H3 Chemistry Definitions

A reference for H3 Chemistry (9813): the additional spectroscopy, stereochemistry and mechanism content, followed by the H2 foundations examined with H3. The three-topic H3 roadmap keeps its own course progress; the H2 material remains part of the examined foundation.

Course hub · Formula list

Molecular orbitals and optical spectra

NMR and mass spectrometry

Stereochemistry

Mechanisms and selectivity

Test the distinction

A pure enantiomer and a racemic mixture contain chiral molecules. The mixture is optically inactive because equal and opposite rotations cancel. An achiral meso compound is inactive for a different reason: its own mirror image is superimposable. Use structure and sample composition together.

H2 foundations examined with H3

Open the H2 definitions

Atoms, orbitals and bonding

Term	Meaning	Remember
Isotope	An atom of the same element as another, with a different neutron number.	Proton number fixes the element; nucleon number can differ.
Atomic orbital	A region described by a one-electron wavefunction, associated with a probability distribution for an electron in an atom.	An orbital is not a circular path; it holds at most two electrons with opposite spins.
First ionisation energy	Energy required to remove one mole of electrons from one mole of isolated gaseous atoms to form one mole of gaseous singly positive ions.	X(g) → X + (g) + e⁻; kJ/mol; the gaseous condition matters.
Electronegativity	Ability of an atom in a bond to attract the bonding electrons towards itself.	A comparison within bonding; different from ionisation energy.
Ionic bond	Electrostatic attraction between oppositely charged ions.	Acts throughout an extended lattice.
Covalent bond	Electrostatic attraction between a shared electron pair and the nuclei of bonded atoms.	A bond may be polar even when the whole molecule is non-polar.
Dative covalent bond	Covalent bond whose shared pair is supplied by one atom.	Once formed, it is not a different kind of attraction.
Metallic bond	Attraction between positive metal ions and delocalised electrons.	Delocalisation helps explain conduction and malleability.
Sigma bond	Covalent bond formed by head-on orbital overlap along the internuclear axis.	A carbon–carbon single bond is a sigma bond.
Pi bond	Covalent bond formed by sideways overlap, with electron density on opposite sides of the internuclear axis.	A double bond contains one sigma and one pi bond.
Hydrogen bond	Attraction between hydrogen bonded to N, O or F and a lone pair on N, O or F in the school model.	Stronger than many other intermolecular attractions, but not an O–H covalent bond.
Permanent dipole	Separation of partial positive and negative charge in a molecule.	Molecular shape determines whether bond dipoles cancel.
Instantaneous dipole–induced dipole attraction	Attraction arising when a fluctuating electron distribution induces a dipole in a neighbouring particle.	Present in all atoms and molecules; strength depends on polarisation and contact.
Bond energy	Energy required to break one mole of a specified covalent bond in gaseous species.	Bond breaking is endothermic; average bond energies give estimates.
Bond length	Equilibrium distance between the nuclei of two bonded atoms.	Compare the specified bonds rather than atomic radii alone.

Moles, reacting quantities and analysis

Term	Meaning	Remember
Mole	Amount containing exactly 6.02214076 × 10²³ specified entities.	State whether the entities are atoms, molecules, ions or electrons.
Relative isotopic mass	Mass of an atom of a particular isotope compared with one twelfth of the mass of a carbon-12 atom.	A ratio, with no unit.
Relative atomic mass	Weighted mean mass of atoms of an element compared with one twelfth of the mass of a carbon-12 atom.	Includes isotope abundances.
Relative molecular / formula mass	Sum of relative atomic masses in a molecule / formula unit.	No unit; distinguish from molar mass in g/mol.
Empirical formula	Simplest whole-number ratio of atoms of each element in a compound.	Derived from mole ratios, not mass ratios directly.
Molecular formula	Actual number of atoms of each element in a molecule.	A whole-number multiple of the empirical formula.
Molar concentration	Amount of solute per unit volume of solution.	Common unit mol/dm³; use solution volume, not solvent volume.
Limiting reactant	Reactant exhausted first according to the balanced reacting ratio.	Compare moles divided by their coefficients, not masses directly.
Oxidation	Loss of electrons or increase in oxidation number.	Oxidising agents are themselves reduced.
Reduction	Gain of electrons or decrease in oxidation number.	Reducing agents are themselves oxidised.
Oxidation number	Formal charge assigned to an atom using electron-allocation rules.	Not necessarily the actual charge on that atom.
Endpoint	Experimentally observed signal used to stop a titration.	Should lie close to, but is conceptually different from, the equivalence point.
Equivalence point	Point at which titrant and analyte have reacted in their stoichiometric ratio.	Does not necessarily have pH 7.

