IB Chemistry Comprehensive Vocabulary Review

Vocabulary flashcards covering core definitions, key principles, theories, and concepts across all units of the IB Chemistry curriculum based on the lecture notes.

Homogeneous Mixture

A mixture with a uniform composition throughout, having no visible phases or boundaries, where all components are equally distributed in the same physical state.

Heterogeneous Mixture

A mixture with a non-uniform composition that contains visible phases or boundaries separating its components.

Recrystallisation

A purification technique used to separate solid impurities from a compound by dissolving the impure mixture in a minimal volume of hot solvent, filtering off insoluble impurities, and cooling the solution to form pure crystals.

Relative Atomic Mass ($A_r$)

The weighted average mass of an atom of an element relative to one-twelfth the mass of a carbon-12 atom, calculated from isotopic masses and their relative percentage abundances.

Aufbau Principle

The principle stating that electrons occupy the lowest energy atomic orbitals available before filling higher energy orbitals.

Pauli Exclusion Principle

The principle stating that an atomic orbital can hold a maximum of two electrons, and these two electrons must have opposite spins.

Hund's Rule

The rule stating that when filling degenerate orbitals of equal energy, electrons occupy each orbital singly with parallel spins before pairing up in any orbital.

First Ionization Energy

The minimum energy required to remove one mole of electrons from one mole of gaseous atoms in their ground state to form one mole of gaseous $+1$ cations.

Avogadro's Law

A gas law stating that equal volumes of all gases measured at the same temperature and pressure contain equal numbers of gas particles.

Ionic Bond

The non-directional electrostatic attraction between oppositely charged positive cations and negative anions in a crystal lattice structure.

Covalent Bond

The electrostatic attraction between positive atomic nuclei and shared pairs of negatively charged valence electrons.

Coordination Bond

A covalent bond (also known as a dative bond) in which both shared bonding electrons are donated by a single atom.

Electronegativity

A relative measure of an atom's ability to attract a shared pair of electrons in a covalent bond toward itself.

London Dispersion Forces

Temporary intermolecular forces existing between all atoms and molecules caused by instantaneous fluctuations in electron density that induce temporary dipoles in neighboring species.

Hydrogen Bonding

An intermolecular attraction occurring between a hydrogen atom covalently bonded to a highly electronegative atom ($F$, $O$, or $N$) and a lone pair of electrons on a nearby electronegative atom.

Formal Charge

The theoretical charge assigned to an atom in a molecule, calculated as $FC = V - (B + L)$, where $V$ is the number of valence electrons, $B$ is the number of bonding electron pairs, and $L$ is the number of non-bonding lone electrons.

Sigma Bond ($σ$)

A covalent bond formed by the direct head-on overlap of atomic orbitals along the internuclear axis, concentrating electron density between the nuclei.

Pi Bond ($π$)

A covalent bond formed by the lateral or sideways overlap of parallel atomic $p$ orbitals, concentrating electron density above and below the internuclear axis.

Hybridisation

The mathematical mixing of atomic orbitals (such as $s$ and $p$ orbitals) within an atom to generate a set of equivalent hybrid orbitals optimized for covalent bonding.

Metallic Bond

The non-directional electrostatic attraction between a lattice of positive metal cations and a surrounding sea of delocalized valence electrons.

Interstitial Alloy

An alloy formed when smaller secondary atoms occupy the vacant spaces (interstices) within the metal lattice of a larger base metal.

Effective Nuclear Charge

The net positive electrostatic nuclear charge experienced by valence electrons, taking into account the shielding effect caused by core electrons.

Electron Affinity

The enthalpy change occurring when one mole of gaseous atoms gains one mole of electrons to form one mole of singly charged negative gaseous ions.

Amphoteric Oxide

An oxide that exhibits both acidic and basic properties, capable of reacting with both strong acids and strong bases to form salts (e.g., $Al_2O_3$).

Ligand

A neutral molecule or anion possessing at least one lone pair of electrons that can be donated to form a coordinate covalent bond with a central transition metal ion.

Homologous Series

A family of organic compounds sharing the same functional group and general chemical formula, possessing similar chemical properties and showing a gradual trend in physical properties.

Structural Isomerism

A phenomenon where molecules share the exact same molecular formula but possess different structural arrangements and connectivity of their constituent atoms.

Stereoisomers

Molecules that have the same molecular formula and sequence of bonded atoms, but differ in the three-dimensional spatial orientation of their atoms.

Enantiomers

A pair of optical stereoisomers that are non-superimposable mirror images of each other due to the presence of an asymmetric chiral carbon atom.

Standard Enthalpy of Formation ($Δ H_f^⊖$)

The enthalpy change that occurs when one mole of a compound is formed from its constituent elements in their standard states under standard conditions ($298\,K$ and $100\,kPa$).

Standard Enthalpy of Combustion ($Δ H_c^⊖$)

The enthalpy change that occurs when one mole of a substance is completely burned in oxygen under standard conditions.

Hess's Law

A law stating that the total enthalpy change for a chemical reaction depends only on the initial and final states and is independent of the pathway taken.

Entropy ($S$)

A thermodynamic quantity that measures the degree of disorder, randomness, or energy dispersion within a physical or chemical system.

Gibbs Free Energy Change ($Δ G^⊖$)

A thermodynamic potential used to calculate the maximum amount of reversible work performable by a system, where a negative value ($Δ G^⊖ < 0$) indicates a spontaneous process.

Activation Energy ($E_a$)

The minimum kinetic energy that colliding reactant molecules must possess in order for a chemical reaction to occur.

Rate-Determining Step

The slowest elementary step in a multi-step reaction mechanism, which dictates the overall rate expression and kinetics of the entire reaction.

Le Chatelier's Principle

A principle stating that if a dynamic equilibrium is subjected to a change in conditions (concentration, temperature, or pressure), the position of equilibrium will shift to oppose that change.

Brønsted–Lowry Acid

A chemical species capable of donating a proton (a hydrogen ion, $H^+$) to another species.

Lewis Base

A chemical species that acts as an electron pair donor to form a coordinate covalent bond.

Buffer Solution

An aqueous solution consisting of a weak acid and its conjugate base (or a weak base and its conjugate acid) that minimizes changes in pH upon the addition of small amounts of strong acid or base.

Voltaic Cell

An electrochemical cell in which electrical energy is produced from a spontaneous exothermic redox reaction occurring at two separate electrodes.

Electrolytic Cell

An electrochemical cell that uses an external source of electrical energy to drive a non-spontaneous redox reaction.

Nucleophile

An electron-rich atom or group of atoms containing a lone pair of electrons or pi bond that is attracted to electron-deficient (positively charged) atomic centers.

Electrophile

An electron-deficient atom, molecule, or ion that is attracted to regions of high electron density and accepts an electron pair to form a covalent bond.

Homolytic Fission

The symmetrical cleavage of a covalent bond where each bonded atom retains one electron from the shared pair, producing two neutral free radicals.

Heterolytic Fission

The unsymmetrical cleavage of a covalent bond where the more electronegative atom takes both shared bonding electrons, producing a cation and an anion.