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)

Le Chatelier's principle states that if a system at equilibrium is subjected to a change in concentration, pressure or temperature, the equilibrium shifts in the direction that tends to counteract (reduce) the effect of that change. Increasing the pressure shifts this equilibrium in the forward direction (towards NH3), because the forward reaction reduces the number of gas moles from 4 (1 + 3) to 2, thereby lowering the pressure.

According to the Bronsted-Lowry concept, an acid is a substance that donates a proton (H+) and a base is a substance that accepts a proton. A conjugate acid-base pair is a pair of species that differ by a single proton; for example, in HCl + H2O -> H3O+ + Cl-, HCl/Cl- and H2O/H3O+ are conjugate acid-base pairs.

acid <=> base + H+

HCl is a strong acid and dissociates completely, so [H+] = 0.001 M = 1 x 10^-3 M. pH = -log[H+] = -log(10^-3) = 3. The solution is acidic (pH < 7).

pH = -log[H+]

The ionic product of water Kw is the product of the molar concentrations of hydrogen and hydroxide ions in water: Kw = [H+][OH-]. At 298 K, Kw = 1.0 x 10^-14 mol^2 L^-2. Taking negative logarithm gives pKw = pH + pOH = 14 at 298 K. For pure (neutral) water, [H+] = [OH-] = 10^-7 M, so pH = 7.

Kw = [H+][OH-] = 1.0 x 10^-14

The common ion effect is the suppression of the degree of dissociation (ionisation) of a weak electrolyte by the addition of a strong electrolyte that provides an ion common to the weak electrolyte. For example, adding NH4Cl (which provides NH4+) to a solution of NH4OH suppresses the ionisation of NH4OH, decreasing the OH- concentration. This is used in qualitative analysis and to control pH.

Oxidation is the loss of electrons or an increase in oxidation number of an element; reduction is the gain of electrons or a decrease in oxidation number. Both occur simultaneously in a redox reaction. For example, in Zn -> Zn2+ + 2e-, zinc is oxidised (oxidation number 0 -> +2).

In KMnO4: K = +1, O = -2 (four O = -8). Let Mn = x. Then +1 + x + (-8) = 0, so x = +7. In K2Cr2O7: 2 K = +2, 7 O = -14. Let each Cr = y. Then +2 + 2y - 14 = 0, so 2y = 12, y = +6. Thus Mn is +7 and Cr is +6.

sum of oxidation numbers = charge on species

An oxidising agent is a substance that oxidises another substance by accepting electrons and is itself reduced; e.g. KMnO4, O2. A reducing agent is a substance that reduces another substance by donating electrons and is itself oxidised; e.g. H2, C (carbon). In a redox reaction the oxidising agent gains electrons and the reducing agent loses electrons.