IB Chemistry HL R3.1 (Proton Transfer Reactions) Flashcards

Brønsted-Lowry acid

A species that donates a proton (H⁺) to another species.

Brønsted-Lowry base

A species that accepts a proton (H⁺) from another species, using a lone pair of electrons.

Proton

An H⁺ ion. In water it exists as the hydronium ion, H₃O⁺, so H⁺(aq) and H₃O⁺(aq) are used interchangeably.

Conjugate acid

The species formed when a base accepts a proton (base + H⁺).

Conjugate base

The species formed when an acid donates a proton (acid − H⁺).

Conjugate acid-base pair

Two species that differ by exactly one H⁺, e.g. NH₄⁺/NH₃ or H₂O/OH⁻.

Conjugate pair strength rule

The stronger the acid, the weaker its conjugate base. The weaker the acid, the stronger its conjugate base.

Amphiprotic species

A species that can both donate and accept a proton (e.g. H₂O, HCO₃⁻, HSO₄⁻, H₂PO₄⁻).

Amphoteric species

A species that can react as both an acid and a base. All amphiprotic species are amphoteric, but not all amphoteric species are amphiprotic (e.g. Al₂O₃ and ZnO have no protons to donate).

Water as amphiprotic

Water can act as an acid (H₂O → H⁺ + OH⁻) or as a base (H₂O + H⁺ → H₃O⁺).

Autoionization (self-ionization) of water

2H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq). A very small fraction of water molecules transfer a proton to other water molecules.

Writing acid ionization in water

HA(aq) + H₂O(l) ⇌ H₃O⁺(aq) + A⁻(aq). Water acts as the base.

Writing base ionization in water

B(aq) + H₂O(l) ⇌ BH⁺(aq) + OH⁻(aq). Water acts as the acid.

Alkali

A base that is soluble in water, producing OH⁻ ions (e.g. NaOH, KOH).

Base vs alkali

All alkalis are bases, but not all bases are alkalis. Insoluble bases such as CuO and Mg(OH)₂ are bases but not alkalis.

Monoprotic acid

An acid that donates one proton per molecule (e.g. HCl, CH₃COOH).

Polyprotic (diprotic/triprotic) acid

An acid that can donate more than one proton per molecule (e.g. H₂SO₄ is diprotic, H₃PO₄ is triprotic).

Basicity of an acid

The number of H⁺ ions one molecule of the acid can donate. Only H bonded to O (or acidic H) counts, so CH₃COOH has a basicity of 1.

Lewis acid (link to R3.4)

An electron-pair acceptor. Lewis base

Ligand and coordinate bond (link)

A Lewis base donating a lone pair to a metal ion forms a coordinate (dative) bond.

pH

pH = −log₁₀[H⁺]. A logarithmic measure of hydrogen ion concentration.

Calculating [H⁺] from pH

[H⁺] = 10^(−pH) mol dm⁻³.

pOH

pOH = −log₁₀[OH⁻]. [OH⁻] = 10^(−pOH).

Logarithmic nature of pH

A change of 1 pH unit means a tenfold change in [H⁺]. A change of 2 units means a hundredfold change.

Ionic product constant of water (Kw)

Kw = [H⁺][OH⁻] = 1.00 × 10⁻¹⁴ at 298 K.

pH + pOH

pH + pOH = 14.00 at 298 K (pKw = 14.00).

Neutral solution

A solution where [H⁺] = [OH⁻]. At 298 K this is pH 7.00.

Effect of temperature on Kw

The ionization of water is endothermic, so heating increases Kw and [H⁺]. pH of pure water falls below 7 at higher temperatures, but the water is still neutral because [H⁺] = [OH⁻].

Acidic, neutral, basic at 298 K

Acidic has pH < 7 ([H⁺] > [OH⁻]). Neutral has pH = 7. Basic has pH > 7 ([OH⁻] > [H⁺]).

pH of strong monoprotic acid

pH = −log[acid], since ionization is essentially complete. For example, 0.010 M HCl has pH 2.00.

pH of strong base

Find [OH⁻] first (from the base formula, e.g. Ba(OH)₂ gives 2 OH⁻), then pOH = −log[OH⁻], then pH = 14 − pOH.

Diluting a strong acid 10 times

pH rises by 1 unit.

Significant figures for pH

The number of decimal places in a pH value equals the number of significant figures in the concentration.

