Brønsted–Lowry Theory of Acids & Bases (Cambridge (CIE) A Level Chemistry): Flashcards

A Brønsted-Lowry base is a species that accepts a proton (H⁺) using its lone pair of electrons.

True.

Water is amphoteric. It donates a proton to ammonia (acting as an acid) and accepts a proton from HCl (acting as a base).

HCl is the Brønsted-Lowry acid: it donates a proton to NH₃.

NH₃ is the Brønsted-Lowry base: it accepts the proton using its lone pair of electrons.

A Brønsted-Lowry base accepts a proton using its lone pair of electrons, and the resulting species is called its conjugate acid.

A hydrogen atom contains one proton, one electron and no neutrons. When hydrogen loses its electron to form H⁺, only the proton remains, so H⁺ is simply a lone proton.

False.

Brønsted-Lowry theory also applies to reactions in the gas phase, such as HCl (g) + NH₃ (g) → NH₄Cl (s).

Water acts as a Brønsted-Lowry base: it accepts a proton from HCl using a lone pair of electrons on the oxygen atom, forming H₃O⁺.

An amphoteric species is one that can act as both a Brønsted-Lowry acid (proton donor) and a Brønsted-Lowry base (proton acceptor). Water is a key example.

A strong acid is an acid that almost completely dissociates in aqueous solution, giving a high concentration of H⁺ ions. Examples include HCl, HNO₃ and H₂SO₄.

False.

Strength refers to the degree of dissociation, not concentration. A weak acid can be concentrated and a strong acid can be dilute. Use dilute/concentrated for concentration.

A weak acid is an acid that only partially dissociates in aqueous solution, establishing an equilibrium. Examples include ethanoic acid, HCN and H₂CO₃.

For a strong acid, the equilibrium position lies to the right and dissociation is essentially complete. For a weak acid, the equilibrium position lies to the left and dissociation is partial.

H₂SO₄ has two ionisable hydrogen ions. The first ionisation (H₂SO₄ → H⁺ + HSO₄^-) is essentially complete, so it acts as a strong acid. The second ionisation (HSO₄^- ⇌ H⁺ + SO₄^2-) is only partial, so HSO₄^- acts as a weak acid.

True.

Group 1 metal hydroxides dissociate almost completely in aqueous solution, giving a high concentration of OH^- ions, so they are strong bases.

Ammonia (NH₃) and amines are weak bases. They are weak because they only partially dissociate in aqueous solution, establishing an equilibrium with a low concentration of OH^- ions.

Hydrogen ions in aqueous solution can be written as H⁺ or as H₃O⁺ if water is included in the equation.

For a strong acid, dissociation is complete, so [H⁺] equals the acid concentration. For a weak acid, only partial dissociation occurs, so [H⁺] is much less than the acid concentration and K_a must be used to find [H⁺].

pH is defined as the negative logarithm (base 10) of the hydrogen ion concentration: pH = -log₁₀ [H⁺ (aq)]. It measures how acidic or alkaline a solution is.

True.

Water dissociates slightly to give equal concentrations of H⁺ and OH^- ions, both at 10^-7 mol dm^-3. Substituting into pH = -log[H⁺] gives pH = 7.

An ionic product of water (K_w) is defined as K_w = [H⁺] [OH^-]. At 298 K, K_w = 1.0 x 10^-14 mol² dm^-6.

Acidic solutions always have [H⁺] greater than 10^-7 mol dm^-3, giving a pH below 7.

K_c = [H⁺] [OH^-] / [H₂O]

Because [H₂O] is effectively constant, it is combined with K_c to give K_w = [H⁺] [OH^-].

False.

Basic solutions have a lower [H⁺] than pure water. [H⁺] is less than 10^-7 mol dm^-3, giving pH > 7.

pH uses a log₁₀ scale. Each unit decrease in pH means [H⁺] is multiplied by 10, because pH = -log[H⁺] and a decrease of 1 in a log scale equals a tenfold increase in the original quantity.

The higher the [OH^-] of a base, the higher the pH of the solution.

pH = -log(10^-3) = 3.

This is acidic because pH 3 is below 7 and [H⁺] > 10^-7 mol dm^-3.

1. pH measurement (strong acid has lower pH).

