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Topic 2: Aqueous solutions and acidity

Describe acids and bases qualitatively, distinguish between strong and weak acids using the extent of ionisation, calculate the pH of strong acid and base solutions, and write balanced equations for the reactions of acids with metals, carbonates and hydroxides

A focused answer to the QCE Chemistry Unit 2 dot point on acidity. Defines acids and bases qualitatively, introduces the pH scale and Kw, distinguishes strong from weak acids by extent of ionisation, calculates pH of strong acid and base solutions, and writes balanced equations for acid reactions with active metals, metal carbonates and metal hydroxides.

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What this dot point is asking

QCAA wants you to define acids and bases qualitatively, distinguish strong from weak acids/bases by extent of ionisation, calculate the pH of strong acid and base solutions of known concentration, and write balanced equations for the three classic reactions of acids (with active metals, with metal carbonates, with metal hydroxides). The fuller Bronsted-Lowry treatment, the weak-acid Ka calculation and buffers come in Unit 3; Unit 2 sets up the framework.

The answer

An acid is a substance that increases the concentration of hydrogen ions (H+, or more accurately the hydronium ion H_3O+) in aqueous solution. A base is a substance that increases the concentration of hydroxide ions (OH-). The pH scale quantifies acidity on a log scale from about 0 (very acidic) through 7 (neutral) to about 14 (very basic).

Acids and bases (Arrhenius framing)

An acid donates H+ to water:

HCl(g)→waterH(aq)++Cl(aq)−HCl_{(g)} \xrightarrow{\text{water}} H^+_{(aq)} + Cl^-_{(aq)}

A base produces OH- in water:

NaOH(s)→waterNa(aq)++OH(aq)−NaOH_{(s)} \xrightarrow{\text{water}} Na^+_{(aq)} + OH^-_{(aq)}

This Arrhenius picture is the foundation; Unit 3 extends it to Bronsted-Lowry (proton transfer) which handles acids and bases that do not contain H+ or OH- explicitly (e.g. NH_3 as a base).

H+ in water is more accurately written as H_3O+ (hydronium ion), reflecting that the proton is solvated by a water molecule. QCAA accepts either H+ or H_3O+ at Unit 2 level.

Strong vs weak acids and bases

Strong acid or base: ionises essentially completely in water. Written with a single arrow. The concentration of dissolved H+ (or OH-) equals the concentration of acid (or base) added.

  • Strong acids: HCl, HBr, HI, HNO_3, H_2SO_4 (first ionisation), HClO_4.
  • Strong bases: NaOH, KOH and other group 1 hydroxides; Ca(OH)_2 (slightly soluble but fully ionised once dissolved).

Weak acid or base: ionises only partially in water. Written with a double arrow. Only a small fraction donates or accepts a proton at any instant.

  • Weak acids: CH_3COOH (ethanoic), HF, H_2CO_3, HCN, NH_4+.
  • Weak bases: NH_3, organic amines, the conjugate bases of weak acids.

Extent of ionisation is set by the position of the dissociation equilibrium, formalised by Ka or Kb in Unit 3. At Unit 2 level you recognise the categorical distinction and predict the qualitative effect on pH (weak acids give less acidic solutions than strong acids of the same concentration).

Strength is independent of concentration: a 0.001 mol/L solution of HCl is still a strong acid (fully ionised), just dilute.

The pH scale and K_w

Water self-ionises slightly:

2H2O(l)⇌H3O(aq)++OH(aq)−2H_2O_{(l)} \rightleftharpoons H_3O^+_{(aq)} + OH^-_{(aq)}

At 25 degrees C the equilibrium gives [H_3O+] = [OH-] = 1.0 x 10^-7 mol/L.

The ion product of water:

Kw=[H3O+][OH−]=1.0×10−14  at 25 degrees CK_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}\;\text{at 25 degrees C}

Definitions on the log scale:

pH=−log⁡10[H3O+]pH = -\log_{10}[H_3O^+]

pOH=−log⁡10[OH−]pOH = -\log_{10}[OH^-]

pH+pOH=14.00  at 25 degrees CpH + pOH = 14.00\;\text{at 25 degrees C}

Acidic solutions have [H_3O+] > 10^-7 mol/L, so pH < 7. Basic solutions have [H_3O+] < 10^-7 mol/L, so pH > 7. Neutral solutions have pH = 7 (at 25 degrees C). A pH change of 1 unit corresponds to a 10-fold change in [H_3O+]; pH 2 is 100 times more acidic than pH 4.

