A-level and IB · Sciences

A-level Chemistry: core concepts

A-level Chemistry flashcards on moles, bonding, energetics, kinetics, equilibria, pH, redox, inorganic trends, organic mechanisms and spectra, with examples.

40 cards6 topicsAbout 15 minutes a roundUK A-level Chemistry, comparable to Grade 12 chemistry plus much of first-year university general chemistry

Independent study material written by StudyBrix, not produced or reviewed by any exam board or by the International Baccalaureate Organization. Check your own specification or subject guide.

Print cards to cut out (PDF)Print a study list (PDF)

About this deck

This deck covers the physical, inorganic and organic chemistry that appears in most A-level Chemistry courses in England: amount of substance, atomic structure, bonding and shapes, energetics and entropy, rates and equilibria, acids, bases and buffers, redox and electrode potentials, periodic trends and transition metals, organic reaction mechanisms and analytical techniques. It also works as a bridge for students starting first-year general chemistry in Canada.

AQA, Edexcel and OCR teach the same national core but arrange it differently, use different data sheets and sometimes slightly different bond enthalpy or electrode potential values. Check your own specification and data booklet, and use the values your exam gives.

You should be comfortable with GCSE chemistry, balancing equations and rearranging formulae. Work every calculation card on paper with units and significant figures before flipping it. For mechanism cards, draw curly arrows from a lone pair or bond to an atom, and say out loud why each step happens.

Common mistakes: using cm^3 instead of dm^3 (or m^3 in pV = nRT); forgetting that enthalpy of formation of an element is zero; writing an equilibrium constant with solids included; and drawing curly arrows that start from an atom rather than an electron pair. Also remember that a catalyst changes the rate, not the position of equilibrium.

Before you start: GCSE chemistry (or Grade 11 chemistry): atomic structure, balancing equations, moles at an introductory level.

All 40 cards

Cover the definitions with a sheet of paper and test yourself, or use the study modes above.

Atoms and amounts

Relative atomic mass: The weighted mean mass of an atom of an element compared with one twelfth of the mass of a carbon-12 atom.Example: Chlorine: 75% of mass 35 and 25% of mass 37 gives 0.75 * 35 + 0.25 * 37 = 35.5.
Amount of substance from mass: Number of moles equals mass in grams divided by molar mass: n = m/M.Example: 4.4 g of CO2 (M = 44 g/mol) is 0.10 mol.
Ideal gas equation: pV = nRT with p in Pa, V in m^3, T in K and R = 8.31 J/(K mol).Example: 0.50 mol at 300 K and 100 kPa: V = 0.50 * 8.31 * 300/100000 = 0.0125 m^3 = 12.5 dm^3.
Titration moles: Moles in solution = concentration (mol/dm^3) times volume (dm^3). Convert cm^3 by dividing by 1000.Example: 25.0 cm^3 of 0.100 mol/dm^3 NaOH contains 0.0250 * 0.100 = 0.00250 mol.
Empirical formula: The simplest whole-number ratio of atoms of each element in a compound, found from masses or percentages.Example: C 40.0%, H 6.7%, O 53.3%: divide by 12, 1 and 16 to get 3.33 : 6.7 : 3.33, so CH2O.
Atom economy: Mass of the desired product divided by the total mass of all products, times 100; it shows how much of the reactants ends up as useful product.Example: CaCO3 -> CaO + CO2 when making CaO: 56/100 = 56%.
Electron configuration exceptions: Chromium and copper fill 3d before completing 4s because a half-full or full d subshell is more stable.Example: Cr is [Ar] 3d5 4s1 and Cu is [Ar] 3d10 4s1.
First ionization energy trend: Energy to remove one electron from each atom in a mole of gaseous atoms. It generally rises across a period and falls down a group.Example: Small dips occur from Mg to Al (3p electron is higher in energy) and from P to S (paired electron repulsion).

