Acids and Bases Common Types of Acids & Bases
Many common acids have a proton attached to an O atom: Carboxylic acids Many bases have a nitrogen atom with a lone pair Arrhenius: H+(aq) + OH-(aq) H2O (l) Acid = H+ produced in aqueous solution e.g. HCl Base = OH- producer e.g. NaOH Bronsted – Lowry: H+ + A- HA Acid = proton donor (H+) e.g. HCl Base = proton acceptor e.g. NH3
Conjugate Acid-Base Pairs
A conjugate base has one less proton than its conjugate acid HSO4 Conjugate bas is SO42 Conjugate acid is H2SO4
Strong acids and bases completely ionise in water: E.g. HCl (aq) + H20 (l) H3O+ (aq) + Cl- (aq) Strong acids: H2SO4, HCl, HBr, HI, HNO3, HClO4 Strong bases: All hydroxides of Groups 1 and 2 (except Be): NaOH, Ca(OH)2 Equilibrium lies completely to the right Examples What is the pH of a 0.1M HCl solution? HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq)
[H3O+] = 0.1M pH = -log10[H3O+] = 1.0 What is the pH of a 0.002M NaOH solution? Completely ionised, so [OH-] = 0.002M pOH = -log10[OH-] = -log10[0.002)] = 2.7 pH = 14 – 2.7 = 11.3
Weak Acids and Equilibrium
HA(aq) + H2O(l) H3O+(aq) + A-(aq) A Strong acid has equilibrium to the right (HA completely ionised) A Weak acid has equilibrium to the right (HA partly/most intact) Equilibrium Equation: Ka (Acid Ionisation Constant)= [H3O+][A-] [HA] pKa = -log10Ka The larger the value of Ka the stronger the acid and the lower the value of pKa The concentration of water [H2O] is assumed to be constant Most acids or bases are weak – they do not completely ionise in water Example Find the pH of 0.1M accetic acid (CH3COOH (HAc)), when pKa = 4.7, Ka = 10-4.7 Conc Initial Change Equil.
HAc(aq) + H2O(l) H3O+(aq) + Ac-(aq) 0.1 large 0 0 -x +x +x 0.1 – x x x
Ka = 10-4.7 = x2/0.1 – x (Assume that the equilibrium constant is very small so that 0.1 – x = 0.1) 10-4.7 = x2/0.1 x = √(10-4.7 x 0.1) = 10-2.85 pH = -log10[H3O+] = -log10[10-2.85] = 2.9 #
Autoionisation of Water H2O(l) + H2O(l) H3O+(aq) + OH-(aq)
Equilibrium constant (Kw) = [H3O+][OH-] = 1.0 x 10-14 pKw = -log10Kw = 14.00 0 At 25 C: Kw = 1.0 x 10-14 reaction is endothermic: more favourable at higher temperature Neutral solution: [H3O+] = [OH-] = 1.0 x 10-7Acidic solution: [H3O+] > 1.0 x 10-7 Basic solution: [H3O+] < 1.0 x 10-7
Temperature Dependence of pH
Kw = 1.0 x 10-14 only at 25oC Reaction is endothermic: more favourable at higher temperature For T > 25oC, Kw > 10-14 For T < 25oC, Kw < 10-14 If T was not 25oC, then pH + pOH does not = 14 and neutral pH is not 7