Physical Chemistry Review Flashcards

Flashcards for reviewing key concepts in physical chemistry, covering atomic structure, amount of substance, bonding, energetics, kinetics, equilibria, and acids/bases.

Atom

The smallest unit of an element that retains the chemical properties of that element.

Isotope

Atoms of the same element that have the same atomic number but different mass numbers due to a different number of neutrons.

Mass number

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

Atomic number

The number of protons in the nucleus of an atom, which determines the element.

Relative atomic mass (Ar)

The weighted average mass of the atoms of an element, taking into account the relative abundance of its isotopes.

Relative isotopic mass

The mass of a specific isotope relative to 1/12th the mass of a carbon-12 atom.

Mass spectrometry

An analytical technique used to measure the mass-to-charge ratio of ions, allowing for the determination of isotopic abundances and relative atomic masses.

Ionisation energy

The energy required to remove one mole of electrons from one mole of gaseous atoms or ions.

Electron configuration

The arrangement of electrons in the various energy levels and subshells within an atom.

Nuclear charge

The positive charge experienced by an electron in an atom, taking into account the shielding effect of inner electrons.

Shielding

The reduction of the effective nuclear charge experienced by an electron due to the presence of other electrons.

Atomic radius

The distance from the nucleus to the outermost electron shell of an atom.

Sub-shells (s, p, d, f)

Regions of space within an electron shell that can hold electrons with specific shapes and energies (s, p, d, f).

Mole (mol)

The SI unit of amount of substance, defined as the amount containing the same number of entities as there are atoms in 12 grams of carbon-12.

Avogadro's constant

The number of entities (atoms, molecules, ions, etc.) in one mole of a substance, approximately 6.022 x 10^23.

Molar mass (Mr)

The mass of one mole of a substance, expressed in grams per mole (g/mol).

Empirical formula

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

Molecular formula

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

Concentration (mol dm³)

The amount of solute dissolved in a given volume of solvent, typically expressed in moles per cubic decimeter (mol/dm^3).

Volume (dm³)

A measure of the three-dimensional space occupied by a substance, often expressed in cubic decimeters (dm^3).

Ideal gas equation

An equation that relates the pressure, volume, temperature, and number of moles of an ideal gas: PV = nRT.

Stoichiometry

The quantitative relationship between reactants and products in a chemical reaction.

Percentage yield

The ratio of the actual yield of a product to the theoretical yield, expressed as a percentage.

Atom economy

A measure of how many of the starting atoms end up in the desired product.

Limiting reagent

The reactant that is completely consumed in a chemical reaction, determining the amount of product that can be formed.

Ionic bonding

The electrostatic attraction between oppositely charged ions.

Covalent bonding

The sharing of electrons between atoms.

Metallic bonding

The attraction between delocalized electrons and positively charged metal ions.

Dative covalent bond

A covalent bond in which both electrons are provided by one atom.

Electronegativity

A measure of the ability of an atom to attract electrons in a chemical bond.

Polar bond

A covalent bond in which the electron density is unevenly distributed, resulting in a partial positive and partial negative charge.

Intermolecular forces

Weak attractive forces between molecules, including van der Waals forces, dipole-dipole interactions, and hydrogen bonding.

Giant ionic lattice

A three-dimensional network of ions held together by strong electrostatic forces.

Giant covalent lattice

A three-dimensional network of atoms held together by strong covalent bonds.

Simple molecular

Molecules held together by weak intermolecular forces.

Metallic structure

A regular arrangement of metal ions surrounded by a sea of delocalized electrons.

Bond enthalpy

The enthalpy change required to break one mole of a particular bond in the gaseous phase.

Bond angle

The angle between two bonds that originate from the same atom in a molecule.

Electron pair repulsion theory

A theory that predicts the shape of molecules based on the repulsion between electron pairs around a central atom.

Enthalpy change (ΔH)

The heat energy exchanged with the surroundings during a chemical reaction at constant pressure.

Exothermic

A reaction that releases heat to the surroundings (ΔH < 0).

Endothermic

A reaction that absorbs heat from the surroundings (ΔH > 0).

Standard conditions

A set of specified conditions (298 K and 100 kPa) used for comparing thermodynamic data.

Standard enthalpy of formation (ΔHf°)

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

Standard enthalpy of combustion (ΔHc°)

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

Hess's law

A law stating that the enthalpy change for a reaction is independent of the path taken.

Bond dissociation enthalpy

The enthalpy change required to break one mole of a specific bond in the gaseous phase.

Mean bond enthalpy

The average enthalpy change for breaking one mole of a particular type of bond in the gaseous phase, averaged over different molecules.

Enthalpy of solution

The enthalpy change when one mole of a solid ionic compound dissolves in water to give an infinitely dilute solution.

Enthalpy of hydration

The enthalpy change when one mole of gaseous ions dissolves in water to form hydrated ions.

Entropy (S)

A measure of the disorder or randomness of a system.

Free energy (Gibbs free energy, ΔG)

A thermodynamic potential that measures the amount of energy available in a system to do useful work at constant temperature and pressure.

Feasibility / spontaneity

The spontaneity of a reaction which is determined by the change in Gibbs free energy (ΔG). A reaction is feasible if ΔG < 0.

Rate of reaction

The speed at which reactants are converted into products in a chemical reaction.

Rate equation

An equation that expresses the rate of a reaction in terms of the concentrations of the reactants.

Rate constant (k)

The proportionality constant in a rate equation, which is temperature-dependent.

Order of reaction

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

Half-life

The time it takes for the concentration of a reactant to decrease to half its initial value.

Activation energy (Ea)

The minimum energy required for a reaction to occur.

Maxwell-Boltzmann distribution

A distribution that shows the range of kinetic energies for the molecules in a sample

Catalyst

A substance that speeds up a reaction without being consumed in the process.

Transition state

A high-energy, unstable state formed during a reaction.

Reaction mechanism

A step-by-step sequence of elementary reactions that make up the overall reaction.

Rate-determining step

The slowest step in a reaction mechanism, which determines the overall rate of the reaction.

Dynamic equilibrium

A state in which the forward and reverse reaction rates are equal.

Le Chatelier's principle

A principle stating that if a change of condition is applied to a system in equilibrium, the system will shift in a direction that relieves the stress.

Equilibrium constant (Kc, Kp)

The ratio of product concentrations to reactant concentrations at equilibrium, each raised to the power of their stoichiometric coefficients.

Mole fraction

The ratio of the number of moles of a component to the total number of moles in a mixture.

Partial pressure

The pressure exerted by a single component in a mixture of gases.

Homogeneous / Heterogeneous equilibria

Equilibria in which all reactants and products are in the same phase/Equilibria in which reactants and products are in different phases.

Acid/Base

Substances that donate protons (H+) or accept electron pairs.

Conjugate acid-base pairs

Acids and bases that differ by the presence or absence of a proton.

pH

A measure of the acidity or alkalinity of a solution, defined as the negative logarithm of the hydrogen ion concentration.

pKa

The negative logarithm of the acid dissociation constant (Ka).

pKb

The negative logarithm of the base dissociation constant (Kb).

Kw (ionic product of water)

The ionic product of water, defined as [H+][OH-], which is 1.0 x 10^-14 at 298 K.

Strong / Weak acids and bases

Acids/bases that completely dissociate in solution and acids/bases that partially dissociate in solution.

Neutralisation

The reaction between an acid and a base, resulting in the formation of a salt and water.

Buffer solution

A solution that resists changes in pH upon the addition of small amounts of acid or base.

Titration curve

A graph showing the change in pH during a titration.