Enthalpy and energy

Term	Meaning	Remember
Enthalpy change of reaction	Heat change for a reaction at constant pressure, for the quantities specified by its equation.	State the reaction equation and molar basis.
Standard conditions	Specified reference conditions: standard pressure 100 kPa, stated temperature and species in their standard states.	Standard does not automatically fix the temperature at 298 K; it must be stated.
Standard enthalpy change of formation	Enthalpy change when one mole of a substance forms from its elements in their standard states under standard conditions.	Fractional coefficients are allowed to form exactly one mole.
Standard enthalpy change of combustion	Enthalpy change when one mole of a substance burns completely in oxygen under standard conditions, with all substances in their standard states.	Specify the product states, especially water.
Standard enthalpy change of neutralisation	Enthalpy change when acid and base react to form one mole of water under standard conditions.	The one-mole basis is water, not always acid.
Lattice energy	Enthalpy change when one mole of an ionic solid forms from its separated gaseous ions.	Formation convention: negative; a lattice dissociation value uses the opposite sign.
Hess’s law	Reaction enthalpy depends only on initial and final states, not the route between them.	Reversing an equation reverses its enthalpy sign.
Exothermic / endothermic	A process releasing heat to / absorbing heat from its surroundings.	Negative / positive Δ H; temperature change depends on which part is measured.
Activation energy	Minimum energy needed by reacting particles for a successful reaction along a specified pathway.	A catalyst changes the pathway; it does not change reaction enthalpy.

Rates and equilibrium

Term	Meaning	Remember
Rate of reaction	Change in concentration of a specified reactant or product per unit time, with the sign or stoichiometric convention stated.	Product formation and reactant disappearance can have different numerical rates.
Rate equation	Experimentally determined relationship between rate and reactant concentrations at fixed temperature.	r = k[A]^m[B]ⁿ; powers need not equal overall equation coefficients.
Order with respect to a reactant	Power of that reactant’s concentration in the rate equation.	Overall order is the sum of the powers.
Rate constant	Proportionality constant in a rate equation at a specified temperature.	Units depend on overall order; temperature and catalyst affect its value.
Half-life	Time for the concentration of a specified reactant to fall to half its initial value.	Successive half-lives are constant for a first-order reaction.
Catalyst	Substance that increases rate through an alternative pathway and is regenerated overall.	It changes neither the equilibrium constant nor the equilibrium composition at fixed temperature.
Homogeneous / heterogeneous catalyst	Catalyst in the same phase as / a different phase from the reactants.	Surface adsorption matters in heterogeneous catalysis.
Enzyme	Biological catalyst, usually a protein, with a specific active site.	Temperature and pH can change its activity.
Dynamic equilibrium	State in a closed system where forward and reverse rates are equal and macroscopic composition stays constant.	Reactant and product concentrations need not be equal.
Le Chatelier’s principle	When an equilibrium system is disturbed, its response tends to oppose the imposed change.	A catalyst speeds approach to equilibrium without shifting it.
Equilibrium constant	Value of the equilibrium composition expression for a specified reaction at a specified temperature.	Use equilibrium concentrations or partial pressures; changing temperature can change the constant.

