chem

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 under standard conditions: $X(g) \rightarrow X^+(g) + e^-$.

Relative atomic mass ($A_r$)

The weighted mean mass of an atom of an element compared to $\frac{1}{12}\text{th}$ of the mass of an atom of carbon-12.

Atomic orbital

A region around the nucleus of an atom that can hold up to two electrons with opposite spins.

Electronegativity

The ability of an atom to attract the shared pair of electrons in a covalent bond.

Metallic bonding

The electrostatic attraction between a regular lattice of positive metal ions and the surrounding sea of delocalised electrons.

Empirical formula

The simplest whole-number ratio of atoms of each element present in a compound.

Percentage atom economy

$$\frac{\text{molar mass of desired product}}{\text{total molar mass of all reactants}} \times 100$$

Homolytic fission

The breaking of a covalent bond where each bonded atom takes one of the shared pair of electrons, forming two free radicals.

Electrophile

An electron-deficient species capable of accepting a lone pair of electrons to form a covalent bond.

Markovnikov's rule

In the electrophilic addition of a hydrogen halide to an unsymmetrical alkene, the hydrogen atom attaches to the carbon atom that already possesses the greater number of hydrogen atoms, yielding the more stable carbocation intermediate.

Mole

The amount of substance containing the same number of particles as there are atoms in exactly $12\,g$ of carbon-12, equal to $6.02 \times 10^{23}$ particles.

Avogadro constant ($L$)

The number of constituent particles (atoms, ions, or molecules) per mole of a substance, which is $6.02 \times 10^{23}\,mol^{-1}$.

Molar mass ($M$)

The mass of one mole of a chemical substance, expressed in units of $g\,mol^{-1}$.

Relative atomic mass ($A_r$)

The weighted mean mass of an atom of an element compared to $\frac{1}{12}\text{th}$ of the mass of an atom of carbon-12.

Relative molecular mass ($M_r$)

The weighted mean mass of a molecule of a compound compared to $\frac{1}{12}\text{th}$ of the mass of an atom of carbon-12.

Empirical formula

The simplest whole-number ratio of atoms of each element present in a compound.

Molecular formula

The actual number of atoms of each element present in one molecule of a compound.

Ideal gas equation

$pV = nRT$, where $p$ is pressure in $Pa$, $V$ is volume in $m^3$, $n$ is moles, $R$ is the ideal gas constant, and $T$ is temperature in $K$.

Molar gas volume at RTP

The volume occupied by one mole of any gas at room temperature and pressure ($293\,K$ and $101.3\,kPa$), which is $24.0\,dm^3\,mol^{-1}$.

Percentage yield

$$\frac{\text{actual yield}}{\text{theoretical yield}} \times 100$$

Percentage atom economy

$$\frac{\text{molar mass of desired product}}{\text{total molar mass of all reactants}} \times 100$$

Concentration of a solution

The amount of solute in moles dissolved per cubic decimeter of solution ($mol\,dm^{-3}$).

Standard solution

A solution of accurately known concentration used in volumetric analysis.

Parts per million (ppm) by mass

$$\frac{\text{mass of solute}}{\text{mass of solution}} \times 10^6$$

Stoichiometry

The molar ratio of reactants and products shown in a balanced chemical equation.

Limiting reactant

The substance that is completely consumed first in a chemical reaction, determining the maximum yield of product.

Water of crystallisation

Water molecules chemically bonded into the regular crystalline lattice of a hydrated compound.

Hydrated salt

A crystalline salt containing chemically incorporated water of crystallisation.

Anhydrous salt

A salt from which all water of crystallisation has been removed.

Percentage uncertainty

$$\frac{\text{uncertainty of apparatus}}{\text{measurement value}} \times 100$$

Atomic number ($Z$)

The number of protons found in the nucleus of an atom of an element.

Mass number ($A$)

The total number of protons and neutrons located in the nucleus of an atom.

Isotopes

Atoms of the same element containing identical numbers of protons but different numbers of neutrons.

