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4.4Unit 4. Rates, Equilibria and Further Organic Chemistry

Acid-base Equilibria — IAL Edexcel Linear (Chemistry) Revision Notes & Topical Questions

5 min read 13 parts Lesson 4 of 56 1 quiz

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Curriculum Equation Key معادلة تفاعل تكوين الإستر (الأسترة) - Esterification Reaction EquationFast Exam Solution
R-COOH + R2-OH ⇌ (conc. H₂SO₄) R-COO-R2 + H₂O
Word Equation & Conditions: حمض كربوكسيلي + كحول (في وجود حمض الكبريتيك المركز كعامل حفاز ونازع للماء) ⇌ إستر + ماء
Location & Specifications: الكيمياء العضوية للثانوية العامة والتحصيلي و Edexcel Chemistry.
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Lesson summary

4.4 Acid-base Equilibria — pH, Ka, Kw, titration curves and buffers — IAL EDEXCEL LINEAR (CHEMISTRY) notes: Use Brønsted-Lowry theory, calculate pH for strong and weak acids and bases, interpret titration curves and indicators, and explain and calculate buffers.Brønsted-Lowry aci…

  • Brønsted-Lowry acid
  • Brønsted-Lowry base
  • Conjugate acid–base pair
  • What the symbols mean
  • Strong acid
  • Buffer solution

Part 1 of 11Brønsted-Lowry acids and bases

Brønsted-Lowry acid
  • Brønsted-Lowry acid a species that donates a proton, e.g. HCl → H+ + Cl−.
Brønsted-Lowry base
  • Brønsted-Lowry base a species that accepts a proton, e.g. OH− + H+ → H2O.
Conjugate acid–base pair
  • Conjugate acid–base pair two species related by the transfer of one proton.
  • CH3COOH + H2O ⇌ CH3COO− + H3O+.
  • CH3COOH (acid) and CH3COO− (its conjugate base) form one pair.
  • H2O (base) and H3O+ (its conjugate acid) form the other.
Water can act as an acid or a base (it is amphiprotic).

Part 2 of 11pH

What the symbols mean:
pH = −log[H+], with [H+] in mol dm−3; [H+] = 10−pH.
What the symbols mean: pH has no unit · [H+] = concentration of hydrogen ions (mol dm−3); each 1 pH unit is a ten-fold change in [H+]
The pH scale is logarithmic with base 10, so each unit is a factor of ten in [H+]; pH 5 is ten times more acidic than pH 6.
pH values are usually given to 2 decimal places.
Diluting a solution tenfold raises the pH of a strong acid by 1 unit.

Part 3 of 11Acid strength, Ka and pKa

Strong acid
  • Strong acid dissociates almost completely, so the equilibrium is far to the right.
  • HCl, HNO3, H2SO4.
Weak acid
  • Weak acid dissociates partially, so an equilibrium forms with its position to the left.
  • Most organic acids, HCN, H2S, H2CO3.
Strength depends on the extent of dissociation, not on concentration.
What the symbols mean:
Ka = [H+][A−] ÷ [HA], in mol dm−3.
What the symbols mean: Ka = acid dissociation constant (mol dm−3) · [H+], [A−], [HA] = equilibrium concentrations (mol dm−3); larger Ka = stronger acid
  • Assumes the H+ from water is negligible.
  • A higher Ka means a more dissociated, stronger acid; e.g. ethanoic acid Ka = 1.74 × 10−5 mol dm−3.
pKa = −log Ka; a smaller pKa means a stronger acid, and most weak acids have pKa between 3 and 7 (ethanoic acid 4.76).

Part 4 of 11pH of strong acids

pH of strong acids
A strong acid is fully ionised, so [H+] = [HA] for a monoprotic acid.
H+ from water is negligible compared with that from the acid.
Dibasic acids have two replaceable protons and react 1:2 with bases, e.g. H2SO4 + 2NaOH → Na2SO4 + 2H2O.
  • H2SO4 → HSO4− + H+ is complete.
  • HSO4− ⇌ SO42− + H+ is suppressed by the H+ already present.
  • So [H+] is less than twice the acid concentration; 0.1 mol dm−3 H2SO4 has pH about 0.98, not 0.69.

Part 5 of 11pH of weak acids

  • Neededthe acid concentration and its Ka.
1
Write the Ka expression.
2
Note [H+] = [A−], so Ka = [H+]2 ÷ [HA].
3
Rearrange: [H+] = √(Ka × [HA]).
4
Calculate pH = −log[H+].
pH of weak acids
Two assumptions
  • [A−] = [H+], because water's H+ is negligible.
  • [HA] at equilibrium ≈ [HA] initially, because dissociation is very small.
The reverse calculation finds Ka from a measured pH: Ka = [H+]2 ÷ c.

Part 6 of 11Ionic product of water and strong bases

What the symbols mean:
H2O ⇌ H+ + OH−; Kw = [H+][OH−] = 1.00 × 10−14 mol2 dm−6 at 298 K.
What the symbols mean: Kw = ionic product of water (1.00 × 10−14 mol2 dm−6 at 298 K) · [H+], [OH−] = ion concentrations (mol dm−3)
  • pKw = −log Kw = 14.00 at 298 K.
  • Kw increases with temperature because dissociation is endothermic, so neutral pH falls below 7 when hot.
  • Pure water is neutral at all temperatures because [H+] = [OH−].
Strong bases are fully ionised, so [OH−] = [base] for a base such as NaOH.
  • [H+] = Kw ÷ [OH−], then pH = −log[H+].
  • To find [OH−] from pH, find [H+] = 10−pH, then [OH−] = Kw ÷ [H+].

