise only to ketones, and tertiary alcohols resist oxidation entirely. Common mistake: forgetting tertiary alcohols simply do not react with the oxidising agent.
Always use curly arrows correctly - they show movement of a pair of electrons, starting from a bond or lone pair and ending at an atom or new bond. Practise drawing full mechanisms for free-radical substitution, electrophilic addition and nucleophilic substitution, as mechanism questions are heavily marked on arrow accuracy.
A chiral carbon has four different groups attached, giving two non-superimposable mirror-image forms called enantiomers.
Enantiomers rotate plane-polarised light in opposite directions and have identical physical properties otherwise.
A racemic mixture (racemate) has equal amounts of both enantiomers and shows no net optical rotation.
Optical isomerism matters in drug synthesis, since only one enantiomer of a drug may be biologically active.
Aldehydes are oxidised by both Fehling's solution (blue Cu2+ to brick-red Cu2O precipitate) and Tollens' reagent (silver mirror from Ag+ to Ag).
Ketones give a negative result with both tests, since they cannot be oxidised easily.
2,4-DNP forms an orange precipitate with any carbonyl (aldehyde or ketone), confirming C=O is present, and its melting point identifies the specific compound.
NaBH4 in water or ethanol reduces aldehydes to primary alcohols and ketones to secondary alcohols.
Carboxylic acids react with carbonates to release CO2 (fizzing), a quick positive test distinguishing them from alcohols and phenols.
Esters are made by refluxing a carboxylic acid with an alcohol using concentrated H2SO4 catalyst (Fischer esterification), an equilibrium reaction.
Esters hydrolyse under acid (reflux with dilute acid, reversible) or base (reflux with NaOH, irreversible, forming a carboxylate salt plus alcohol).
Benzene undergoes electrophilic substitution, not addition, because the delocalised ring of six pi electrons is unusually stable (extra stabilisation ~150 kJ/mol vs a theoretical Kekule structure).
Nitration uses conc. HNO3 with conc. H2SO4 catalyst at 50 degrees C, generating the NO2+ electrophile.
Friedel-Crafts acylation uses an acyl chloride with AlCl3 catalyst under reflux, anhydrous conditions, forming a ketone attached to the ring.
Mass spectrometry: the molecular ion peak (M+) gives Mr; fragmentation peaks help identify substructures.
Infrared spectroscopy: O-H (broad, 2500-3300 for acids, 3200-3550 for alcohols), C=O (1680-1750, sharp), N-H (3300-3500).
NMR: 13C gives number of distinct carbon environments; 1H integration gives ratio of protons; splitting patterns follow the n+1 rule; TMS is the reference standard at 0 ppm.
D2O shake removes O-H and N-H peaks from an 1H NMR spectrum, confirming their position by disappearance.
Don't confuse Fehling's (aldehyde only) with 2,4-DNP (any carbonyl) — 2,4-DNP does NOT distinguish aldehydes from ketones.
Always state conditions for esterification and Friedel-Crafts reactions; examiners award marks for reflux, catalyst, and anhydrous conditions.
Remember base hydrolysis of esters is one-way (not an equilibrium) because the carboxylate ion formed won't re-form the ester.