Calculating pH of a strong acid solution

Strong acid fully ionises, so [H_3O+] equals the formal acid concentration.

  • 0.10 mol/L HCl: [H_3O+] = 0.10 mol/L; pH = -log(0.10) = 1.00.
  • 0.0010 mol/L HNO_3: [H_3O+] = 1.0 x 10^-3 mol/L; pH = 3.00.
  • 1.5 mol/L H_2SO_4 (first ionisation only): [H_3O+] approximately 1.5 mol/L; pH approximately -0.18. (Second ionisation of HSO_4- is weak and contributes more H+; treat fully only in Unit 3.)

Limit: very dilute strong acids (below about 10^-6 mol/L) cannot ignore the contribution of water's own ionisation, so the pH approaches but never exceeds 7. QCE problems usually stay above this limit.

Calculating pH of a strong base solution

Strong base fully ionises, so [OH-] equals the formal base concentration. Then either:

  • Use pOH = -log[OH-] and pH = 14 - pOH.
  • Or use [H_3O+] = K_w / [OH-] and pH = -log[H_3O+].

Both routes give the same answer.

  • 0.10 mol/L NaOH: [OH-] = 0.10 mol/L; pOH = 1.00; pH = 13.00.
  • 0.020 mol/L Ca(OH)_2: each formula unit gives 2 OH-, so [OH-] = 0.040 mol/L; pOH = 1.40; pH = 12.60.

Reactions of acids

Three reaction types are required at QCE Unit 2 level. All produce a salt; two also produce a gas; one also produces only water (neutralisation).

Acid + active metal -> salt + hydrogen gas.

M(s)+2H(aq)+→M(aq)2++H2(g)  (for a 2+ metal)M_{(s)} + 2H^+_{(aq)} \rightarrow M^{2+}_{(aq)} + H_{2(g)}\;\text{(for a 2+ metal)}

Worked: Zn(s) + 2HCl(aq) -> ZnCl_2(aq) + H_2(g). Effervescence; test the gas with a lit splint, hear a squeaky pop.

Active metals (above hydrogen in the activity series): K, Na, Ca, Mg, Al, Zn, Fe. Below hydrogen (Cu, Ag, Au) do not react with dilute non-oxidising acids.

Acid + metal carbonate (or hydrogencarbonate) -> salt + water + carbon dioxide.

CaCO3(s)+2HCl(aq)→CaCl2(aq)+H2O(l)+CO2(g)CaCO_{3(s)} + 2HCl_{(aq)} \rightarrow CaCl_{2(aq)} + H_2O_{(l)} + CO_{2(g)}

Effervescence; test the gas by bubbling through limewater (Ca(OH)_2) - it goes milky as CaCO_3 reprecipitates.

Acid + metal hydroxide -> salt + water (neutralisation).

HCl(aq)+NaOH(aq)→NaCl(aq)+H2O(l)HCl_{(aq)} + NaOH_{(aq)} \rightarrow NaCl_{(aq)} + H_2O_{(l)}

No gas, no precipitate. Net ionic equation H+(aq) + OH-(aq) -> H_2O(l) is the same for any strong acid plus strong base.

Acid + ammonia is sometimes added as a fourth class: HCl(aq) + NH_3(aq) -> NH_4Cl(aq); the gaseous demonstration NH_3(g) + HCl(g) -> NH_4Cl(s) gives a white smoke ring.

Indicators

pH indicators are weak acids or bases whose protonated and deprotonated forms have different colours. Common ones for QCE level:

Indicator pH range Colour change (low pH to high pH)
Methyl orange 3.1 to 4.4 Red to yellow
Bromothymol blue 6.0 to 7.6 Yellow to blue
Phenolphthalein 8.3 to 10.0 Colourless to pink
Universal indicator 1 to 13 Red, orange, yellow, green, blue, violet

Choice of indicator for a titration depends on the pH at the equivalence point of the specific acid-base combination (covered in Unit 3).

Common traps

Confusing strength and concentration. Strong/weak refers to extent of ionisation; concentrated/dilute refers to mol/L. A 0.001 mol/L HCl solution is dilute but the acid is strong.