Part 7 of 11Analysing pH data

1
The relative strengths of acids are compared using the pH of equimolar solutions at the same temperature; the higher the pH, the weaker the acid.
2
For bases, the higher the pH, the stronger the base.
3
pH of salt solutions
  • NaCl and KNO3 (strong acid + strong base) are pH 7.
  • CH3COONa (weak acid + strong base) is alkaline.
  • NH4Cl (strong acid + weak base) is acidic.
  • CH3COONH4 (weak acid + weak base of similar strength) is neutral.
4
Dilution
  • A strong acid rises by 1 pH unit per tenfold dilution.
  • A weak acid rises by about 0.5 per tenfold dilution.
  • Very dilute HCl (1 × 10−8 mol dm−3) has pH close to 7, because water's H+ can no longer be ignored.
5
A pH meter should be calibrated with buffer solutions.

Part 8 of 11Titration curves and indicators

Titration curves and indicators
A pH curve plots pH against the volume of titrant; it is S-shaped and the midpoint of the steep section is the equivalence point.
From the curve you can read:
  • the starting pH of the acid,
  • the pH at the equivalence point,
  • the volume of base at the equivalence point,
  • the pH range of the vertical section.
Acid / baseStarting pHEquivalence pHVertical region
Strong / strongabout 17about 3 to 11
Weak / strongabout 3about 9about 7 to 11
Strong / weakabout 1about 5about 3 to 7
Weak / weakabout 3about 7No sharp vertical region
Titration curves and indicators
An indicator is a weak acid HA(colour 1) ⇌ H+ + A−(colour 2); KIn is its dissociation constant.
  • It changes colour over about pKIn ± 1; methyl orange has pKIn = 3.7.
  • The colour change must fall within the vertical section of the curve.
  • Methyl orange is unsuitable for a weak acid–strong base titration (pH 8–10); phenolphthalein is used.
Weak acid with strong base (NaOH + CH3COOH)
  • The initial rise is steep, then ethanoate ions form a buffer so pH rises gradually.
  • At half-equivalence [HA] = [A−], so pH = pKa.
  • The equivalence point is above pH 7.

Part 9 of 11How a buffer works

Buffer solution
  • Buffer solution a solution that resists changes in pH when small amounts of acid or alkali are added.
It contains a weak acid and its conjugate base (pH below 7), or a weak base and its conjugate acid (pH above 7).
Ethanoic acid and sodium ethanoate
  • The acid is only partly ionised, giving a reserve of CH3COOH.
  • The salt is fully ionised, giving a large reserve of CH3COO−.
Adding H+
  • CH3COO− reacts with H+ to form CH3COOH and the equilibrium shifts left.
  • Both concentrations change little, so the pH stays almost constant.
Adding OH−
  • OH− reacts with H+ to form water and more CH3COOH ionises, shifting the equilibrium right.
  • Both concentrations change little, so the pH stays almost constant.
A mixture of excess weak acid with NaOH gives a buffer because the acid is partly neutralised.

Part 10 of 11Buffer calculations and applications

1
From the Ka expression, [H+] = Ka × [acid] ÷ [base (salt)].
2
Taking logs gives pH = pKa + log([A−] ÷ [HA]) (the Henderson–Hasselbalch equation).
3
To make a buffer of a given pH
  • Find [H+] = 10−pH.
  • Use the equation to find the ratio [acid] ÷ [base] needed.
  • For pH 5.00 with ethanoic acid (Ka = 1.74 × 10−5), the ratio is 0.575.
  • Mix solutions of the two to give that ratio.
4
Blood
  • HCO3− buffers blood at pH 7.35–7.45.
  • CO2 + H2O ⇌ H+ + HCO3−; extra H+ shifts the equilibrium left and a fall in H+ shifts it right.
  • Uncontrolled H+ would cause acidosis.
5
Food
  • Most microorganisms thrive near neutral pH (6.6–7.5) and survive between about pH 4.4 and 9.0.
  • Buffer capacity is the amount of acid or base needed to change the pH significantly.
  • Protein-rich foods have a higher buffer capacity because amino acids are both acidic and basic.

Part 11 of 11More from the specification

Aqueous
  • Dissolved in water.
Hydroxide
  • A compound containing the OH− ion, formed by Group I metals with water.
Ionisation
  • The removal of an electron from an atom or molecule.

Key terms

TermMeaning
Brønsted-Lowry acidA proton donor.
Brønsted-Lowry baseA proton acceptor.
Conjugate acid-base pairTwo species that differ by one proton.
pHpH = −log10[H+(aq)].
KaThe acid dissociation constant, [H+][A-]/[HA].
KwThe ionic product of water, [H+][OH-] = 1.00 × 10-14 mol2 dm-6 at 298 K.
Buffer solutionA solution that resists change in pH when small amounts of acid or alkali are added.
Half-equivalence pointThe point where half the weak acid is neutralised and pH = pKa.

Avoid these mistakes

1
Saying strength and concentration are the same — strength is the extent of dissociation.
2
Forgetting the two assumptions in weak acid pH calculations.
3
Using methyl orange for a weak acid–strong base titration — its range is too low.
4
Saying pure water is only neutral at 25 °C — it stays neutral because [H+] = [OH-], although the pH changes.
5
Treating H2SO4 as fully giving 2 H+ — the second ionisation is incomplete.

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