Strong acid

An acid that ionizes (dissociates) completely in aqueous solution. The equilibrium lies fully to the right (→).

Weak acid

An acid that ionizes only partially in aqueous solution. The equilibrium (⇌) lies mainly to the left.

Strong base

A base that ionizes or dissociates completely in aqueous solution.

Weak base

A base that only partially reacts with water to produce OH⁻. The equilibrium (⇌) lies mainly to the left.

Common strong acids

HCl, HBr, HI, HNO₃, H₂SO₄ (first ionization), HClO₄.

Common weak acids

Ethanoic acid (CH₃COOH), carbonic acid (H₂CO₃), HF, HCN, other carboxylic acids, H₃PO₄.

Common strong bases

Group 1 hydroxides (LiOH, NaOH, KOH) and Ba(OH)₂ (also Ca(OH)₂, Sr(OH)₂ largely).

Common weak bases

Ammonia (NH₃) and amines (e.g. CH₃NH₂, C₂H₅NH₂).

Strength vs concentration

Strength is the extent of ionization. Concentration is the amount of solute per dm³. A dilute strong acid can have a higher pH than a concentrated weak acid.

Strength vs concentration terminology

“Strong/weak” describes ionization. “Concentrated/dilute” describes how much solute is present.

Distinguishing strong from weak (equal concentrations)

The strong acid has lower pH, higher electrical conductivity, and faster reaction with metals and carbonates (higher [H⁺]).

Same neutralization amount

Equal volumes and concentrations of a strong and a weak monoprotic acid neutralize the same amount of base, because the weak acid keeps ionizing as H⁺ is removed (Le Châtelier).

Why ethanoic acid conducts poorly

It is only partly ionized, so few mobile ions are present.

Acid dissociation constant (Ka)

Ka = [H⁺][A⁻]/[HA] for HA ⇌ H⁺ + A⁻. A larger Ka means a stronger acid.

Base dissociation constant (Kb)

Kb = [BH⁺][OH⁻]/[B] for B + H₂O ⇌ BH⁺ + OH⁻. A larger Kb means a stronger base.

Water is not in Ka or Kb expressions

[H₂O] is essentially constant, so it is incorporated into the constant.

pKa

pKa = −log Ka. A smaller pKa means a stronger acid.

pKb

pKb = −log Kb. A smaller pKb means a stronger base.

Ka from pKa

Ka = 10^(−pKa).

Relationship Ka × Kb

For a conjugate pair, Ka × Kb = Kw = 1.00 × 10⁻¹⁴ at 298 K.

Relationship pKa + pKb

For a conjugate pair, pKa + pKb = 14.00 at 298 K.

Conjugate strength using Ka

A weak acid has a weak conjugate base. A very weak acid has a conjugate base that is a stronger (but still weak) base. A strong acid has a conjugate base with no basic character (e.g. Cl⁻).

pH of a weak acid (approximation)

If ionization is small, [HA] ≈ initial concentration and [H⁺] = [A⁻], so [H⁺] = √(Ka × [HA]).

Weak acid approximation conditions

Valid when the degree of ionization is small (typically < 5%, i.e. [HA]/Ka > about 400). State the assumption in the answer.

pH of a weak base (approximation)

[OH⁻] = √(Kb × [B]), then pOH = −log[OH⁻] and pH = 14 − pOH.

Percentage ionization

(amount ionized / initial amount) × 100. For a weak acid this is ([H⁺]/[HA]₀) × 100. It increases as concentration decreases.

ICE table

Initial, Change, Equilibrium table used to find equilibrium concentrations for Ka and Kb calculations.

Using Ka to compare acid strength

Compare Ka or pKa values at the same temperature. Do not compare pH values unless the concentrations are equal.

Effect of temperature on Ka and Kb

These constants are temperature dependent. Ka, Kb and Kw all change with temperature.

Acid + metal

Acid + reactive metal → salt + hydrogen gas (e.g. Mg + 2HCl → MgCl₂ + H₂).

Acid + metal oxide or hydroxide

Acid + base → salt + water (neutralization), e.g. CuO + H₂SO₄ → CuSO₄ + H₂O.

Acid + carbonate

Acid + carbonate → salt + water + carbon dioxide (e.g. CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂).

Acid + hydrogencarbonate

Acid + HCO₃⁻ → salt + water + carbon dioxide.

Acid + ammonia

Produces an ammonium salt only, with no water formed (NH₃ + HCl → NH₄Cl).

