PAU Chemistry (Cantabria) flashcards

32 cards. Read the term and the definition here. The deck at the top of this page is optional.

Bohr energy levels

Electrons occupy quantized orbits; energy is absorbed or emitted only when an electron moves between levels, giving line spectra. ΔE = hν and 1/λ = R(1/n₁² − 1/n₂²).

Heisenberg uncertainty principle

Position and momentum of an electron cannot both be known precisely at the same time, so Bohr orbits are replaced by probability distributions (orbitals) in the quantum mechanical model.

Hund's rule

When filling degenerate orbitals, electrons occupy them singly with parallel spins before pairing. This maximizes total spin and explains paramagnetism of atoms with unpaired electrons.

First ionization energy

Energy required to remove one mole of electrons from one mole of gaseous atoms. It increases across a period and decreases down a group as atomic radius grows.

VSEPR theory

Electron pairs around a central atom repel each other and adopt the arrangement that minimizes repulsion. Lone pairs occupy more space than bonding pairs, compressing bond angles.

Hydrogen bond

A strong dipole–dipole attraction between H bonded to N, O, or F and a lone pair on a nearby N, O, or F atom. It explains water's high boiling point and ice's low density.

London dispersion forces

Weak, temporary attractions caused by instantaneous dipoles in all molecules. They grow with the number of electrons and molecular surface area, so heavier halogens are more polarizable.

Lattice energy

Enthalpy released when one mole of an ionic solid forms from gaseous ions. It increases with higher ionic charges and smaller ionic radii, raising melting point.

Hess's law

The enthalpy change of a reaction is independent of the route taken. It allows ΔH to be calculated by summing known reactions, reversing signs when a step is reversed.

Standard enthalpy of formation

Enthalpy change when one mole of a compound forms from its elements in their standard states. It is zero for elements and used in ΔH°rxn = ΣΔHf°(products) − ΣΔHf°(reactants).

Entropy

A measure of the dispersal of energy or disorder of a system. It increases with temperature, phase changes to gas, more gas moles, and greater molecular complexity.

Gibbs free energy

ΔG = ΔH − TΔS determines spontaneity. A negative ΔG means the process is spontaneous at that temperature; ΔG = 0 corresponds to equilibrium.

Rate law and order

An equation v = k[A]ᵐ[B]ⁿ relating rate to concentrations. Exponents are determined experimentally, not from stoichiometric coefficients, and the overall order is their sum.

Half-life

Time for the concentration of a reactant to fall to half its initial value. For first-order reactions t½ = ln2/k, independent of initial concentration.

Arrhenius equation

k = A·e^(−Ea/RT) shows that the rate constant increases with temperature and decreases with activation energy. A catalyst lowers Ea without changing ΔH or K.

Le Chatelier's principle

When a system at equilibrium is disturbed by changes in concentration, pressure, or temperature, it shifts in the direction that partially counteracts the change. Catalysts do not shift it.

Strong versus weak acids

Strong acids dissociate completely in water, giving pH from initial concentration. Weak acids dissociate partially, described by Ka, with [H₃O⁺] ≈ √(Ka·C₀) for dilute solutions.

Buffer solution

A mixture of a weak acid and its conjugate base (or weak base and conjugate acid) that resists pH change. Henderson–Hasselbalch gives pH = pKa + log([base]/[acid]).

Salt hydrolysis

Reaction of a dissolved salt's ions with water to change pH. Salts of weak acids give basic solutions; salts of weak bases give acidic solutions; strong–strong salts are neutral.

Equivalence point in titration

The point where the moles of acid equal the moles of base added. Its pH is not necessarily 7; it depends on the strength of the acid and base involved.

Solubility product Ksp

The equilibrium constant for dissolution of a sparingly soluble solid, with the solid omitted. For MX, Ksp = s²; for MX₂ or M₂X, Ksp = 4s³.

Common-ion effect

Adding an ion already present in the dissolution equilibrium shifts it left, decreasing the salt's solubility through Le Chatelier's principle. Ksp itself stays constant.

Ion product Q

The product of ion concentrations raised to their stoichiometric powers at a given instant. If Q > Ksp, a precipitate forms; if Q < Ksp, the solution remains unsaturated.

Oxidation number

A bookkeeping charge assigned to an atom in a compound. Oxidation is an increase in oxidation number, reduction is a decrease, and the sum equals the overall charge.

Standard cell potential

E°cell = E°cathode − E°anode using standard reduction potentials. A positive E°cell indicates a spontaneous galvanic reaction linked to ΔG° = −nFE°.

Nernst equation

E = E° − (0.0592/n)log Q at 25 °C, correcting cell potential for nonstandard concentrations. It shows that potential falls as products accumulate.

Faraday's laws of electrolysis

The amount of substance deposited or liberated is proportional to the charge passed. m = (M·I·t)/(n·F), with F = 96485 C/mol and n the electrons per ion.

Markovnikov's rule

In electrophilic addition of HX to an alkene, the hydrogen adds to the carbon bearing more hydrogens, placing the halogen on the more substituted carbon and stabilizing the carbocation.

Optical isomerism

Stereoisomerism in molecules with a chiral carbon bearing four different groups. The two enantiomers are non-superimposable mirror images and rotate plane-polarized light oppositely.

Functional group priority

When naming, the principal group with highest priority gives the suffix: carboxylic acid > ester > aldehyde > ketone > alcohol > amine > alkene > alkane, with lowest locants.

Addition polymerization

Monomer alkenes join through their double bonds without losing atoms, forming long chains such as polyethylene from ethene. It contrasts with condensation polymerization, which eliminates small molecules.

Electrophilic aromatic substitution

A reaction in which an electrophile replaces a hydrogen on a benzene ring, as in nitration with HNO₃/H₂SO₄. The ring's delocalized π electrons are the electron-rich site attacked.