Key Chemistry Definitions and Concepts

Empirical formula

the simplest ratio of atoms in a molecule.

Molecular formula

The total number of atoms of each element in a molecule.

Structural formula

A formula which shows the arrangement of atoms in the molecule of a compound.

Displayed formula

the symbols of each atom joined by lines representing all the bonds in the molecule.

Skeletal formula

No hydrogen atoms are shown, just the bonds between the atoms.

Addition

A reaction where two or more molecules combine to give a single product.

Elimination

A reaction where a small molecule such as HCl or H2O is removed from a molecule.

Condensation

A reaction where a small molecule such as HCl or H2O is produced as two molecules join together.

Substitution

A reaction where one atom or group of atoms is replaced by another.

Oxidation

The addition of oxygen, removal of electrons or increase in oxidation state.

Reduction

The removal of oxygen, addition of electrons or decrease in oxidation state.

Hydrolysis

Decomposition reaction with water as one of the reactants.

Nucleophile

A species which donates pairs of electrons - it is usually attracted to positive charge.

Electrophile

A species which accepts pairs of electrons - it is usually attracted to negative charge.

Chiral centre

A carbon atom with 4 different atoms or groups attached.

Enantiomers

A pair of optically active molecules whose mirror images cannot be superimposed.

Racemic mixture

A mixture containing equal amounts of a pair of enantiomers.

Free radical

A species with an unpaired electron.

Homolytic fission

Where a bond splits giving equal numbers of electrons back, producing radicals.

Heterolytic fission

Where a bond splits giving unequal numbers of electrons back, producing ions.

Co-ordinate (dative)

The sharing of a pair of electrons between two atoms where the electrons have come from the same atom.

Electronegativity

The power of an atom that is covalently bonded to attract the bonding pair of electrons towards itself.

σ Bond

Formed through end on end orbital overlap.

π Bond

Formed through side on side overlap of P orbitals, forming electron density clouds above and below the plane of the σ Bond.

Geometric isomers

Formed by the lack of rotation around a double bond, molecules cannot be superimposed on each other.

Structural isomers

Compounds with the same molecular formula but different structural formula.

Chain isomerism

Isomerism caused by changes to the carbon chain.

Optical isomers

Formed when a chiral carbon exists in a molecule giving 2 non-superimposable mirror images.

Lattice energy, Δ Hlatθ

The enthalpy change when 1 mole of an ionic compound is formed from its gaseous ions under standard conditions (25℃, 1 a.t.m).

Standard enthalpy change of atomisation, Δ Hatθ

Enthalpy change when 1 mole of gaseous atom is formed from its element under standard conditions.

Electron affinity, ΔHeaθ

enthalpy change when 1 mole of electrons is added to one mole of atoms or ions in the gaseous state under standard conditions.

Ionisation energy, ΔHI.eθ

the energy required to remove 1 mole of electrons from 1 mole of gaseous atoms of an element to form 1 mole of uniposiƟve (+1) ions.

Hess' law

The total energy change of a reaction is independent of the route taken.

Polarising power

the ability of a cation to attract electrons and distort the electron cloud of an anion.

Standard enthalpy change of solution, Hsolθ

enthalpy change when 1 mole of an ionic solid dissolves in sufficient water to form an infinitely (very dilute) solution.

Standard enthalpy change of hydration, Hhydθ

enthalpy change when 1 mole of a specified gaseous ion dissolves in sufficient water to form a very dilute solution.

Standard enthalpy of neutralisation, Hneutθ

the enthalpy change when one mole of water is formed in the reaction between an acid and a base under standard conditions, 1 atm and 298 K.

Average bond energy

the average energy needed to break a specific covalent bond.

Bond energy

The energy required to break one mole of a particular covalent bond in the gaseous state.

Specific heat capacity

the energy needed to raise 1 g of a substance by 1 °C.

Entropy

measure of dispersal of energy at a specific temperature/ measure of randomness or disorder of a system.

Solubility product, Ksp

the product of the concentrations of each ion in a saturated solution of a sparingly soluble salt at 298 K, raised to the power of their relative concentrations.

