AP Intermediate Chemistry (BIEAP) flashcards

32 cards. Read the term and the definition here. The deck at the top of this page is optional.

Bohr energy formula

For a hydrogen-like atom, En = -13.6 Z^2/n^2 eV. Energy is negative, becomes less negative as n rises, and the spacing between levels decreases with increasing n.

de Broglie wavelength

A particle of mass m moving at speed v has wavelength λ = h/mv, where h is Planck's constant. This wave nature explains electron diffraction and the quantisation of angular momentum in Bohr's model.

Aufbau (n+l) rule

Orbitals fill in order of increasing (n+l); when two orbitals share the same (n+l), the one with lower n fills first. This gives 4s before 3d and 5s before 4d.

Hund's rule

Within a set of degenerate orbitals, electrons occupy separate orbitals with parallel spins before any pairing occurs. This minimises electron repulsion and maximises exchange energy, stabilising half-filled configurations.

Diagonal relationship

Certain elements of period 2 resemble the element diagonally below them in period 3 because of similar charge-to-radius ratios: Li-Mg, Be-Al, B-Si. They show similar compound formulas and often amphoteric behaviour.

Ionisation enthalpy trend

Ionisation enthalpy rises across a period as nuclear charge increases and falls down a group as size and shielding increase. It is endothermic, and successive values satisfy IE1 < IE2 < IE3.

Isoelectronic species radius

Species with the same number of electrons have smaller radii as nuclear charge increases, because the nucleus pulls the same electron cloud more strongly. Thus Al3+ < Mg2+ < Na+ < F- < O2-.

VSEPR theory

Electron pairs around a central atom arrange to minimise repulsion. Lone pair-lone pair repulsion is strongest, so lone pairs compress bond angles, as in NH3 (107°) and H2O (104.5°).

Bond order from MO theory

Bond order equals half the difference between bonding and antibonding electrons. Higher bond order gives shorter, stronger bonds, and unpaired electrons in antibonding orbitals explain the paramagnetism of O2.

Hydrogen bonding

A strong dipole-dipole attraction between H bonded to F, O or N and a lone pair on a nearby electronegative atom. It raises boiling points and explains water's high specific heat and surface tension.

Graham's law of effusion

The rate of effusion or diffusion of a gas is inversely proportional to the square root of its molar mass: r1/r2 = √(M2/M1). Lighter gases therefore escape faster through a small hole.

van der Waals equation

A real-gas equation written as (P + an^2/V^2)(V - nb) = nRT, where a corrects for intermolecular attraction and b for finite molecular volume. It predicts critical behaviour and deviations from ideality.

Critical temperature

The highest temperature at which a gas can be liquefied by pressure alone. Above Tc the substance exists only as a supercritical fluid and no meniscus forms between liquid and gas.

Hess's law

The enthalpy change of a reaction is independent of the route taken, so thermochemical equations can be added and reversed. It allows calculation of ΔH from standard enthalpies of formation or combustion.

Gibbs energy criterion

ΔG = ΔH - TΔS decides spontaneity at constant temperature and pressure: negative ΔG means spontaneous, zero means equilibrium, positive means non-spontaneous. ΔG° = -RT ln K links it to equilibrium.

Standard enthalpy of formation

The enthalpy change when one mole of a compound forms from its elements in their standard states at 298 K and 1 bar. It is defined as zero for elements in their standard states.

Le Chatelier's principle

When a system at equilibrium is disturbed by a change in concentration, pressure or temperature, it shifts in the direction that partially counteracts the change. Catalysts do not shift the position of equilibrium.

Solubility product Ksp

The equilibrium constant for dissolution of a sparingly soluble salt. For AB, Ksp = s^2; for AB2, Ksp = 4s^3. Precipitation occurs when the ionic product exceeds Ksp.

Buffer solution

A mixture of a weak acid and its conjugate base, or a weak base and its conjugate acid, that resists pH change on adding small amounts of acid or base. Its pH follows the Henderson-Hasselbalch equation.

Nernst equation

Cell potential depends on concentration: E = E° - (0.0591/n) log Q at 298 K. At equilibrium E becomes zero and E° = (0.0591/n) log K relates potential to the equilibrium constant.

Faraday's laws of electrolysis

The mass deposited at an electrode is proportional to charge passed, Q = It, and to molar mass divided by electrons transferred: m = MIt/(nF), with F = 96500 C mol^-1.

First-order half-life

For a first-order reaction, t½ = 0.693/k and is independent of initial concentration. A plot of ln[A] against time is a straight line with slope -k.

Arrhenius equation

Rate constant varies with temperature as k = A e^(-Ea/RT), where Ea is activation energy and A the pre-exponential factor. A catalyst lowers Ea and increases k without changing ΔH.

Freundlich adsorption isotherm

Empirical relation x/m = k p^(1/n) for adsorption of a gas on a solid, with 0 < 1/n < 1. A plot of log(x/m) against log p is linear with slope 1/n.

Hardy-Schulze rule

Coagulation of a lyophobic sol is caused by ions of opposite charge, and the coagulating power increases sharply with the charge of the ion, so Al3+ is far more effective than Na+.

Amphoteric oxide examples

Oxides such as BeO, Al2O3, ZnO and PbO react with both acids and bases to form salts. They are typical of elements near the metal-nonmetal boundary of the periodic table.

Diborane B2H6

A boron hydride with two bridging hydrogen atoms and four terminal hydrogens. Its bonding involves three-centre two-electron banana bonds, so it is electron-deficient and dimerises readily.

Crystal field splitting

In an octahedral complex the five d orbitals split into a lower t2g set and a higher eg set separated by Δo. Strong-field ligands such as CN- give large splitting and low-spin complexes.

Chelate effect

Polydentate ligands form extra-stable complexes because ring formation releases several solvent molecules, increasing entropy. EDTA, a hexadentate ligand, is a classic example.

SN1 versus SN2

SN1 proceeds via a carbocation in two steps with racemisation, favoured by tertiary halides and polar protic solvents. SN2 is a single bimolecular step with backside attack and inversion, favoured by primary halides.

Markovnikov's rule

In electrophilic addition of HX to an unsymmetrical alkene, hydrogen adds to the carbon bearing more hydrogens, giving the more stable carbocation intermediate. Peroxides reverse this for HBr only.

Tollens' test

Ammoniacal silver nitrate gives a silver mirror with aldehydes but not with ketones, because aldehydes are easily oxidised. It distinguishes ethanal from propanone in the laboratory.