Bonding and energetics

Electron pair repulsion shapes: Electron pairs around a central atom spread out to minimize repulsion; lone pairs repel more than bonding pairs and squeeze bond angles.Example: CH4 is tetrahedral at 109.5 degrees, NH3 pyramidal at about 107 and H2O bent at about 104.5.
Hydrogen bonding: A strong intermolecular attraction between a lone pair on N, O or F and a hydrogen atom bonded to N, O or F in another molecule.Example: Water boils at 100 degrees C, much higher than H2S, because of these bonds.
Calorimetry: Heat transferred q = mc delta T; dividing by moles reacted and adding a sign gives the enthalpy change.Example: 50 g water rises 6.0 K: q = 50 * 4.18 * 6.0 = 1254 J. For 0.025 mol reacting, delta H = -50.2 kJ/mol.
Hess's law: The total enthalpy change for a reaction is the same whatever route is taken, as long as start and end are the same.Example: Enthalpy of reaction = sum of enthalpies of formation of products minus the sum for reactants.
Mean bond enthalpy calculation: delta H is about the bonds broken minus the bonds made, using average values, so the answer is approximate.Example: H2 + Cl2 -> 2HCl: 436 + 243 - 2 * 432 = -185 kJ/mol.
Gibbs free energy: delta G = delta H - T delta S. A reaction is feasible when delta G is zero or negative. Convert delta S to kJ before combining.Example: CaCO3 decomposition: delta H = +178 kJ/mol, delta S = +161 J/(K mol), so it becomes feasible above 178000/161 = 1106 K.
Lattice enthalpy: The enthalpy change when one mole of an ionic solid forms from its gaseous ions (or the reverse, depending on the definition your board uses). Found with a Born-Haber cycle.Example: MgO has a far more exothermic lattice formation than NaCl because its ions are smaller and doubly charged.

Kinetics and equilibria

Maxwell-Boltzmann distribution: A curve showing how many molecules have each energy. Raising the temperature moves it right, so many more molecules exceed the activation energy.Example: A 10 K rise can roughly double the rate of many reactions.
Rate equation: rate = k[A]^m[B]^n, where the orders m and n are found by experiment, not from the balanced equation.Example: Doubling [A] with [B] fixed makes the rate four times faster, so the reaction is second order in A.
Arrhenius equation: k = A e^(-Ea/RT). Taking logs gives ln k = ln A - Ea/(RT), so a plot of ln k against 1/T has gradient -Ea/R.Example: A gradient of -6000 K gives Ea = 6000 * 8.31 = 49900 J/mol, about 50 kJ/mol.
Le Chatelier's principle: When a system at equilibrium is disturbed, the position shifts to oppose the change.Example: N2 + 3H2 <=> 2NH3 (exothermic): raising pressure shifts right; raising temperature shifts left.
Equilibrium constant Kc: The ratio of product concentrations to reactant concentrations at equilibrium, each raised to its coefficient. Only temperature changes its value.Example: For N2 + 3H2 <=> 2NH3, Kc = [NH3]^2/([N2][H2]^3) with units mol^-2 dm^6.
Partial pressure: The pressure one gas in a mixture would exert alone; it equals mole fraction times total pressure.Example: 0.25 mole fraction in a 200 kPa mixture gives a partial pressure of 50 kPa.

