Physical chemistry

• Dynamic equilibrium


  Occurs in a closed system when the rate of reverse reaction equals rate of forward reaction, so the composition of reactants and products appears to stay constant
• Homogenous equilibrium


  All species in the same phase
• Heterogenous equilibrium


  Species in more than one phase
• Equilibrium constant (K)


  Expressed in terms of concentration (or pressure for gases) for the reaction: aA + bB ⇌ cC + dD, K = [C]^c[D]^d / [A]^a[B]^b
• Concentrations of pure solids or liquids are constant and given the value 1 in the equilibrium constant expression
• Equilibrium constant is for a given temperature and is independent of concentration, pressure or the presence of a catalyst
• When dilute hydrochloric acid is added to a yellow solution of sodium chromate

  The solution turns orange as the equilibrium shifts to the right, increasing the concentration of Cr2O7^2- and decreasing the concentration of CrO4^2-
• Le Chatelier's Principle


  When a reaction at equilibrium is subjected to change, the composition alters to minimise the effects of the change
• For an endothermic reaction (ΔH > 0)


  Increasing temperature causes an increase in the equilibrium constant K
• For an exothermic reaction (ΔH < 0)

  Increasing temperature causes a decrease in the equilibrium constant K
• High value of K


  Means equilibrium lies to the right, with more product
• A catalyst has no effect on the equilibrium constant K, it only speeds up the rate of establishment of equilibrium
• Equilibrium constant (K)


  For the reaction: 2SO2(g) + O2(g) ⇌ 2SO3(g), K = [SO3]^2 / ([SO2]^2[O2])
• Le Chatelier's Principle


  When a system at equilibrium is subjected to a change, the system will shift to counteract the change and re-establish equilibrium
• Increasing temperature for the reaction 2SO2(g) + O2(g) ⇌ 2SO3(g)


  Shifts the equilibrium to the left, decreasing [SO3] and increasing [SO2] and [O2]
• Increasing temperature for the reaction 2SO2(g) + O2(g) ⇌ 2SO3(g)


  Decreases the value of the equilibrium constant K, since the reaction is exothermic
• pH


  A measure of the concentration of hydrogen ions (H+) in a solution, pH = -log[H+]
• For strong acids and bases, the small number of H+ or OH- ions from water can be ignored in pH calculations
• Brønsted-Lowry definition of acids and bases

  An acid is a substance capable of donating a proton, a base is a substance capable of accepting a proton
• Ionic product of water (Kw)

  Kw = [H+][OH-] = 1.0 x 10^-14 at 25°C
• As temperature increases


  The value of Kw increases, as the water dissociation reaction is endothermic
• Ionisation of water


  For every molecule which ionises, one H+ and one OH- ion are produced, hence the [H+] in mol l-1 must equal the [OH-] in mol l-1, i.e. the number of H+ and OH- ions in water are equal
• Substitution of [OH-] by [H+] in the equilibrium expression

  1. [H+]2 = 10-14 mol2 l-2

  2. Taking square roots [H+] = 10-7 mol l-1

  3. [OH-] = 10-7 mol l-1

  4. Kw = [H+][OH-] = 1.0 x 10-14
• Calculating pH of alkali solutions

  • 0.01 mol l-1 NaOH [OH-] = 10-2, [H+] = 10-12, pH = 12
  • 0.5 mol l-1 NaOH [H+] = 2 x 10-14 pH = -0.3 + 14 = 13.7
• Kw is always quoted as 1.0 x 10-14 at 25°C since the value varies with temperature
• H2O(l) ⇌ H+(aq) + OH-(aq)


  As the reaction is endothermic, an increase in temperature moves the equilibrium to the right and a decrease moves it to the left
• Calculating pH of solutions

  1. 0.35 mol l-1 HNO3

  2. 0.14 mol l-1 H2SO4 (assume fully ionised)

  3. 0.78 mol l-1 NaOH
• Conjugate acid


  The species formed when a base accepts a proton
• Conjugate base


  The species left when an acid donates a proton
• Water is amphoteric as it can act both as a proton acceptor and a proton donor
• Strong acids dissociate completely in solution
• Strong bases dissociate completely in solution
• Dissociation constant (Ka)


  A measure of the strength of an acid, the equilibrium constant from the acid dissociation equation
• Strong acids have no meaningful Ka as the equilibrium lies completely to the right
• Weak acids dissociate less than 5% in aqueous solution
• Kb


  The equilibrium constant for the base equilibrium, but Ka is generally used to calculate pH for both acids and bases
• pH of a salt solution

  Depends on the strength of the acid and base from which it was formed
• Calculating pH for a weak acid

  1. Using the equilibrium constant (dissociation constant) Ka

  2. [H+] = √(Ka * c)

  3. pH = -log[H+]
• 0.02 mol l-1 benzoic acid C6H5COOH, a monobasic acid, was found to have a pH of 2.94. Calculate the Ka of this weak acid.
• Buffer solution


  A solution in which the pH remains approximately constant when small amounts of acid or base are added or the solution is diluted