2. Electrical conductivity (strong acid conducts better due to higher [H⁺]).

3. Reactivity with a metal (strong acid produces H₂ more vigorously).

True.

HCl is a strong acid and fully dissociates, giving [H⁺] = 0.1 mol dm^-3 and pH = 1. Ethanoic acid only partially dissociates, so [H⁺] is lower and pH ≈ 2.9.

A neutralisation reaction is one in which an acid and a base react to form a salt and water. The net ionic equation is: H⁺ (aq) + OH^- (aq) → H₂O (l).

Hydrochloric acid forms chloride salts, sulfuric acid forms sulfate salts, and nitric acid forms nitrate salts.

A strong acid has a higher concentration of H⁺ ions because it fully dissociates. More H⁺ ions are available to react with Mg, so more H₂ gas is produced at a faster rate.

False.

Spectator ions (e.g. Na⁺ and Cl^-) do not take part in forming water. Only H⁺ and OH^- combine to form H₂O (l).

A stronger acid dissociates more completely, producing a higher concentration of H⁺ and A^- ions in solution. More ions means greater electrical conductivity.

In a neutralisation reaction, the hydrogen ion from the acid reacts with the hydroxide ion from the base to form water.

Carbon dioxide gas is also produced: acid + metal carbonate → salt + water + carbon dioxide.

An equivalence point is the point in a titration at which the moles of acid exactly equal the moles of alkali. It is located within the steep vertical region of the pH titration curve.

The equivalence point pH is 7 because all the H⁺ ions have been neutralised by OH^- ions, leaving only water and a neutral salt in solution.

False.

The equivalence point is above pH 7 (around pH 9) because the conjugate base (e.g. CH₃COO^-) is a relatively strong base, making the solution slightly alkaline at the equivalence point.

When a strong acid is titrated with a weak alkali (e.g. NH₃), the equivalence point pH is approximately 5.5 because the conjugate acid (NH₄⁺) is a weak acid, making the solution slightly acidic at the equivalence point.

There is no vertical region. Instead there is a point of inflexion at the equivalence point, making it difficult to read the equivalence point accurately from the curve.

False.

A weak acid only partially dissociates, so [H⁺] is low. The initial pH of a weak acid solution is approximately 2-3, not 1.

The pH rises sharply past the equivalence point as excess OH^- ions from the strong alkali accumulate, driving the pH up to approximately 13–14.

In a titration, the end point is the point at which the indicator changes colour, while the equivalence point is where moles of acid equal moles of alkali.

A strong acid starts at pH 1-2 and shows a steep vertical section at the equivalence point (pH 7). A weak acid starts at pH 2-3, has a more gradual initial rise and the equivalence point is above pH 7 (around pH 9).

An indicator is a substance that changes colour over a specific pH range. In titrations, it is used to signal the endpoint, which should coincide with the equivalence point.

Methyl orange changes colour between pH 3.1 and 4.4 (red in acid, yellow in alkali).

Phenolphthalein changes colour between pH 8.3 and 10.0 (colourless in acid, pink in alkali).

False.

Phenolphthalein changes colour at pH 8.3-10.0, but the equivalence point for a strong acid + weak alkali titration is around pH 5.5. Only methyl orange is suitable here.

For a weak acid + strong alkali titration, only phenolphthalein is a suitable indicator because its pH range of 8.3-10.0 falls within the vertical region of the curve near the equivalence point.

The vertical region of the titration curve spans a wide pH range that includes both pH 3.1–4.4 (methyl orange) and pH 8.3–10.0 (phenolphthalein). Both indicators change colour within this steep section, so either gives an accurate endpoint.

False.

A weak acid + weak alkali titration has no sharp vertical region. Neither indicator changes colour within a clear endpoint and a different method should be used.

In the vertical region, a tiny addition of titrant causes a large pH change. If the indicator's colour-change range falls here, the colour change is sharp and sudden, giving an accurate endpoint. Outside this region, the pH changes gradually and the endpoint is imprecise.

Methyl orange is red in acidic solution and yellow in alkaline solution, changing colour between pH 3.1 and 4.4.

Methyl orange should be chosen. The equivalence point pH is around 5.5, which falls within methyl orange's colour-change range of pH 3.1-4.4. Phenolphthalein changes colour above pH 8, which is outside the vertical region.