Acid-Base Equilibria

This lesson covers: 

1. The definitions of acids and bases
2. The difference between strong and weak acids and bases
3. Neutralisation reactions between acids and bases
4. Reactions of acids with metals and metal compounds
5. The Brønsted-Lowry definitions of acids and bases

Acids are proton donors, bases are proton acceptors

An acid is defined as a substance that donates protons (H⁺) when dissolved in water. These protons combine with water molecules to form hydronium ions (H₃O⁺), indicating the acidic nature of the solution.

Examples of common acids include:

• Hydrochloric acid (HCl)
• Sulfuric acid (H₂SO₄)
• Nitric acid (HNO₃)
• Ethanoic acid (CH₃COOH)

A base, in contrast, is a substance that accepts protons. When a base dissolves in water, it can produce hydroxide ions (OH^-), making the solution alkaline. Bases that dissolve in water are specifically called alkalis.

Examples of common alkalis include:

• Sodium hydroxide (NaOH)
• Potassium hydroxide (KOH)
• Aqueous ammonia (NH₃ in H₂O)

Aqueous ammonia is a special case; it doesn’t produce hydroxide ions directly but accepts protons from water to form ammonium (NH₄⁺) and hydroxide ions (OH^-):

NH₃ + H₂O ➔ NH₄⁺ + OH^-

Strength depends on extent of dissociation

The strength of an acid or a base is determined by its ability to dissociate in water. This dissociation process can be reversible, as shown in the equations below:

• Acid dissociation: HA + H₂O ⇌ H₃O⁺ + A^-
• Base dissociation: B + H₂O ⇌ BH⁺ + OH^-

Strong acids and bases:

Strong acids, such as hydrochloric acid (HCl), and strong bases, like sodium hydroxide (NaOH), dissociate completely in water. This results in a significant release of H₃O⁺ and OH^- ions, respectively, with the forward reaction being predominantly favoured:

• HCl_(aq) ➔ H⁺_(aq) + Cl^-_(aq)
• NaOH_(aq) ➔ Na⁺_(aq) + OH^-_(aq)

Weak acids and bases:

Conversely, weak acids and bases, like ethanoic acid (CH₃COOH) and ammonia (NH₃), only partially dissociate in water, releasing fewer H₃O⁺ and OH^- ions. In these cases, the reverse reaction is favoured:

• CH₃COOH_(aq) ⇌ H⁺_(aq) + CH₃COO^-_(aq)
• NH_3(aq) + H₂O_(l) ⇌ NH₄⁺_(aq) + OH^-_(aq)

Differences between behaviour of strong and weak acids

The differences between strong and weak acids can be observed qualitatively through their reactions with reactive metals and by measuring their pH values or conductivity.

1. Reaction with reactive metals

• Strong acids - React vigorously with reactive metals (e.g., magnesium or zinc), producing a significant amount of hydrogen gas. This vigorous reaction is due to the high concentration of H₃O⁺ ions in strong acids, which readily react with the metal.
• Weak acids -  React slowly or not at all with reactive metals, producing little to no hydrogen gas due to the lower concentration of H₃O⁺ ions.

1. pH values

• Strong acids - Have very low pH values (typically below 2) due to high concentrations of H₃O⁺ ions. This can be measured with a pH meter, universal indicator, or litmus paper, all of which will show a highly acidic value.
• Weak acids - Have higher pH values (typically between 2 and 6), indicating lower concentrations of H₃O⁺ ions.

1. Conductivity

• Strong acids - Are good conductors of electricity because they have a high concentration of free H₃O⁺ ions and anions in solution to carry electrical charge
• Weak acids - Are poor conductors of electricity due to having fewer free ions in solution to carry electrical charge.

Neutralisation produces salt and water

Neutralisation is the reaction between an acid and an alkali, resulting in the formation of water and a salt. This reaction specifically involves the combination of H⁺ ions from the acid with OH^- ions from the alkali to form water (H₂O).

The general ionic equation for this process is:

H⁺_(aq) + OH^-_(aq) ➔ H₂O_(l)

A salt is an ionic compound formed when the H⁺ ions in an acid are replaced by metal ions or other positive ions, such as ammonium ions (NH₄⁺). In a neutralisation reaction, the salt is comprised of the cation from the base, which can be a metal or ammonium ion, and the anion from the acid.

The type of salt produced depends on the specific acid used in the reaction:

For instance, the reaction between hydrochloric acid and sodium hydroxide yields sodium chloride and water:

HCl_(aq) + NaOH_(aq) ➔ NaCl_(aq) + H₂O_(l)

Acids react with metals and metal compounds

1. Acids and reactive metals - Acids react with certain metals, producing a salt and hydrogen gas:

Metal + acid ➔ salt + hydrogen

For example, calcium reacting with sulfuric acid: Ca_(s) + H₂SO_4(aq) ➔ CaSO_4(aq) + H_2(g)

1. Acids and metal oxides - These reactions yield a salt and water:

Metal oxide + acid ➔ salt + water

An example is zinc oxide with hydrochloric acid: ZnO_(s) + 2HCl_(aq) ➔ ZnCl_2(aq) + H₂O_(l)

1. Acids and metal hydroxides - Similar to oxides, these reactions produce a salt and water:

Metal hydroxide + acid ➔ salt + water

For example, potassium hydroxide reacting with nitric acid: KOH_(aq) + HNO_3(aq) ➔ KNO_3(aq) + H₂O_(l)

1. Acids and metal carbonates - These reactions produce a salt, water, and carbon dioxide:

Metal carbonate + acid ➔ salt + water + carbon dioxide

An example is sodium carbonate with hydrochloric acid: Na₂CO_3(s) + 2HCl_(aq) ➔ 2NaCl_(aq) + H₂O_(l) + CO_2(g)

Brønsted-Lowry acids are proton donors, bases are proton acceptors

According to the Brønsted-Lowry theory:

• An acid is defined as a substance that donates a proton (H⁺).
• A base is defined as a substance that accepts a proton.

For instance, when a Brønsted-Lowry acid (HA) is dissolved in water, it donates a proton to a water molecule, forming a hydronium ion (H₃O⁺):

HA_(aq) + H₂O_(l) ⇌ H₃O⁺_(aq) + A^-_(aq)

Conversely, when a Brønsted-Lowry base (B) is mixed with water, it accepts a proton from a water molecule, forming a hydroxide ion (OH^-):

B_(aq) + H₂O_(l) ⇌ BH⁺_(aq) + OH^-_(aq)