Acids, bases and buffers

Term	Meaning	Remember
Arrhenius acid / base	Substance increasing hydrogen-ion / hydroxide-ion concentration in aqueous solution.	This model is limited to aqueous systems.
Brønsted–Lowry acid / base	Proton donor / proton acceptor.	Identify the species that actually transfers or accepts the proton.
Conjugate acid–base pair	Two species differing by one proton.	Their charges differ by one unit.
Strong / weak acid	Acid that ionises essentially completely / partially in water under the stated conditions.	Strength and concentration describe different things.
Strong / weak base	Base reacting essentially completely / partially with water to generate hydroxide ions, or a hydroxide that dissociates essentially completely / partially.	Compare like species and the appropriate equilibrium.
pH	Negative base-10 logarithm of the numerical hydrogen-ion concentration in mol/dm³ in the school concentration model.	pH measures acidity; it is not a direct measure of total acid concentration.
Acid dissociation constant, Kₐ	Equilibrium constant for an acid transferring a proton to water.	For HA, Kₐ = [H⁺][A⁻]/[HA]; depends on temperature.
Base dissociation constant, K_b	Equilibrium constant for a base reacting with water to produce its conjugate acid and hydroxide ions.	For B, K_b = [BH⁺][OH⁻]/[B].
Ionic product of water, K_w	Product of hydrogen-ion and hydroxide-ion concentrations in the school model at a specified temperature.	Neutrality means equal ion concentrations, not pH 7 at every temperature.
Buffer	Solution that resists changes in pH when small amounts of acid or alkali are added.	Contains appreciable amounts of a weak acid/base and its conjugate partner; capacity is finite.
Acid–base indicator	Weak acid/base whose conjugate forms have different colours.	Choose its transition range within the steep part of the titration curve.
Amphoteric substance	Substance that can react as both an acid and a base.	Identify its role in each particular reaction.

Organic chemistry and polymers

Term	Meaning	Remember
Functional group	Atom or group responsible for characteristic reactions of an organic compound.	The same molecule can have more than one functional group.
Homologous series	Family with the same functional group and general formula, similar chemical reactions and a gradation of physical properties.	Successive members differ by CH₂.
Structural isomers	Compounds with the same molecular formula but different connectivity.	Atoms are joined in different orders.
Stereoisomers	Compounds with the same connectivity but different spatial arrangements.	Compare with structural isomers.
Cis–trans isomerism	Stereoisomerism arising from restricted rotation, with groups on the same or opposite sides of a reference plane.	Each alkene carbon needs two different substituents for alkene cis–trans isomerism.
Addition / substitution	Reaction adding groups across a multiple bond / replacing an atom or group.	Distinguish product connectivity, not just reactant names.
Elimination	Reaction removing atoms or groups to create a multiple bond.	Often competes with substitution for a suitable substrate.
Condensation	Reaction joining molecules with elimination of a small molecule.	Ester and amide formation are examples.
Hydrolysis	Bond cleavage involving reaction with water.	Acidic or alkaline conditions can give different final species.
Addition polymerisation	Joining unsaturated monomers without eliminating small molecules.	The repeat unit is not the unchanged alkene molecule.
Condensation polymerisation	Joining suitable monomers with loss of small molecules.	A difunctional monomer can form a chain.
Denaturation	Loss of a protein’s functional three-dimensional structure.	Does not normally mean hydrolysis of its primary peptide sequence.

