Mission 1: Is water an acid or a base?

Acids give protons, bases take them

• A Brønsted–Lowry acid is a proton (H⁺) donor.
• A Brønsted–Lowry base accepts a proton, using a lone pair.
• HNO₃ + NH₃ → NH₄⁺ + NO₃⁻: HNO₃ donates H⁺; NH₃ accepts it.
• H⁺(aq) and H₃O⁺(aq) are two ways of writing the same hydrated proton.

Base or alkali?

• An alkali is a base that dissolves in water, releasing OH⁻(aq): NaOH, KOH, NH₃(aq).
• Every alkali is a base; not every base is an alkali.
• Iron(III) oxide neutralizes acids, so it is a base. It is insoluble, so not an alkali.
• A base need not contain OH: NH₃, O²⁻ and CO₃²⁻ accept H⁺ with a lone pair.

Conjugate acid–base pairs

• A conjugate pair differs by exactly one proton.
• Conjugate base: remove one H⁺; the charge falls by one. HSO₄⁻ → SO₄²⁻.
• Conjugate acid: add one H⁺; the charge rises by one. NH₃ → NH₄⁺.
• Every proton transfer has two pairs: HNO₃/NO₃⁻ and NH₄⁺/NH₃.
• The acid and base that react are not a pair; each partner is across the arrow.

Amphiprotic species

• Amphiprotic: can act as a Brønsted–Lowry acid or base, depending on its partner.
• It needs an H it can lose and a lone pair to accept H⁺.
• Examples: H₂O, HSO₄⁻, H₂PO₄⁻, HPO₄²⁻.
• In any one reaction it plays one role only.

Writing both equations

• Water as a base: HNO₃ + H₂O → H₃O⁺ + NO₃⁻.
• Water as an acid: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻.
• HPO₄²⁻ as a base: HPO₄²⁻ + H₃O⁺ ⇌ H₂PO₄⁻ + H₂O.
• HPO₄²⁻ as an acid: HPO₄²⁻ + OH⁻ ⇌ PO₄³⁻ + H₂O.
• Check: as a base it gains one H; as an acid it loses one. Charges balance.

Mission 2: Lemon juice is about pH 2; black coffee is about pH 5. How many times more H⁺ is in the lemon juice?

What pH measures

• pH = −log₁₀[H⁺], with [H⁺] in mol dm⁻³. pH has no units.
• Rearranged: [H⁺] = 10⁻ᵖᴴ. Only the pH form is printed in the data booklet.
• Lower pH means higher [H⁺].
• One pH unit is a tenfold change in [H⁺]; two units, a hundredfold.
• Very concentrated solutions can have pH below 0 or above 14.

Non-integer pH: worked examples

• [H⁺] = 3.2 × 10⁻⁴ mol dm⁻³: pH = −log₁₀(3.2 × 10⁻⁴) = 3.49.
• pH 4.60: [H⁺] = 10⁻⁴·⁶⁰ = 2.5 × 10⁻⁵ mol dm⁻³.
• Use the log key (base 10), not ln.
• Sense check: pH 3.49 lies between 3 and 4, so [H⁺] lies between 10⁻⁴ and 10⁻³.

Measuring pH

• Universal indicator: match its colour to a chart to estimate pH, to about 1 unit.
• pH meter (probe): a precise reading, usually to 0.01.
• Calibrate the meter first with buffer solutions of known pH.
• Rinse the probe with distilled water between samples.

Water ionizes, a little: Kw

• Water ionizes slightly: H₂O(l) ⇌ H⁺(aq) + OH⁻(aq).
• Kw = [H⁺][OH⁻] = 1.00 × 10⁻¹⁴ at 298 K (data booklet).
• Every aqueous solution contains both H⁺ and OH⁻.
• Acidic: [H⁺] > [OH⁻]. Neutral: [H⁺] = [OH⁻]. Basic: [H⁺] < [OH⁻].
• At 298 K, neutral means both are 1.00 × 10⁻⁷ mol dm⁻³: pH 7.00.
• Kw changes with temperature, so neutral is pH 7.00 only at 298 K.

Using Kw: worked example

• The product is fixed, so [H⁺] up tenfold means [OH⁻] down tenfold.
• [OH⁻] = Kw ÷ [H⁺] and [H⁺] = Kw ÷ [OH⁻].
• 0.020 mol dm⁻³ KOH is fully ionized: [OH⁻] = 0.020 mol dm⁻³.
• [H⁺] = 1.00 × 10⁻¹⁴ ÷ 0.020 = 5.0 × 10⁻¹³ mol dm⁻³; pH = 12.30.

Mission 3: Can a strong acid have a higher pH than a weak acid?

