CAPE Chemistry Unit 1 flashcards

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

Avogadro constant

The number of particles (atoms, molecules or ions) in one mole of a substance, equal to 6.02 × 10^23 mol⁻¹. Used to convert between moles and actual particle counts.

Molar volume of gas

The volume occupied by one mole of any gas under specified conditions: 22.4 dm³ at STP (273 K, 1.01 × 10^5 Pa) and 24.0 dm³ at RTP (298 K, 1 atm).

Empirical formula

The simplest whole-number ratio of atoms of each element present in a compound. Found by dividing percentage composition by atomic mass, then dividing by the smallest ratio.

Limiting reagent

The reactant that is completely consumed first and therefore determines the maximum amount of product formed. Identified by dividing moles by stoichiometric coefficient and taking the smallest value.

Atom economy

The mass of the desired product expressed as a percentage of the total mass of all reactants, calculated from relative molecular masses in the balanced equation. Measures how efficiently atoms are used.

Ideal gas equation

pV = nRT, where p is pressure in pascals, V is volume in cubic metres, n is moles, R = 8.31 J K⁻¹ mol⁻¹ and T is absolute temperature in kelvin.

Relative atomic mass

The weighted mean mass of the naturally occurring isotopes of an element compared with one-twelfth of the mass of a carbon-12 atom. A dimensionless quantity.

Isotopes

Atoms of the same element with the same number of protons but different numbers of neutrons. They have identical chemical properties but different masses and physical properties.

First ionisation energy

The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions. Increases across a period and decreases down a group.

Hund's rule

When filling degenerate orbitals of equal energy, electrons occupy separate orbitals singly with parallel spins before any orbital receives a second electron. Explains the 2p³ configuration of nitrogen.

Lattice energy

The enthalpy change when one mole of an ionic solid is formed from its gaseous ions. Its magnitude increases with higher ionic charges and smaller ionic radii.

Hydrogen bonding

A strong intermolecular attraction between a hydrogen atom covalently bonded to N, O or F and a lone pair on a nearby N, O or F atom. Explains the high boiling point of water.

Electronegativity

The tendency of an atom in a covalent bond to attract the shared pair of electrons towards itself. Increases across a period and decreases down a group.

Metallic bonding

Electrostatic attraction between a lattice of positive metal ions and a sea of delocalised valence electrons. Explains electrical conductivity and malleability of metals.

Periodic table groups

Vertical columns of elements whose atoms have the same number of outer-shell electrons. Group number correlates with valence electron count and hence similar chemical properties.

Amphoteric oxide

An oxide that reacts with both acids and bases to form salts, such as aluminium oxide and zinc oxide. Occurs at the metals/non-metals boundary of Period 3.

Transition metal

A d-block element that forms at least one stable ion with a partially filled d subshell. Shows variable oxidation states, coloured complexes and catalytic activity.

Halogen displacement

A reaction in which a halogen higher in Group VII displaces a halide ion of a halogen lower in the group from solution, because oxidising power decreases down the group.

Diagonal relationship

The similarity in chemical behaviour between certain elements diagonally adjacent in the periodic table, such as lithium and magnesium, arising from comparable charge-to-radius ratios.

Enthalpy of formation

The enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions. Elements in standard states have zero enthalpy of formation.

Hess's law

The total enthalpy change of a chemical reaction is independent of the route taken, provided initial and final conditions are the same. Enables calculation of unmeasurable enthalpy changes.

Bond enthalpy

The average energy required to break one mole of a particular covalent bond in the gaseous state. Used to estimate reaction enthalpy: bonds broken minus bonds formed.

Calorimetry equation

q = mcΔT, where q is heat energy transferred, m is mass of solution, c is specific heat capacity (4.18 J g⁻¹ K⁻¹ for water) and ΔT is temperature change.

Born–Haber cycle

An energy cycle applying Hess's law to ionic compound formation, combining atomisation, ionisation energy, electron affinity and lattice enthalpy to calculate an unknown enthalpy change.

Gibbs free energy

ΔG = ΔH − TΔS. A negative ΔG indicates that a reaction is thermodynamically feasible; a positive value indicates it is not feasible at that temperature.

Rate equation

An expression of the form rate = k[A]^m[B]^n, where m and n are reaction orders determined experimentally. The overall order is the sum of m and n.

Half-life

The time taken for the concentration of a reactant to fall to half its initial value. For a first-order reaction it is constant and independent of concentration.

Activation energy

The minimum energy that colliding particles must possess for a reaction to occur. Catalysts lower activation energy by providing an alternative reaction pathway.

Le Chatelier's principle

When a system at equilibrium is disturbed by a change in concentration, pressure or temperature, the equilibrium shifts to oppose that change and restore equilibrium.

Equilibrium constant K_c

The ratio of product concentrations raised to their stoichiometric powers to reactant concentrations raised to their powers, at equilibrium. Depends only on temperature.

Brønsted–Lowry acid

A proton donor. After donation it forms its conjugate base, which differs by one hydrogen ion and can accept a proton back in the reverse reaction.

pH

The negative logarithm to base ten of the hydrogen ion concentration in mol dm⁻³. A decrease of one pH unit corresponds to a tenfold increase in hydrogen ion concentration.

Buffer solution

A solution that resists changes in pH when small amounts of acid or base are added. Usually a weak acid with its conjugate base; pH = pK_a + log([salt]/[acid]).

Standard electrode potential

The potential of a half-cell measured against the standard hydrogen electrode (0.00 V) under standard conditions of 298 K and 100 kPa with 1 mol dm⁻³ solutions.

Oxidation number

A number assigned to an element in a compound representing the electrons lost or gained. Increases during oxidation and decreases during reduction; used to balance redox equations.