Ionic equation for neutralization

H⁺(aq) + OH⁻(aq) → H₂O(l).

Enthalpy of neutralization

The enthalpy change when an acid reacts with a base to form 1 mol of water. For strong acid plus strong base it is about −57 kJ mol⁻¹.

Weak acid neutralization enthalpy

Less exothermic than −57 kJ mol⁻¹, because some energy is needed to ionize the weak acid.

Salt

An ionic compound formed when the H⁺ of an acid is replaced by a metal ion or ammonium ion.

Test for CO₂ from carbonate reaction

Bubble the gas through limewater. It turns milky or cloudy (CaCO₃ forms).

Strong acid + strong base salt (e.g. NaCl)

Neutral solution, pH about 7. Neither ion reacts with water.

Strong acid + weak base salt (e.g. NH₄Cl)

Acidic solution. The cation NH₄⁺ is a weak acid that donates H⁺ to water.

Weak acid + strong base salt (e.g. CH₃COONa)

Basic solution. The anion CH₃COO⁻ is a weak base that accepts H⁺ from water, producing OH⁻.

Weak acid + weak base salt

pH depends on relative Ka and Kb values; it is close to neutral if they are similar.

Small, highly charged metal cations (e.g. Al³⁺, Fe³⁺)

Make water acidic because the hydrated ion [M(H₂O)₆]³⁺ donates H⁺ from its coordinated water ligands.

Anions of weak acids

Act as bases in water (e.g. CO₃²⁻, CH₃COO⁻, F⁻).

Anions of strong acids

Cl⁻, Br⁻, NO₃⁻ and ClO₄⁻ are neutral in water.

Acid-base indicator

A weak acid (or base) whose acid form (HIn) and conjugate form (In⁻) have different colors. The color change shows the pH change.

How an indicator works

HIn(aq) ⇌ H⁺(aq) + In⁻(aq). In acid, equilibrium shifts left and the HIn color shows. In base, it shifts right and the In⁻ color shows.

Indicator pKa and end point

The color changes sharply near pH = pKa of the indicator, where [HIn] = [In⁻].

Indicator pH range

Approximately pKa ± 1. The range must lie on the steep (vertical) part of the pH curve.

Methyl orange

Red in acid, yellow in base. Changes at about pH 3.2 to 4.4.

Phenolphthalein

Colorless in acid, pink in base. Changes at about pH 8.2 to 10.0.

Bromothymol blue

Yellow in acid, blue in base. Changes at about pH 6.0 to 7.6.

Universal indicator

A mixture of indicators giving a continuous color range across a wide pH range (approximate pH only).

Titration

A quantitative technique in which a solution of known concentration (titrant) reacts with a measured volume of another solution (analyte) to find an unknown concentration.

Equivalence point

The point where the reactants have been added in exactly the stoichiometric ratio (moles of acid and base in the ratio of the equation).

End point

The point where the indicator changes color. It should be as close as possible to the equivalence point.

Standard solution

A solution of accurately known concentration.

Titration procedure essentials

Rinse the burette with titrant and the pipette with the analyte solution. Rinse the conical flask with deionized water only. Remove the funnel and fill the tip. Take concordant results (within 0.10 cm³).

Why rinse the flask with water only

Water does not change the number of moles of analyte, but rinsing with solution would add extra moles.

Titration calculation steps

Find moles of the known solution (n = cV), use the mole ratio from the balanced equation, then find the unknown concentration c = n/V.

Primary standard

A pure, stable, non-hygroscopic substance of known molar mass used to make an accurate standard solution (e.g. anhydrous Na₂CO₃, potassium hydrogen phthalate).

Why NaOH is not a primary standard

It absorbs water and CO₂ from the air, so its concentration is not known accurately.

Back titration

Adding excess reagent to the sample, then titrating the leftover reagent. Used for insoluble substances or slow reactions.

Strong acid + strong base curve

Starts at low pH, rises sharply through a vertical section around pH 7, then levels at high pH. Equivalence point pH = 7. Suitable indicators include methyl orange, bromothymol blue, phenolphthalein.

Weak acid + strong base curve

Starts at higher pH than for a strong acid. Has a buffer region and an equivalence point above pH 7 (about 8 to 9). Suitable indicator

Strong acid + weak base curve

Starts at low pH and the equivalence point is below pH 7 (about 5 to 6). Suitable indicator