Strong and weak

• Strong acids and bases ionize completely in water.
• Strong acids: HCl, HBr, HI, HNO₃, H₂SO₄. Strong bases: group 1 hydroxides.
• Weak acids and bases ionize only partly: the equilibrium lies to the left.
• Weak: carboxylic acids such as HCOOH, carbonic acid, NH₃ and amines.
• HCOOH(aq) ⇌ HCOO⁻(aq) + H⁺(aq). NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq).

Strength is not concentration

• Concentration: amount of acid per dm³. Strength: the extent of ionization.
• Diluting a strong acid raises its pH, but it stays strong.
• At equal concentration, a weak acid has a higher pH and lower conductivity than a strong one.
• It still neutralizes the same amount of base: as H⁺ is used up, more of the acid ionizes.

Which way does the equilibrium lie?

• An acid–base equilibrium lies on the side of the weaker acid and weaker base.
• The stronger the acid, the weaker its conjugate base.
• HNO₃ + H₂O → H₃O⁺ + NO₃⁻ is effectively complete: NO₃⁻ barely accepts H⁺.
• HCOOH + H₂O ⇌ HCOO⁻ + H₃O⁺ lies left: H₃O⁺ is a stronger acid than HCOOH.

Mission 4: Why can an antacid tablet make you burp?

Oxides and hydroxides: salt + water

• In neutralization, a base accepts the H⁺ ions from an acid.
• Acid + metal hydroxide → salt + water: KOH + HNO₃ → KNO₃ + H₂O.
• Acid + metal oxide → salt + water: Fe₂O₃ + 6HCl → 2FeCl₃ + 3H₂O.
• Count the protons: each O²⁻ takes 2H⁺, so Fe₂O₃ needs 6 HCl.
• A salt forms when an acid's H is replaced by a metal ion or another cation, such as NH₄⁺.

Carbonates and hydrogencarbonates: CO₂ too

• Acid + carbonate or hydrogencarbonate → salt + water + carbon dioxide.
• CO₃²⁻ accepts two H⁺; HCO₃⁻ accepts one.
• K₂CO₃ + 2HNO₃ → 2KNO₃ + H₂O + CO₂.
• KHCO₃ + HNO₃ → KNO₃ + H₂O + CO₂.
• The fizzing is the CO₂ escaping.

Ammonia, amines and organic acids

• Ammonia or an amine + acid → a salt only; no water forms.
• NH₃ + HNO₃ → NH₄NO₃. CH₃CH₂NH₂ + HCl → CH₃CH₂NH₃Cl.
• The N lone pair takes the H⁺; there is no OH⁻ or O²⁻ to make water.
• Weak organic acids react in the same ways: 2HCOOH + K₂CO₃ → 2HCOOK + H₂O + CO₂.
• Metal + acid gives hydrogen. That is redox, not neutralization: see the redox lesson.

The parent acid and base of a salt

• A salt's anion comes from its parent acid; its cation from its parent base.
• KNO₃: parent acid HNO₃, parent base KOH.
• (NH₄)₂SO₄: parent acid H₂SO₄, parent base NH₃.
• HCOOK (potassium methanoate): parent acid HCOOH, parent base KOH.
• Parents are the neutral acid and base, not the ions in the salt.

The pH curve: strong acid, strong base

• Flask: 25.0 cm³ of 0.100 mol dm⁻³ HCl. Burette: 0.100 mol dm⁻³ NaOH.
• Intercept: the pH of the flask solution before any base is added, 1.00.
• pH rises slowly, then jumps from about 3 to about 11 near 25.0 cm³.
• Equivalence point: amounts match the equation. Here it is pH 7 (at 298 K), mid-jump.
• After it, pH rises slowly towards 13 (about 12.5 by 50 cm³), as excess NaOH builds up.

Why the jump is so steep

• pH is a log scale. Early on, most of the H⁺ is still there, so pH barely moves.
• At 24.0 cm³, 1/25 of the H⁺ is left, in 49.0 cm³: pH about 2.7.
• At 26.0 cm³, a little OH⁻ is in excess: pH about 11.3.
• So 2 cm³ moves the pH by about 8.6 units: the near-vertical section.

The big ideas

• Acids donate H⁺; bases accept it with a lone pair. Conjugate pairs differ by one H⁺.
• Amphiprotic species, such as water, act as acid or base: one role per reaction.
• pH = −log₁₀[H⁺]: one unit is tenfold. Kw = [H⁺][OH⁻] = 1.00 × 10⁻¹⁴ at 298 K.
• Strong means fully ionized; weak, partly. Strength is not concentration.
• Oxides and hydroxides give salt + water; carbonates add CO₂; ammonia and amines give a salt only.
• Strong acid–strong base curve: starts at the flask's pH, jumps steeply, equivalence at pH 7 (298 K).