Although fluorine is smaller, its 2p subshell is very compact, so the incoming electron experiences strong inter-electronic repulsion from the already crowded 2p electrons. In chlorine, the larger 3p subshell accommodates the incoming electron with less repulsion, so more energy is released. Hence chlorine has a more negative (more exothermic) electron gain enthalpy than fluorine.

Electronegativity is the tendency of an atom in a molecule to attract the shared pair of electrons in a covalent bond towards itself. It increases across a period (as size decreases and nuclear charge increases) and decreases down a group (as size increases). Fluorine is the most electronegative element. The Pauling scale is commonly used.

All four species have 10 electrons (isoelectronic). For isoelectronic species, the greater the nuclear charge (number of protons), the smaller the radius. Nuclear charges: O2- (8), F- (9), Na+ (11), Mg2+ (12). Increasing order of radius: Mg2+ < Na+ < F- < O2-.

An ionic (electrovalent) bond is formed by the complete transfer of one or more electrons from one atom to another, producing oppositely charged ions held by electrostatic attraction (e.g. NaCl). A covalent bond is formed by the mutual sharing of electron pairs between atoms (e.g. H2 or Cl2). Ionic bonds form between metals and non-metals; covalent bonds form between non-metals.

Both molecules are based on sp3 hybridisation with an ideal angle of 109.5 deg. NH3 has one lone pair, while H2O has two lone pairs. Lone pair-lone pair repulsion is stronger than lone pair-bond pair repulsion, and water has an extra lone pair, so its bond pairs are pushed closer together. Hence the bond angle of H2O (104.5 deg) is smaller than that of NH3 (107 deg).

N2 has 14 electrons. Number of bonding electrons Nb = 10 and antibonding electrons Na = 4. Bond order = (Nb - Na)/2 = (10 - 4)/2 = 3. This corresponds to a nitrogen-nitrogen triple bond, which explains the very high stability and bond dissociation energy of N2.

Bond order = (Nb - Na)/2

A coordinate bond is a covalent bond in which the shared pair of electrons is contributed by only one of the two bonded atoms (the donor). In the ammonium ion (NH4+), nitrogen in NH3 has a lone pair which it donates to a proton (H+) that has no electrons, forming the fourth N-H bond as a coordinate bond. Once formed, all four N-H bonds are identical.

NH3 + H+ -> NH4+

A hydrogen bond is a weak electrostatic attraction between a hydrogen atom covalently bonded to a highly electronegative atom (F, O or N) and the lone pair of another electronegative atom. Conditions: hydrogen must be attached to a small, highly electronegative atom. Types: intermolecular hydrogen bonding (between different molecules, e.g. in water and HF) and intramolecular hydrogen bonding (within the same molecule, e.g. o-nitrophenol).

Oxygen is much more electronegative and smaller than sulphur, so water molecules form strong intermolecular hydrogen bonds, whereas H2S molecules are held only by weak van der Waals (dipole) forces. Extra energy is needed to break the hydrogen bonds in water, so water has a much higher boiling point than H2S even though H2S has a higher molar mass.

A sigma bond is formed by the head-on (axial) overlap of orbitals along the internuclear axis, while a pi bond is formed by the sidewise (lateral) overlap of parallel p orbitals. A sigma bond is stronger because axial overlap is more effective and greater, giving a larger region of electron density between the nuclei; pi bonds have smaller lateral overlap and are weaker and more reactive.

Boyle's law: at constant temperature, the volume of a fixed mass of gas is inversely proportional to its pressure, i.e. V is proportional to 1/P, so PV = constant. Charles's law: at constant pressure, the volume of a fixed mass of gas is directly proportional to its absolute (Kelvin) temperature, i.e. V is proportional to T, so V/T = constant.

PV = constant (Boyle); V/T = constant (Charles)

The ideal gas equation is PV = nRT, where P = pressure, V = volume, n = number of moles, T = absolute temperature and R = universal gas constant. R = 0.0821 L atm K^-1 mol^-1 = 8.314 J K^-1 mol^-1.

PV = nRT

The van der Waals equation is (P + a n^2/V^2)(V - nb) = nRT. The constant 'a' corrects for the intermolecular forces of attraction (it accounts for the pressure being lower than ideal), and the constant 'b' corrects for the finite volume actually occupied by the gas molecules (excluded volume). Real gases deviate from ideal behaviour at high pressure and low temperature.

(P + a n^2/V^2)(V - nb) = nRT

A system is the specified part of the universe under study; the surroundings are the rest of the universe outside the system that can interact with it. The three types are: open system (exchanges both matter and energy with surroundings), closed system (exchanges only energy, not matter) and isolated system (exchanges neither matter nor energy).

The first law of thermodynamics states that energy can neither be created nor destroyed, only transformed from one form to another; the total energy of an isolated system remains constant. Mathematically, delta U = q + w, where delta U is the change in internal energy, q is the heat added to the system (positive when absorbed) and w is the work done on the system (positive when done on the system).

delta U = q + w

Enthalpy (H) is the total heat content of a system at constant pressure, defined as H = U + PV. For a reaction at constant pressure and temperature, delta H = delta U + P delta V. For ideal gases, P delta V = delta ng RT, where delta ng is the change in the number of moles of gaseous species. Hence delta H = delta U + delta ng RT.

delta H = delta U + delta ng RT

Hess's law states that the total enthalpy change of a reaction is the same whether the reaction takes place in one step or in several steps, provided the initial and final conditions are the same. It follows from the fact that enthalpy is a state function. Applications: it is used to calculate enthalpies of formation, bond enthalpies and reaction enthalpies that cannot be measured directly.

Entropy (S) is a thermodynamic state function that measures the degree of randomness or disorder of a system. When ice (a highly ordered solid) melts into water (a more disordered liquid), disorder increases, so the entropy increases and delta S is positive.

delta S = q(rev) / T

Standard enthalpy of formation is the enthalpy change when one mole of a compound is formed from its constituent elements in their standard states (at 298 K and 1 bar); e.g. for CO2 it is -393.5 kJ/mol. Standard enthalpy of combustion is the enthalpy change when one mole of a substance is completely burnt in excess oxygen under standard conditions; it is always negative (exothermic).

A reversible reaction is one that proceeds in both forward and backward directions under the same conditions. Characteristics of chemical equilibrium: (1) it is dynamic in nature, i.e. the forward and backward reactions continue at equal rates; (2) the observable properties (concentration, pressure, colour) remain constant with time; (3) it can be attained from either direction and is disturbed by changing conditions.

rate(forward) = rate(backward)