Relative isotopic mass

The mass of an individual atom of an isotope relative to $\frac{1}{12}\text{th}$ of the mass of an atom of carbon-12.

Mass spectrometry: Electron impact ionisation

High-energy electrons fired from an electron gun collide with gaseous atoms, knocking out an electron: $M(g) + e^- \rightarrow M^+(g) + 2e^-$.

Mass spectrometry: Acceleration

Positive ions are accelerated by an electric field so that all ions attain the same kinetic energy.

Mass spectrometry: Ion drift (Time of Flight)

Ions travel along a flight tube without an electric field; lighter ions travel faster than heavier ions.

Mass spectrometry: Detection

Positive ions hit a detector plate, receive electrons, and produce a current proportional to ion abundance.

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: $X(g) \rightarrow X^+(g) + e^-$.

Second ionisation energy

The energy required to remove one mole of electrons from one mole of gaseous $1+$ ions to form one mole of gaseous $2+$ ions: $X^+(g) \rightarrow X^{2+}(g) + e^-$.

General trend in first ionisation energy across a period

Increases across a period due to increased nuclear charge and reduced atomic radius with similar electron shielding.

General trend in first ionisation energy down a group

Decreases down a group due to increased atomic radius and greater electron shielding outweighing the increase in nuclear charge.

First ionisation energy anomaly between Group 2 and Group 3 ($Mg$ to $Al$)

The removed electron in Group 3 occupies a higher-energy $3p$ orbital which is shielded by the $3s$ subshell, lowering the required energy.

First ionisation energy anomaly between Group 5 and Group 6 ($N$ to $O$)

The removed electron in Group 6 is paired in a $2p$ orbital, where spin repulsion between paired electrons reduces the energy required for removal.

Atomic orbital

A region of space around the nucleus that can hold up to two electrons with opposite spins.

s-orbital characteristics

A spherical orbital with a maximum electron capacity of two.

p-orbital characteristics

A dumbbell-shaped orbital; three degenerate p-orbitals exist per principal shell ($p_x, p_y, p_z$), holding up to six electrons in total.

Pauli exclusion principle

An orbital can hold a maximum of two electrons, and they must possess opposite spins.

Hund's rule

Electrons occupy degenerate orbitals singly with parallel spins before pairing up to minimise electron-electron repulsion.

Aufbau principle

Electrons enter and fill the lowest-energy available orbitals first before occupying higher levels.

Electron configuration of Chromium ($Cr$, $Z=24$)

$1s^2 2s^2 2p^6 3s^2 3p^6 3d^5 4s^1$, forming a half-filled $3d$ subshell for extra stability.

Electron configuration of Copper ($Cu$, $Z=29$)

$1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^1$, forming a completely filled $3d$ subshell for extra stability.

Ionic bonding

The electrostatic attraction between oppositely charged ions arranged in a giant ionic lattice.

Factors strengthening ionic bonding

Higher ionic charge and smaller ionic radius, which lead to higher charge density and stronger electrostatic attraction.

Polarising power of a cation

The ability of a cation to attract and distort the electron cloud of a neighbouring anion.

Polarisability of an anion

The ease with which the electron cloud of an anion can be distorted by a cation, which increases with larger ionic radius and higher negative charge.

Fajans' rules

Covalent character in an ionic compound increases with a cation of high charge and small radius, and an anion of high charge and large radius.

Covalent bond

The strong electrostatic attraction between two positive atomic nuclei and a shared pair of electrons.

Dative covalent (coordinate) bond

A covalent bond where both electrons in the shared pair originate from the same bonded atom.

Condition for dative bond formation

One species must have a lone pair of electrons, and the other must have an empty orbital to accept the pair.

Electronegativity

The ability of an atom to attract the bonding pair of electrons in a covalent bond.

Most electronegative element

Fluorine, which holds a value of $4.0$ on the Pauling electronegativity scale.

Non-polar covalent bond

A bond between atoms with identical or near-identical electronegativities, giving symmetric electron density.

Polar covalent bond

A bond formed between atoms of different electronegativities, causing an asymmetric electron distribution and a permanent dipole.

