KSEAB II PUC Chemistry key terms

The SI base unit for amount of substance, defined as exactly 6.02214076 × 10^23 elementary entities. It links the mass of a sample to the number of particles it contains.

The number of constituent particles in one mole of a substance, fixed at 6.02214076 × 10^23 per mole. It converts between particle counts and amounts in moles.

The mass of one mole of a substance, expressed in grams per mole and numerically equal to its relative molecular or formula mass. It converts mass to moles in stoichiometric calculations.

The reactant that is completely consumed first in a chemical reaction, fixing the maximum amount of product formed. Other reactants remain in excess once it is used up.

The number written before a species in a balanced chemical equation, showing the relative number of moles reacting or produced. Ratios of these coefficients give mole-to-mole conversion factors.

The formula showing the simplest whole-number ratio of atoms of each element in a compound. It is derived from percentage composition and may differ from the molecular formula.

The rule that electrons occupy the lowest-energy available orbital first when an atom is built up. It gives the standard order of orbital filling in the ground state.

The rule that degenerate orbitals of a subshell are singly occupied with parallel spins before any pairing occurs. It minimises electron-electron repulsion and explains observed magnetic behaviour.

No two electrons in an atom can have the same set of all four quantum numbers. Consequently an orbital holds at most two electrons, and they must have opposite spins.

It is impossible to determine both the exact position and the exact momentum of a particle simultaneously; the product of the uncertainties is at least h/4π. It rules out fixed electron orbits.

A set of numbers (principal, azimuthal, magnetic and spin) that specifies the energy, shape, orientation and spin state of an electron in an atom. No two electrons share all four.

The mixing of atomic orbitals of similar energy on the same atom to give an equal number of equivalent hybrid orbitals with identical energy and shape, explaining observed molecular geometry and bond angles.

Valence Shell Electron Pair Repulsion theory predicts molecular shape by arranging electron pairs around the central atom so that repulsions are minimised, with lone pairs occupying more space than bonding pairs.

For a diatomic molecule, half the difference between the number of bonding and antibonding electrons; it indicates bond strength and length. A higher bond order generally means a shorter, stronger bond.

A bonding model in which atomic orbitals combine to form molecular orbitals that extend over the whole molecule. Electrons fill these orbitals by Aufbau, Pauli and Hund rules, allowing prediction of magnetism.

A measure of molecular polarity, equal to the product of charge and separation distance and expressed in debye units. It is a vector sum, so symmetric molecules may have zero net moment.

Energy cannot be created or destroyed, only transferred as heat or work, so the change in internal energy of a system equals q + w. It defines the relation between heat, work and internal energy.

A state function equal to internal energy plus the product of pressure and volume. At constant pressure, the change in enthalpy equals the heat exchanged by the system.

The enthalpy change of a reaction is independent of the route taken, depending only on the initial and final states. It allows calculation of unknown enthalpy changes from known ones.

A thermodynamic state function measuring the dispersal or randomness of energy in a system. Its change, together with enthalpy change and temperature, determines the spontaneity of a process.

A state function defined as G = H − TS; a process is spontaneous at constant temperature and pressure when the change in G is negative. It combines enthalpy and entropy effects into one criterion.

When a system at equilibrium is disturbed by a change in concentration, pressure or temperature, the system shifts in the direction that partially counteracts the change. It predicts the direction of equilibrium shift.

The ratio of product concentrations (or partial pressures) to reactant concentrations, each raised to its stoichiometric coefficient, at equilibrium at a given temperature. Its value indicates the extent of reaction.

A species that accepts electrons from another species and is itself reduced in the process. Common examples include permanganate and dichromate ions in acidic solution.

The formal charge an atom would carry if all its bonds were treated as fully ionic. It is used to identify oxidation and reduction and to balance redox equations.

A reaction involving simultaneous oxidation (loss of electrons, increase in oxidation number) and reduction (gain of electrons, decrease in oxidation number). Electron transfer links the two half-reactions.

Properties of dilute solutions that depend on the number of solute particles rather than their nature: relative lowering of vapour pressure, elevation of boiling point, depression of freezing point and osmotic pressure.

The ratio of the observed colligative effect to that expected for a non-dissociating solute, reflecting the number of particles a formula unit produces in solution. It is greater than one for electrolytes.

The excess pressure that must be applied to a solution to prevent net flow of solvent across a semipermeable membrane. It equals iCRT for dilute solutions and is useful for molar mass determination.

The potential of a half-cell measured against the standard hydrogen electrode under standard conditions (1 M, 1 bar, 298 K). Comparing values predicts cell emf and the direction of redox reactions.