Common ion effect

the reduction in the solubility of a dissolved salt achieved by adding a solution of a compound which has an ion in common with the dissolved salt.

Acid (Bronsted-Lowery Theory)

proton donor.

Base (Bronsted-Lowery Theory)

proton acceptor.

Strong acid/base

strong electrolytes that dissociate completely in water.

Weak acid/base

weak electrolytes that dissociate partially in water.

Acid-base indicator

a dye or mixture of dyes that changes colour over a specific pH range.

Buffer solution

solution which resists changes in pH when small quantities of acid or alkali are added.

Kw

[H+][OH-], at 298 K, Kw = 1.0 x 10-14.

pH + pOH

14.

d - block element

An element whose electron configuration differs from the previous element by an electron in a d sub shell.

Transition metal

A d-block element which can form one or more stable ions which have a partially filled d sub shell.

Ligand

is an ion or molecule with a functional group that binds to a central metal atom by forming a dative covalent bond to form a coordination complex.

Bidentate ligand

A ligand which can form 2 dative covalent bonds with the central metal ion.

Polydentate ligand

A ligand which can form more than 2 dative covalent bonds with the central metal ion.

Rate of reaction

the change in the concentration of a reactant or product with time. (unit: mol dm-3 s-1)

Order of reaction w.r.t. a particular reactant

the power to which the concentration of that reactant is raised in the rate equation.

Overall order of reaction

the sum of the powers of the concentration in the experimentally determined rate equation.

Half-life

time required for the concentration of a reactant to decrease to half of its initial concentration.

Reaction mechanism

a sequence of simple steps proposed in the theory to account for the overall chemical reaction that takes place and it must be consistent with the observed kinetics.

Rate-determining step

the slowest step in the sequence of steps leading to product formation.

Intermediate

species that appear in a reaction mechanism but not in the overall balanced equation.

Catalyst

substance that increases the rate of a chemical reaction without itself being consumed.

Homogeneous catalyst

catalyst that is in the same phase as the reaction mixture.

Heterogeneous catalyst

catalyst that is in a different phase to the reaction mixture.

Autocatalyst

product of a chemical reaction that acts as catalyst in the reaction.

Adsorption

molecules become bonded to the atoms on the surface of a solid.

Desorption

the bonds between molecules and atoms on the surface of solid are broken off.

Electrolysis

decomposition of a compound into its elements when an electric current passes through an electrolyte.

Electrolyte

the compound that is decomposed.

Electrode

a rod which conducts electricity to and from the electrolyte.

1 Faraday (F)

the quantity of electric charge carried by 1 mole of single charged ions.

Electrode potential, E

the difference in potential between metal/metal ion system and another system.

Standard electrode potential, Eθ

potential difference between a standard hydrogen electrode and a metal (the electrode) which is immersed in a solution containing metal ions at 1.0 mol dm-3 concentration at 25℃ and at 1 atmospheric pressure.

Standard reduction potential

a measure of a standard half-cell's tendency to accept electron with reference to the standard hydrogen electrode with all containing metal ions at 1.0 mol dm-3 concentration at 25℃ and at 1 atmospheric pressure.

Standard cell potential

the difference between the Eθ values of the two standard half cells with all containing metal ions at 1.0 mol dm-3 concentration at 25℃ and at 1 atmospheric pressure.

Fuel cell

electrochemical cell in which a fuel gives up electrons at one electrode and oxygen gains electrons at the other electrode.

Faraday's 1st Law of Electrolysis

the mass of substance liberated at an electrode during electrolysis is directly proportional to the quantity of electricity passed through the electrolyte.

Faraday's 2nd Law of Electrolysis

the number of faradays required to discharge one mole of an ion at an electrode equals to the number of charges on the ion.

Q

Q = It

Eθcell

Eθcell = Eθcathode - Eθanode = Reduction - Oxidation

The Nernst Equation

E = Eθ + RTzF ln ([Oxidised form]/[Reduced form])

At 298K

E = Eθ + 0.059z log10 ([Oxidised form]/[Reduced form])