London (dispersion) forces

Weak intermolecular attractions caused by instantaneous, temporary dipoles that induce complementary dipoles in neighbouring particles.

Factors increasing London force strength

An increased total number of electrons and greater surface contact area between unbranched molecules.

Permanent dipole-dipole interactions

Weak electrostatic attractions between the permanent delta-positive end of one polar molecule and the delta-negative end of an adjacent polar molecule.

Hydrogen bonding

A strong intermolecular attraction between a hydrogen atom covalently bonded to $N$, $O$, or $F$ and a lone pair on a neighbouring $N$, $O$, or $F$ atom.

Anomalously high boiling point of water

Caused by extensive hydrogen bonding between water molecules, which requires substantial energy to break relative to other group 16 hydrides.

Anomalous low density of ice

Hydrogen bonds hold water molecules in an open, hexagonal crystal lattice that spaces molecules further apart than in the liquid phase.

Metallic bonding

The electrostatic attraction between a regular lattice of positive metal ions and delocalised valence electrons.

Electrical conductivity in metals

Mobile delocalised electrons flow through the metallic lattice when an electrical potential difference is applied.

Structure and bonding of diamond

A giant covalent lattice where each carbon atom forms four strong covalent bonds in a tetrahedral arrangement with $109.5^\circ$ bond angles.

Structure and bonding of graphite

Planar hexagonal layers of carbon atoms bonded to three others, with delocalised electrons between layers and weak London forces holding layers together.

Graphene

A single two-dimensional layer of carbon atoms arranged in a hexagonal honeycomb lattice with delocalised pi electrons.

Hydrocarbon

A compound containing only carbon and hydrogen atoms.

Saturated hydrocarbon

A hydrocarbon containing only single carbon-carbon ($C-C$) bonds.

Unsaturated hydrocarbon

A hydrocarbon containing one or more multiple carbon-carbon bonds ($C=C$ or $C\equiv C$).

Homologous series

A series of organic compounds with identical functional groups, the same general formula, similar chemical properties, and each differing by $-CH_2-$.

Functional group

An atom or group of atoms in an organic molecule primarily responsible for its specific chemical reactions.

General formula of alkanes

$$C_n H_{2n+2}$$

General formula of cycloalkanes

$$C_n H_{2n}$$

General formula of alkenes with one double bond

$$C_n H_{2n}$$

Structural formula

A chemical representation showing the arrangement of atoms in a molecule carbon by carbon without drawing every bond.

Displayed formula

A diagram showing all the atoms and every covalent bond present in a molecule.

Skeletal formula

A simplified drawing where vertices and line ends represent carbon atoms, and hydrogen atoms attached to carbon are omitted.

Structural isomerism

Compounds that have the same molecular formula but different structural arrangements of their atoms.

Chain isomerism

Isomers with the same molecular formula but different branching along the carbon backbone.

Positional isomerism

Isomers sharing the same carbon skeleton and functional group, but with the functional group located at different carbon positions.

Functional group isomerism

Compounds with identical molecular formulae that contain completely different functional groups.

Fractional distillation of crude oil

The continuous physical separation of crude oil into fractions containing hydrocarbons of similar boiling points.

Thermal cracking

The thermal breakdown of long-chain alkanes at high temperatures ($700-1000\,^\circ C$) and high pressures ($7000\,kPa$) to yield high alkene fractions.

Catalytic cracking

The breakdown of long-chain alkanes at $450\,^\circ C$ using a zeolite catalyst to yield branched alkanes, cycloalkanes, and aromatic compounds.

Complete combustion of an alkane

Reaction of an alkane with excess oxygen producing exclusively carbon dioxide and water: $C_n H_{2n+2} + \left(\frac{3n+1}{2}\right)O_2 \rightarrow nCO_2 + (n+1)H_2O$.

Incomplete combustion hazard

Produces carbon monoxide ($CO$), a colourless, odourless toxic gas that binds irreversibly to haemoglobin, preventing oxygen transport.

Free radical

An atom, ion, or molecule with an unpaired valence electron, rendering it chemically highly reactive.