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5.1.3 Acids, bases and buffers

Definitions

Term Definition
Monobasic acid 1 hydrogen ion can be replaced per molecule in acid-base reaction
Dibasic acid 2 hydrogen ions can be replaced per molecule in acid-base
Tribasic acid 3 hydrogen ions can be replaced per molecule in acid-base reaction
End point Where the indicator changes colour

Brønsted–Lowry acids and bases

Brønsted-Lowry acids and bases

  • Brønsted-Lowry acid = a proton donor
  • Brønsted-Lowry base = a proton acceptor

Conjugate acid-base pairs

  • Two species that can be interconverted by transfer of a proton (cannot be more than 1 protons)
  • \(\text{Conjugate acid (aq)} \rightleftharpoons H^+(aq) + \text{Conjugate base (aq)}\)

pH and [H+(aq)]

pH formula

  • \(pH = -\log [H^+]\)
  • \([H^+] = 10^{-pH}\)

Acid dissociation constant \(K_a\)

  • For reaction \(HA(aq) \rightleftharpoons H^+(aq) + A^-(aq)\)
    • \(K_a = \frac{[H^+(aq)][A^-(aq)]}{[HA(aq)]}\)
  • Unit = \(mol \cdot dm^{-3}\)
  • Changes with temperature, recorded values normally standardised at 25°C

\(pK_a\) values

  • \(pK_a = -\log K_a\)
  • \(K_a = 10^{-pK_a}\)
  • Used when \(K_a\) is small
  • Stronger acid = higher \(K_a\) + lower \(pK_a\)

pH calculations

  • pH for strong monobasic acid
    • Assume \([H^+(aq)] = [HA(aq)]\) (= acid concentration)
  • pH for weak monobasic acid
    • \(K_a \approx \frac{[H^+(aq)]^2}{[HA(aq)]}\)
    • Precise formula: \(K_a = \frac{[H^+(aq)]_{eqm}[A^-(aq)]_{eqm}}{[HA(aq)]_{start} - [H^+(aq)]_{eqm}}\)
    • Assumptions:
      1. Equal concentration of \(H^+\) and \(A^-\): \([H^+(aq)]_{eqm} \approx [A^-(aq)]_{eqm}\) (Not valid for very weak acids / very dilute solutions)
      2. Dissociation is negligible: \([HA(aq)]_{eqm} \approx [HA(aq)]_{start}\) (Not valid for stronger weak acids with \(K_a > 10^{-2} \ mol \cdot dm^{-3}\) / very dilute solutions)

Ionic product of water \(K_w\)

  • How water ionises: \(H_2O(l) \rightleftharpoons H^+(aq) + OH^-(aq)\)
  • \(K_w = [H^+(aq)][OH^-(aq)]\)
  • \(K_w\) at 298K = \(1.00 \times 10^{-14} \ mol^2 \cdot dm^{-6}\) (varies with temperature, can assume it is this value unless stated otherwise)

Buffers: action, uses and calculations

Buffer solution

  • A system that minimises pH changes on addition of small amounts of acid or alkali
  • Contains a weak acid + its conjugate base in a salt

Preparing buffer solutions

  • Mixing weak acid + one of its salts
    • e.g. \(CH_3COOH + CH_3COONa\)
    • Weak acid partially dissociates when added to water \(\rightarrow\) provide the weak acid component
    • Salts of weak acids = ionic compounds \(\rightarrow\) completely dissolve + dissociate into ions in water \(\rightarrow\) provide the conjugate base
  • Adding a strong alkali to an excess of a weak acid
    • e.g. excess \(CH_3COOH + NaOH\)
    • Weak acid partially neutralised by the alkali \(\rightarrow\) conjugate base formed
    • Some weak acid left unreacted
    • Result: mixture of salt of weak acid + unreacted weak acid

How buffer solutions work

  • Equilibrium established e.g. \(HA(aq) \rightleftharpoons H^+(aq) + A^-(aq)\)
  • Acid / \(H^+(aq)\) added
    • \([H^+(aq)]\) increases
    • \(H^+(aq)\) ions react with the conjugate base (\(A^-\))
    • Equilibrium position shifts to the left
    • Most of the \(H^+(aq)\) ions added is used up so pH change is minimised
  • Alkali / \(OH^-(aq)\) added
    • \([OH^-(aq)]\) increases
    • \(OH^-(aq)\) reacts with \(H^+\) to form water so \(H^+\) is used up
    • The equilibrium moves to RHS to replace most of the \(H^+\) used up and minimise the pH change
    • HA dissociates \(\rightarrow\) shifts the equilibrium to the right \(\rightarrow\) restore most of the \(H^+(aq)\) ions
    • Overall reaction: \(HA(aq) + OH^-(aq) \rightarrow H_2O(l) + A^-(aq)\)

Calculating pH of a buffer solution

  • Acid + alkali: calculate how much \(A^-\) is in the solution and how much \(HA\) is left
  • Acid + salt: assume all \(A^-\) from salt
  • Assume \([HA]\) stays constant in both cases
  • \([H^+(aq)] = K_a \times \frac{[HA(aq)]}{[A^-(aq)]}\)
  • Quick formula: \(pH = pK_a + \log \frac{[\text{salt}]}{[\text{acid}]}\)

Control of blood pH

  • Blood plasma needs to have a pH between 7.35 to 7.45
  • pH below 7.35 \(\rightarrow\) acidosis (fatigue, shortness breath, shock, death)
  • pH above 7.45 \(\rightarrow\) alkalosis (muscle spasms, light headedness, nausea)
  • pH maintained by carbonic acid (\(H_2CO_3\)) - hydrogencarbonate (\(HCO_3^-\)) buffer system: \(H_2CO_3 \rightleftharpoons H^+ + HCO_3^-\)

Neutralisation

pH titration curve

  • Weak base - steeper gradient for basic section
  • Weak acid - steeper gradient for acidic section
  • Exported image
  • (Acid to base: mirror the shapes)

Equivalence point

  • Halfway up the vertical section
  • Where the amount of acid exactly neutralises the amount of alkali

Indicator colour change explanation

  • Indicator = weak acid with distinctively different colour to conjugate base
  • \(HIn(aq) \rightleftharpoons H^+(aq) + In^-(aq)\)
  • Equilibrium position shifts toward LHS in acidic conditions
  • Equilibrium position shifts toward RHS in basic conditions

Common indicators

  • Exported image
  • (Most indicators change colour over a range of 2 pH units)

Choice of suitable indicators for titration

  • Choose indicator with a colour change range that coincides with the vertical section of the pH titration curve
  • Ideally same end point and equivalence point
  • End point not in vertical section: range of volumes for colour change is too big
  • Titration Curves ALevel ChemistryStudent
  • No indicator suitable for weak acid-weak base titrations as there is no vertical section + even at steepest requires several \(cm^3\) to pass through a range of 2 pH units