Unit 8 pH of acidic and alkaline solutions
Key unit competence: Interpret acid–base behaviour using different theories, explain the differences between strong and weak acids and bases, and calculate pH.
Key ideas
- Theories: Arrhenius (acids give H+ in water, bases give OH−); Brønsted–Lowry (acid = proton donor, base = proton acceptor; conjugate pairs: CH3COOH/CH3COO−, NH4+/NH3); Lewis (acid = electron-pair acceptor, e.g. BF3, AlCl3, H+; base = electron-pair donor, e.g. NH3).
- Strong acids and bases ionise completely (HCl, HNO3, H2SO4; NaOH, KOH); weak ones only partly (ethanoic acid, ammonia). For HA ⇌ H+ + A−: Ka = [H+][A−] ÷ [HA]; pKa = −log Ka (smaller pKa, stronger acid). For bases, Kb; for a conjugate pair Ka × Kb = Kw.
- Ionic product of water: Kw = [H+][OH−] = 1.0 × 10−14 mol2 dm−6 at 25 °C. pH = −log[H+]; pOH = −log[OH−]; pH + pOH = 14 at 25 °C. Neutral water: pH 7 at 25 °C (Kw increases with temperature, so neutral pH falls when hot).
- Weak acid: [H+] = √(Kac) (α small); degree of ionisation α = √(Ka ÷ c). Weak base: [OH−] = √(Kbc).
- Salt hydrolysis: salts of a strong acid and weak base are acidic (NH4Cl); of a weak acid and strong base are alkaline (CH3COONa, Na2CO3); of strong acid and strong base are neutral (NaCl). For CH3COO−: Kb = Kw ÷ Ka.
- A buffer resists changes in pH when small amounts of acid or alkali are added: a weak acid with its salt (ethanoic acid + sodium ethanoate) or a weak base with its salt (NH3 + NH4Cl). pH = pKa + log([salt] ÷ [acid]) (Henderson–Hasselbalch). Blood is buffered at pH 7.4 by H2CO3/HCO3−.
Common misconceptions
- "A strong acid and a concentrated acid mean the same thing." No: 'Strong' means fully ionised; 'concentrated' means a lot of acid per dm3.
- "pH 7 is always neutral." No: Only at 25 °C; at higher temperature Kw is larger and neutral pH is below 7.
- "A buffer keeps the pH completely constant." No: It only resists change for small additions; large amounts of acid or alkali overwhelm it.
Worked examples
Example 1. Explain, with equations, how an ammonia/ammonium chloride buffer works.
Solution
The buffer contains much NH3 (weak base) and much NH4+ (its conjugate acid, from NH4Cl).
Add a little acid: NH3 + H+ → NH4+; the added H+ is removed.
Add a little alkali: NH4+ + OH− → NH3 + H2O; the added OH− is removed.
The ratio [NH3]/[NH4+] changes very little, so the pH (about 9.25) hardly changes.
Example 2. Calculate the pH of the solution formed by mixing 25.0 cm3 of 0.20 mol dm−3 HCl with 25.0 cm3 of 0.10 mol dm−3 NaOH.
Solution
n(HCl) = 0.0050 mol; n(NaOH) = 0.0025 mol.
Excess H+ = 0.0025 mol in 50.0 cm3.
[H+] = 0.0025 ÷ 0.0500 = 0.050 mol dm−3.
pH = −log 0.050 = 1.30.
Practice questions
Define a Brønsted–Lowry base.
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A proton (H+) acceptor.
Common mistake: Saying it releases OH−.
Identify the conjugate pairs in NH3 + H2O ⇌ NH4+ + OH−.
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NH4+/NH3 and H2O/OH−.
Common mistake: Pairing NH3 with H2O.
Write the Ka expression for ethanoic acid.
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Ka = [H+][CH3COO−] ÷ [CH3COOH]
Common mistake: Including [H2O].
Calculate [H+] in a solution of pH 4.5.
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[H+] = 10−4.5 = 3.16 × 10−5 mol dm−3.
Common mistake: Writing 4.5 × 10−1.
Explain why AlCl3 is a Lewis acid.
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Aluminium has an empty orbital and can accept an electron pair.
Common mistake: Saying it donates H+.
Why is a solution of Na2CO3 alkaline?
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CO32− accepts protons from water: CO32− + H2O ⇌ HCO3− + OH−.
Common mistake: Saying sodium is alkaline.
Give one natural and one laboratory buffer.
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Blood (H2CO3/HCO3−); ethanoic acid/sodium ethanoate (or NH3/NH4Cl).
Common mistake: Giving HCl/NaCl.
Calculate the pH of 1.0 × 10−3 mol dm−3 KOH.
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pOH = 3; pH = 11.
Common mistake: Giving pH 3.
Why does a 0.1 mol dm−3 HCl solution have a lower pH than 0.1 mol dm−3 ethanoic acid?
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HCl is fully ionised, so [H+] is much higher.
Common mistake: Saying HCl is more concentrated.
Write the equation that defines Kw.
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Kw = [H+][OH−].
Common mistake: Including [H2O] in the expression.
Matching and drag-and-drop
Matching
Match each item in column A with the correct answer in column B. Write the letter next to each number.
Column A Column B 1. pH is defined as: a. −log[H+] 2. A smaller pKa means: b. Ethanoic acid and sodium ethanoate 3. Which mixture is a buffer? c. a stronger acid 4. A solution of ammonium chloride is: d. acidic Show answer
1 → a (−log[H+]) · 2 → c (a stronger acid) · 3 → b (Ethanoic acid and sodium ethanoate) · 4 → d (acidic)
- Negative logarithm of hydrogen ion concentration.
- Larger Ka.
- Weak acid and its conjugate base.
- NH4+ donates protons to water.
Drag and drop
Choose the correct word or phrase from the box for each gap. One word in the box is not needed.
NH4+ · acidic · alkaline · proton donor- A solution of sodium ethanoate is .
- According to Brønsted–Lowry, an acid is a .
- The conjugate acid of NH3 is .
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1: alkaline · 2: proton donor · 3: NH4+
- CH3COO− accepts protons.
- It gives H+.
- Base plus H+.
Matching
Match each item in column A with the correct answer in column B. Write the letter next to each number.
Column A Column B 1. A solution of sodium chloride is: a. neutral 2. For a conjugate acid–base pair, Ka × Kb equals: b. ionises completely in water 3. A strong acid is one that: c. Kw 4. The pH of 0.010 mol dm−3 NaOH at 25 °C is: d. 12 Show answer
1 → a (neutral) · 2 → c (Kw) · 3 → b (ionises completely in water) · 4 → d (12)
- Strong acid + strong base salt.
- At 25 °C, 1.0 × 10−14.
- e.g. HCl.
- pOH = 2; pH = 14 − 2.
Drag and drop
Choose the correct word or phrase from the box for each gap. One word in the box is not needed.
7.4 · CH3COO− · OH− · 1.0 × 10−14 mol2 dm−6- The conjugate base of CH3COOH is .
- The value of Kw at 25 °C is .
- The pH of human blood is about .
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1: CH3COO− · 2: 1.0 × 10−14 mol2 dm−6 · 3: 7.4
- Acid minus H+.
- [H+][OH−].
- Buffered by hydrogencarbonate.