Forgetting that base concentration gives [OH-], not [H+]. Calculate pOH first or convert via K_w.

Reporting pH to too many decimal places. pH is a logarithm; only the digits after the decimal count as significant figures. pH = 1.60 has 2 sig fig (the leading 1 is the order-of-magnitude digit).

Writing the wrong number of equivalents for diprotic acid or base. H_2SO_4 contributes 2 H+ per formula unit if fully ionised; Ca(OH)_2 contributes 2 OH-. In Unit 2 the second ionisation of H_2SO_4 is sometimes simplified.

Predicting gas evolution from neutralisation. Acid + hydroxide -> salt + water only. No gas. Carbon dioxide comes only from carbonates and hydrogencarbonates; hydrogen comes only from active metals.

In one sentence

Acids increase [H_3O+] and bases increase [OH-] in aqueous solution; the pH scale (pH = -log[H_3O+], with K_w = [H_3O+][OH-] = 1.0 x 10^-14 at 25 degrees C, so pH + pOH = 14) measures acidity logarithmically; strong acids and bases are fully ionised (so [H_3O+] or [OH-] equals the formal concentration), weak acids and bases only partially ionised (giving less extreme pH values than the same concentration of strong acid or base); and acids react with active metals to give hydrogen, with carbonates to give carbon dioxide, and with metal hydroxides to give a neutralised salt solution and water.

Past exam questions, worked

Real questions from past QCAA papers on this dot point, with our answer explainer.

2024 QCAA-style4 marks(a) Calculate the pH of a 0.0250 mol/L solution of hydrochloric acid. (b) Calculate the pH of a 0.0250 mol/L solution of sodium hydroxide. (c) Predict, with reasoning, whether the pH of a 0.0250 mol/L solution of ethanoic acid (a weak acid) will be lower, higher than, or equal to your answer in (a).
Show worked answer →

A 4-mark answer needs both pH calculations and the qualitative comparison.

(a) HCl is a strong acid (fully ionised). [H_3O+] = 0.0250 mol/L.

pH = -log_10[H_3O+] = -log(0.0250) = 1.60.

(b) NaOH is a strong base (fully ionised). [OH-] = 0.0250 mol/L.

pOH = -log(0.0250) = 1.60. pH = 14.00 - 1.60 = 12.40.

Alternative: [H_3O+] = K_w / [OH-] = 1.0 x 10^-14 / 0.0250 = 4.0 x 10^-13 mol/L. pH = -log(4.0 x 10^-13) = 12.40.

(c) Ethanoic acid (weak). Ethanoic acid is only partially ionised; not all 0.0250 mol/L of the acid donates a proton. So [H_3O+] is much less than 0.0250 mol/L, and pH is higher (less acidic) than the strong acid at the same concentration. The pH of 0.0250 mol/L CH_3COOH is about 3.1, which is greater than 1.60.

Markers reward correct pH calculations to 2 dp, the use of pOH or K_w, and the explicit "partial ionisation -> lower [H_3O+] -> higher pH" reasoning. The Bronsted-Lowry model and Ka treatment are extended in Unit 3.

2023 QCAA-style3 marksWrite balanced molecular equations and identify the gas evolved (if any) for the reactions of dilute hydrochloric acid with (a) magnesium metal, (b) solid calcium carbonate, and (c) aqueous sodium hydroxide.
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A 3-mark answer needs all three equations, balanced with state symbols, and the gases identified where relevant.

(a) Acid + active metal. Mg(s) + 2HCl(aq) -> MgCl_2(aq) + H_2(g). Gas evolved: hydrogen.

(b) Acid + metal carbonate. CaCO_3(s) + 2HCl(aq) -> CaCl_2(aq) + H_2O(l) + CO_2(g). Gas evolved: carbon dioxide.

(c) Acid + metal hydroxide (neutralisation). HCl(aq) + NaOH(aq) -> NaCl(aq) + H_2O(l). No gas evolved.

The three reactions illustrate the three classic acid behaviours: displacement of H_2 by an active metal, decomposition of carbonate to release CO_2 and water, and neutralisation to give a salt and water.

Markers reward balanced equations, correct state symbols, and explicit identification of the gas (or lack of one).

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