Acids, bases and redox

pH of a strong acid: pH = -log10[H+]; a strong monoprotic acid ionizes fully, so [H+] equals the acid concentration.Example: 0.020 mol/dm^3 HCl: pH = -log10(0.020) = 1.70.
Acid dissociation constant Ka: Ka = [H+][A-]/[HA] for a weak acid. Assuming [H+] = [A-] and little ionization, [H+] = sqrt(Ka * concentration).Example: 0.10 mol/dm^3 ethanoic acid, Ka = 1.74 * 10^-5: [H+] = 1.32 * 10^-3, so pH = 2.88.
Ionic product of water: Kw = [H+][OH-], which is 1.00 * 10^-14 mol^2 dm^-6 at 298 K. It lets you find the pH of strong bases.Example: 0.050 mol/dm^3 NaOH: [H+] = 10^-14/0.050 = 2.0 * 10^-13, so pH = 12.70.
Buffer solution: A mixture that resists pH change when small amounts of acid or base are added, usually a weak acid with its conjugate base.Example: When [HA] = [A-], pH = pKa; an ethanoic acid and sodium ethanoate mixture sits near pH 4.76.
Oxidation state: A number showing how many electrons an atom has lost or gained compared with the element, using set rules (O is usually -2, H usually +1).Example: In MnO4-, x + 4(-2) = -1, so manganese is +7.
Standard cell potential: E(cell) = E(positive electrode) - E(negative electrode), using standard electrode potentials. A positive value means the reaction is feasible.Example: Zn and Cu half-cells: 0.34 - (-0.76) = +1.10 V.

Inorganic chemistry

Silver nitrate test for halides: Acidified silver nitrate gives coloured silver halide precipitates that are told apart by colour and by solubility in ammonia.Example: Chloride gives white, bromide cream and iodide yellow; only silver chloride dissolves easily in dilute ammonia.
Group 2 trends: Down the group, atomic radius rises, first ionization energy falls and reactivity with water increases; hydroxides get more soluble and sulfates less soluble.Example: Barium sulfate is so insoluble it is used in barium meals for X-rays.
Complex ion: A central metal ion bonded to ligands by coordinate (dative) bonds; ligands donate a lone pair.Example: [Cu(H2O)6]2+ is octahedral; with excess concentrated HCl it becomes tetrahedral [CuCl4]2-, turning yellow-green.
Colour in transition metal complexes: Ligands split the d orbitals; electrons absorb visible light to jump between them and we see the complementary colour.Example: Changing the ligand, coordination number or oxidation state changes the energy gap and so the colour.

Organic chemistry and analysis

Free-radical substitution: A chain reaction of an alkane with a halogen in UV light: initiation, propagation and termination steps.Example: Cl2 -> 2Cl. (initiation); Cl. + CH4 -> HCl + CH3. and CH3. + Cl2 -> CH3Cl + Cl. (propagation).
Electrophilic addition: An electron-rich C=C double bond attacks an electrophile, forming a carbocation that a negative ion then joins.Example: Propene + HBr gives mainly 2-bromopropane through the more stable secondary carbocation.
Nucleophilic substitution: A nucleophile with a lone pair attacks the delta positive carbon of a haloalkane and the halide ion leaves.Example: Bromoethane + warm aqueous NaOH -> ethanol + bromide ion.
E/Z isomerism: Stereoisomerism from restricted rotation about C=C when each carbon carries two different groups; Z has the higher priority groups on the same side.Example: But-2-ene exists as E (methyl groups opposite) and Z (methyl groups on the same side).
Optical isomerism: Non-superimposable mirror-image molecules, caused by a chiral carbon bonded to four different groups; they rotate plane-polarized light in opposite directions.Example: Butan-2-ol has a chiral centre at carbon 2 (H, OH, CH3 and C2H5).
Electrophilic substitution of benzene: The delocalized ring attacks an electrophile, then loses H+ so the stable ring is restored.Example: Nitration: concentrated HNO3 and H2SO4 form NO2+, giving nitrobenzene.
Infrared spectroscopy: Bonds absorb infrared at characteristic wavenumbers, which identifies functional groups.Example: A very broad band from 2500 to 3300 cm^-1 plus a sharp C=O near 1700 cm^-1 suggests a carboxylic acid.
Proton NMR splitting: A peak is split into n + 1 lines by n hydrogens on neighbouring carbons.Example: Ethanol: CH3 is a triplet (2 neighbours) and CH2 is a quartet (3 neighbours); OH is usually a singlet.
Carbon-13 NMR: Each chemically different carbon environment gives one peak, so counting peaks shows molecular symmetry.Example: Propanone has 2 peaks; propanal has 3.