GTLE Subject Specialism Chemistry flashcards

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

Effective nuclear charge

The net positive charge experienced by an electron after subtracting shielding by other electrons. Across a period Z_eff rises, contracting the atom and raising ionisation energy.

Hund's rule

Within a subshell, electrons occupy separate orbitals with parallel spins before pairing. This maximises exchange energy and explains the paramagnetism of atoms with partly filled p or d subshells.

Successive ionisation energy trend

Each successive ionisation energy is larger because removing an electron from an increasingly positive ion requires more energy. A large jump reveals removal from a shell closer to the nucleus.

VSEPR lone pair repulsion

Lone pairs occupy more angular space than bonding pairs, so they compress bond angles. Methane is 109.5 degrees, ammonia 107, and water 104.5 in a tetrahedral electron-domain arrangement.

Formal charge rule

Formal charge equals valence electrons minus nonbonding electrons minus half the bonding electrons. The preferred resonance structure minimises formal charges and puts negative charge on the more electronegative atom.

Hydrogen bonding

An unusually strong dipole attraction between H bonded to N, O or F and a lone pair on a nearby N, O or F. It raises boiling points and makes ice less dense than water.

Limiting reagent

The reactant that is completely consumed first, determined by dividing each reactant's moles by its stoichiometric coefficient. The smallest quotient limits the theoretical yield of product.

Molar volume at STP

One mole of an ideal gas occupies 22.7 dm³ at 273 K and 100 kPa, or 22.4 dm³ at 273 K and 1 atm. Use n = V/V_m for gas stoichiometry.

Percentage yield

Actual mass of product divided by theoretical mass predicted from the limiting reagent, multiplied by 100. Atom economy instead compares product molar mass with total reactant molar mass.

Graham's law

The rate of effusion or diffusion of a gas is inversely proportional to the square root of its molar mass. Lighter gases such as helium escape faster than heavier ones at the same temperature.

van der Waals equation

A real-gas correction: (P + an²/V²)(V − nb) = nRT. The a term corrects for attractive forces, and b corrects for the finite volume occupied by molecules.

Critical temperature

The highest temperature at which a substance can be liquefied by pressure alone. Above it, the gas and liquid phases are indistinguishable and no amount of pressure condenses the gas.

Hess's law

The enthalpy change of a reaction is independent of the route taken. It allows ΔH to be calculated from standard enthalpies of formation or from bond enthalpy sums.

Standard enthalpy of formation

The enthalpy change when one mole of a compound forms from its elements in their standard states at 298 K and 100 kPa. Elements in standard states are assigned zero.

Gibbs energy equation

ΔG equals ΔH minus TΔS, with temperature in kelvin. Negative ΔG indicates a spontaneous process; the sign of ΔH and ΔS together determines the temperature dependence.

Le Chatelier's principle

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

Reaction quotient Q

The same expression as the equilibrium constant but using instantaneous concentrations or pressures. If Q is less than K the forward reaction is favoured; if greater, the reverse.

First-order half-life

For a first-order reaction, half-life equals 0.693 divided by the rate constant and is independent of initial concentration. Concentration decays exponentially with time.

Henderson–Hasselbalch equation

pH equals pKa plus the logarithm of the conjugate base concentration divided by the weak acid concentration. It describes buffer pH and is most reliable near the pKa.

Salt hydrolysis

Reaction of dissolved ions with water that alters pH. Salts of weak acids with strong bases give basic solutions; salts of weak bases with strong acids give acidic solutions.

Buffer capacity

The amount of added acid or base a buffer can absorb before pH changes appreciably. It increases with the absolute concentrations of the conjugate acid-base pair, not just their ratio.

Nernst equation

E equals E° minus (RT/nF) ln Q, often written with 0.0592/n at 25 °C. It gives cell potential at non-standard concentrations and links potential to equilibrium.

Standard hydrogen electrode

The reference electrode with H⁺ at 1 mol dm⁻³ and hydrogen gas at 1 bar over platinum, assigned exactly 0.00 V. All standard reduction potentials are measured relative to it.

Faraday's electrolysis law

The moles of substance produced or consumed equal charge divided by (n times F), where n is electrons per formula unit and F is 96,485 C per mole.

Markovnikov's rule

In addition of HX to an alkene, hydrogen adds to the carbon with more hydrogens, forming the more stable carbocation. Peroxide conditions reverse this regiochemistry via radical intermediates.

SN2 mechanism

A one-step bimolecular substitution in which the nucleophile attacks from the side opposite the leaving group. Rate depends on both substrate and nucleophile, and the product is inverted.

SN1 mechanism

A two-step unimolecular substitution via a carbocation intermediate, favoured by tertiary and benzylic halides. Rate depends only on substrate, and stereochemistry is largely racemised.

Electrophilic aromatic substitution

Aromatic rings attack electrophiles, temporarily breaking aromaticity before deprotonation restores it. Activating ortho-para directors have lone pairs; deactivating meta directors are electron-withdrawing.

Chiral centre

A carbon atom bonded to four different groups, giving non-superimposable mirror images. A racemic mixture contains equal amounts of enantiomers and shows no net optical rotation.

Carboxylic acid acidity

Carboxylic acids are stronger than alcohols because the conjugate base is stabilised by resonance over two oxygen atoms. Electron-withdrawing substituents such as chlorine further increase acidity.

Ligand field splitting

The energy gap Δ between d orbitals when ligands approach a metal ion. Strong-field ligands such as cyanide cause large splitting and low-spin configurations; weak-field ligands favour high-spin.

Chelate effect

Multidentate ligands such as EDTA bind a metal ion at several sites, giving complexes far more stable than comparable monodentate ligands, largely through favourable entropy of displacement.