An equation relating electrode potential to standard potential and the reaction quotient: E = E° − (RT/nF) ln Q. It shows how concentration changes shift cell potential.

A device that converts chemical energy of a spontaneous redox reaction into electrical energy, consisting of two half-cells joined by a salt bridge. Oxidation occurs at the anode, reduction at the cathode.

An experimentally determined equation expressing reaction rate as a function of reactant concentrations, each raised to a power called the order with respect to that reactant. Orders need not match stoichiometric coefficients.

The time required for the concentration of a reactant to fall to half its initial value. For a first-order reaction it is constant and equals 0.693/k, independent of initial concentration.

The minimum extra energy that colliding reactant molecules must possess to form products. It equals the difference between transition-state energy and reactant energy and is lowered by a catalyst.

The steady decrease in atomic and ionic radii across the lanthanoid series due to poor shielding of the 4f electrons. It causes close similarity in size of elements after the lanthanoids.

An element whose atom or common ions have an incomplete d subshell, typically the d-block metals. They show variable oxidation states, coloured ions, catalytic activity and complex formation.

An ion or molecule that donates one or more electron pairs to a central metal ion to form a coordinate bond in a complex. Ligands are classified by the number of donor atoms they use.

The number of donor atoms directly bonded to the central metal ion in a coordination entity. It determines the geometry, commonly octahedral (6) or tetrahedral/square planar (4).

A model explaining bonding in complexes by electrostatic interaction between the metal ion and ligands, which splits d-orbital energies. The splitting pattern accounts for colour and magnetic properties.

Two nucleophilic substitution pathways: SN1 proceeds via a carbocation intermediate in two steps with first-order kinetics, while SN2 is a one-step backside attack with second-order kinetics and inversion of configuration.

In addition of HX to an unsymmetrical alkene, the hydrogen adds to the carbon with more hydrogens, giving the more stable carbocation intermediate. Peroxides can reverse this to the anti-Markovnikov product.

An electron-rich species that donates an electron pair to an electron-deficient centre, such as a carbocation or carbonyl carbon, forming a new bond. It is attracted to electrophiles.

An electron-deficient species that accepts an electron pair from a nucleophile to form a bond. Common examples are carbocations and the nitronium ion in aromatic substitution.

Base- or acid-catalysed reaction in which an enolate from one carbonyl compound attacks another carbonyl carbon, giving a β-hydroxy aldehyde or ketone that may dehydrate to an α,β-unsaturated product.

Compounds such as acid chlorides, esters, amides and anhydrides formed by replacing the hydroxyl group of a carboxylic acid. They undergo nucleophilic acyl substitution, with reactivity decreasing as the leaving group worsens.

A carbohydrate that can reduce Tollens' or Fehling's reagent because it has a free aldehyde or ketone group, or a hemiacetal that can open to one. Sucrose is non-reducing.

Cyclic stereoisomers of a sugar that differ only in the configuration at the anomeric carbon, the hemiacetal carbon formed on ring closure. They are designated α or β.

An amide linkage formed between the carboxyl group of one amino acid and the amino group of another with loss of water. Its partial double-bond character makes it planar and rigid.

A molecule carrying both positive and negative charges simultaneously, as amino acids do in neutral solution. The pH at which the net charge is zero is the isoelectric point.

A chain reaction in which unsaturated monomers join without loss of any small molecule, giving a polymer whose repeating unit has the same atoms as the monomer. Polyethylene and PVC form this way.

Stepwise polymerisation of bifunctional monomers with elimination of a small molecule such as water or HCl. Nylon-6,6 and polyesters are typical products.

A polymer formed from two or more different monomers, such as styrene-butadiene rubber or the nylon salt of hexamethylenediamine and adipic acid. Monomer ratio controls its properties.

A substance that neutralises excess stomach hydrochloric acid, often a weak base or buffer system such as magnesium hydroxide or sodium bicarbonate. It relieves hyperacidity without greatly altering pH.

A volumetric procedure in which a solution of known concentration is added to a measured volume of analyte until the equivalence point is signalled, usually by an indicator or instrument. It determines unknown concentration.

A highly pure, stable, non-hygroscopic substance used to prepare a solution of accurately known concentration. Examples include oxalic acid and potassium hydrogen phthalate.

The observable change, such as a colour change of indicator, that signals completion of a titration. It should be as close as possible to the stoichiometric equivalence point.

A systematic scheme of adding reagents to separate and identify cations as characteristic precipitates, often grouped by behaviour with HCl, H2S and ammonium sulphide. Confirmatory tests follow group separation.