NCEA L3 Chemistry flashcards

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

Equilibrium constant Kc

For aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d / ([A]^a[B]^b). Pure solids and liquids are excluded. Kc depends only on temperature, not on concentrations or pressure.

Reaction quotient Q

The same expression as Kc but using current, non-equilibrium concentrations. If Q < Kc the reaction proceeds forward; if Q > Kc it proceeds in reverse; if Q = Kc the system is at equilibrium.

Le Chatelier's principle

When a system at equilibrium is disturbed, it shifts in the direction that opposes the change. Adding reactant shifts right; increasing pressure shifts toward fewer gas moles; heating favours the endothermic direction.

Catalyst effect on equilibrium

A catalyst lowers the activation energy and speeds up both forward and reverse reactions equally. It shortens the time to reach equilibrium but does not change the equilibrium position or the value of Kc.

Weak acid Ka expression

Ka = [H3O+][A-]/[HA]. For a weak monoprotic acid of concentration c, [H3O+] ≈ sqrt(Ka × c), so pH ≈ ½(pKa − log c). Smaller Ka means a weaker acid.

Buffer solution

A solution of a weak acid and its conjugate base (or weak base and conjugate acid) that resists pH change on adding small amounts of acid or base. pH ≈ pKa + log([base]/[acid]).

Titration equivalence point pH

At 25 °C, strong acid–strong base gives pH 7; weak acid–strong base gives pH above 7 because the conjugate base hydrolyses; weak base–strong acid gives pH below 7 because the conjugate acid hydrolyses.

Solubility product Ks

For AgCl(s) ⇌ Ag+(aq) + Cl-(aq), Ks = [Ag+][Cl-]. A precipitate forms when the ion product exceeds Ks. Ks depends on temperature and is unaffected by adding common ions except through the ion product.

Gibbs free energy

ΔG = ΔH − TΔS. A reaction is spontaneous when ΔG is negative. Endothermic reactions (ΔH > 0) can be spontaneous if ΔS is positive and the temperature is high enough for TΔS to exceed ΔH.

Oxidation number rules

Free elements are 0; a monatomic ion equals its charge; oxygen is usually −2 (peroxide −1); hydrogen is usually +1 (hydride −1). The sum of oxidation numbers equals the overall charge of the species.

Oxidation and reduction

Oxidation is loss of electrons and an increase in oxidation number; reduction is gain of electrons and a decrease in oxidation number. The oxidant is reduced and the reductant is oxidised.

Standard electrode potential E°

A measure of the tendency of a half-cell to be reduced, measured against the standard hydrogen electrode at 1 mol L-1, 1 atm, 25 °C. More positive E° means a stronger oxidant.

Cell potential E°cell

E°cell = E°cathode − E°anode. A positive value means the cell reaction is spontaneous as written. In both galvanic and electrolytic cells, oxidation occurs at the anode and reduction at the cathode.

Faraday's law of electrolysis

Moles of electrons = Q/F, where F = 96 485 C mol-1. The mass deposited is (Q × M)/(n × F), where M is molar mass and n is the number of electrons per ion.

Tollens' and Fehling's tests

Both detect aldehydes, which are oxidised to carboxylate ions. Tollens' reagent gives a silver mirror; Fehling's solution gives a brick-red precipitate of copper(I) oxide. Ketones do not react.

Bromine water test

Bromine water is decolourised by alkenes and alkynes through electrophilic addition across the double or triple bond. Alkanes do not decolourise it under normal conditions without UV light.

Amine basicity

Amines are basic because the nitrogen lone pair accepts a proton. Amides are far weaker bases because the lone pair is delocalised into the adjacent carbonyl group, reducing its availability.

Esterification and hydrolysis

Carboxylic acid plus alcohol gives an ester and water; the reaction is acid-catalysed and reversible. Ester hydrolysis is base-promoted (saponification, irreversible) or acid-catalysed (reversible).

Reaction order from data

Compare experiments where one reactant concentration changes while others stay constant. If doubling that concentration doubles the rate, the reaction is first order in that reactant; if it quadruples the rate, second order.

Rate constant units

Units of k depend on overall order: zero order mol L-1 s-1; first order s-1; second order mol-1 L s-1; third order mol-2 L2 s-1. Overall order is the sum of individual orders.

First-order half-life

For a first-order reaction, t½ = ln2/k ≈ 0.693/k. The half-life is constant and independent of initial concentration, which is a distinguishing feature of first-order kinetics.

Arrhenius equation

k = A e^(−Ea/RT). A plot of ln k against 1/T is a straight line with slope −Ea/R and intercept ln A. A larger activation energy Ea makes k more sensitive to temperature.

Rate-determining step

The slowest step in a reaction mechanism. It controls the overall rate, and the experimental rate law involves only species up to and including this step, not products formed afterwards.

Hess's law

The enthalpy change of a reaction is independent of the route taken. Unknown ΔH values can be found by combining known equations, reversing the sign when an equation is reversed and scaling when it is multiplied.

Bond enthalpy calculation

ΔH = Σ(bonds broken) − Σ(bonds formed). Using average bond enthalpies gives an estimate, because tabulated values are averaged across many different molecules and states.

Structural isomerism types

Chain isomers differ in carbon skeleton branching; position isomers differ in the location of a functional group or multiple bond; functional group isomers have different functional groups, such as alcohols and ethers.

Optical isomerism

Occurs when a carbon atom is bonded to four different groups (a chiral centre). The two enantiomers are non-superimposable mirror images and rotate plane-polarised light in opposite directions.

Geometric isomerism

Occurs about a C=C double bond when each carbon carries two different groups. Cis isomers have similar groups on the same side; trans isomers have them on opposite sides. Restricted rotation prevents interconversion.

Hydrogen bonding

A strong dipole–dipole attraction between H bonded to N, O, or F and a lone pair on N, O, or F of another molecule. It raises boiling points, as in water, ammonia, and hydrogen fluoride.

VSEPR molecular shapes

Electron domains repel to maximise separation: 2 domains linear; 3 trigonal planar; 4 tetrahedral (109.5°). Lone pairs compress angles, giving trigonal pyramidal (~107°) or bent (~104.5°) shapes.

Polarity of molecules

A molecule is polar only if it has polar bonds and an asymmetric shape so bond dipoles do not cancel. CO2 is nonpolar (linear); H2O is polar (bent); CCl4 is nonpolar (tetrahedral).

Periodic trends

Atomic radius decreases across a period and increases down a group. First ionisation energy generally increases across a period and decreases down a group, with dips at group 13 and group 16 elements.