Further H2 terms

Term	Meaning	Remember
Ideal gas	Model with negligible particle volume and negligible intermolecular forces except during elastic collisions.	Best approached at low pressure and high temperature.
Partial pressure	Pressure a gas component would exert alone in the mixture’s total volume at the same temperature.	pᵢ = xᵢ pₜₒₜₐₗ for an ideal mixture.
Entropy	Thermodynamic quantity related to the number of microscopic arrangements accessible to a system.	Unit J/(mol K) for molar entropy; particle number, phase and energy distribution matter.
Gibbs free-energy change	Thermodynamic change Δ G = Δ H-TΔ S at a stated temperature.	Negative Δ G favours forward change at constant temperature and pressure; it does not imply a fast reaction.
Standard enthalpy change of atomisation	Enthalpy change forming one mole of gaseous atoms from an element in its standard state.	Molar basis is atoms, not necessarily one mole of the element’s molecules.
Standard enthalpy change of hydration	Enthalpy change when one mole of gaseous ions becomes hydrated aqueous ions under standard conditions.	Usually exothermic; stronger charge density strengthens hydration.
Standard enthalpy change of solution	Enthalpy change dissolving one mole of a substance in sufficient solvent under standard conditions.	Combine the specified lattice and hydration conventions consistently.
Electron affinity	Enthalpy change when one mole of gaseous atoms gains one mole of electrons to form gaseous anions.	First and successive electron affinities need not have the same sign.
Rate-determining step	Slow step that controls the rate in an appropriate multi-step mechanism.	A proposed mechanism must also reproduce the observed rate law.
Lewis acid / base	Electron-pair acceptor / electron-pair donor.	Forms an adduct through a dative bond; no proton is required.
pKₐ	Negative base-10 logarithm of the numerical Kₐ value in the stated concentration convention.	Lower pKₐ corresponds to larger Kₐ and stronger acid.
Solubility product, Kₛₚ	Equilibrium product of ion concentrations, raised to stoichiometric powers, for a saturated solution in contact with its solid.	Pure solid is omitted; solubility is not generally equal to Kₛₚ.
Common-ion effect	Suppression of an ionisation or dissolution equilibrium by adding a species that supplies an ion already present.	Account for the added ion before evaluating equilibrium concentrations.
Standard electrode potential	Electrode potential of a specified half-cell relative to the standard hydrogen electrode under standard conditions.	Reduction potentials are tabulated; reverse the chemistry when an electrode is oxidised.
Standard hydrogen electrode	Reference half-cell of hydrogen gas at standard pressure, hydrogen-ion concentration 1 mol/dm³ and a platinum electrode, at stated temperature.	Assigned potential 0 V.
Electrolysis	Chemical change at electrodes driven by an external electric current.	Oxidation occurs at the anode; reduction at the cathode.
Faraday constant	Charge per mole of electrons.	Approximately 9.65 × 10⁴ C/mol.
Transition element	Element whose atom or at least one of its ions has an incomplete d subshell.	Zinc is not a transition element under this definition.
Ligand	Species that donates an electron pair to a central metal atom or ion to form a coordinate bond.	A multidentate ligand can donate through several donor atoms.
Coordination number	Number of donor atoms directly bonded to the central metal atom or ion.	Count donor atoms, not merely ligand molecules.
Complex ion	Charged species containing a central metal atom/ion coordinated to ligands.	Charge belongs to the whole species, including the ligands.
Ligand exchange	Replacement of one ligand by another at a central metal.	Need not change metal oxidation number.
Electrophile	Species that accepts an electron pair to form a covalent bond.	Curly arrows point towards its receiving atom or bond.
Nucleophile	Species that donates an electron pair to form a covalent bond.	May be negatively charged or neutral with a lone pair.
Free radical	Species with an unpaired electron.	Use single-headed arrows for single-electron movement.
Heterolytic / homolytic fission	Bond cleavage giving both bonding electrons to one atom / one electron to each atom.	Produces ions / radicals in the simple models.
Chiral molecule	Molecule not superimposable on its mirror image.	One tetrahedral carbon with four different groups is a common cause, not a universal test.
Enantiomers	Non-superimposable mirror-image stereoisomers.	A pure enantiomer rotates plane-polarised light; a racemic mixture has no net rotation.
Racemic mixture	Equal amounts of a pair of enantiomers.	Optical rotations cancel, but the mixture is not a single achiral molecule.
Amino acid	Organic compound containing amino and carboxyl functional groups.	Speciation depends on pH; a zwitterion has both positive and negative charges.
Zwitterion	Species with positive and negative charges at different sites and zero overall charge.	It is not a neutral structure without formal charges.
Peptide bond	Amide linkage joining amino-acid residues.	Condensation forms it; hydrolysis breaks it.

H3 stopping point

Use the supplied Data Booklet for characteristic IR absorptions. Quantitative LCAO and mathematical steady-state treatment are not required. The NMR energy discussion is qualitative; Fischer projections, fac/mer identification and solvent effects in the SN1/SN2 comparison are outside the specified additional scope.

Return to the H3 course hub for spectra, structures and mechanism reasoning.