Acid Strength

Acid strength reflects an acid's capacity to donate protons (H⁺ ions). The more readily an acid releases H⁺ ions, the stronger it is. This property affects the degree of ionization, impacting the reactivity and pathways of chemical reactions. Acid strength is particularly important in titrations for determining the concentration of unknown solutions.

Acid strength refers to the ability of an acid to lose its H⁺ ion (proton) in a solution.

In the Bronsted-Lowry concept, acids and bases are defined by their ability to donate or accept H⁺ ions (protons).

• Acid Strength: Refers to the ease with which an acid loses its proton (H⁺). The easier the proton is removed, the stronger the acid.

For example, if an acid is represented by the chemical formula HA, it dissociates into a proton (H⁺) and an anion (A⁻):
HA → H⁺ + A⁻

NOTE

• The strength of an inorganic acid depends on the oxidation state of the atom bonded to the proton.
• Acid strength varies with the solvent used. For example, HCl is a strong acid in water but acts as a weak acid in glacial acetic acid.

1.0Acid Strength Trends in the Periodic Table

1. Within a Group:

• As the size of atom A increases down a group, the H–A bond strength decreases.
• This weaker bond leads to higher acid strength.

Example: HF < HCl < HBr < HI

1. Within a Period:

• Moving left to right across a period, atomic size decreases while electronegativity increases.
• This results in stronger bond polarity, increasing acid strength.

By understanding these factors and trends, the acid strengths of different compounds can be effectively compared and predicted.

2.0Strong and Weak Acids

Acids are categorised based on their ability to dissociate in a solution:

1. Strong Acids:

• Fully dissociate in a solution.
• High degree of dissociation.
• Highly corrosive and capable of causing severe burns.
• Represented by: HA + S ⇌ SH⁺ + A⁻ (where S is the solvent).
• Common examples include Hydrochloric acid (HCl), Sulfuric acid (H₂SO₄), Nitric acid (HNO₃), Hydrobromic acid (HBr), Perchloric acid (HClO₄)

2. Weak Acids:

• Partially dissociate in a solution.
• Low degree of dissociation, leaving undissociated molecules in the solution.
• Mildly corrosive and often present in food and biological systems.
• Represented by: HA ⇌ H⁺ + A⁻
• Common examples include acetic acid (CH₃COOH), benzoic acid (C₆H₅COOH), oxalic acid (H₂C₂O₄), phosphoric acid (H₃PO₄), methanoic acid (HCOOH), and hydrofluoric acid (HF).

Note:

• Strong acids are completely ionised in water, whereas weak acids leave undissociated molecules in solution.

3.0Acid Dissociation 

The strength of a weak acid is determined by its dissociation constant (Ka), given by:
Ka = [H⁺][A⁻] / [HA]

• A larger Ka value indicates a stronger acid with higher ionisation in water.
• A smaller Ka value represents a weaker acid with lower ionisation.

4.0Factors Affecting Acid Strength

1. Bond Strength:

• The ease of dissociation is determined by the strength of the bond between H⁺ and A⁻.

2. Weaker bonds: Require less energy to break, leading to stronger acids.
3. Stronger bonds: Require more energy to dissociate, resulting in weaker acids.
4. Bond Polarity:

• Polarity depends on the electronegativity of the atoms in the bond.

5. Highly polar bonds: Allow easier proton dissociation, increasing acid strength.
6. Low polarity: Makes proton release more difficult, resulting in weaker acids.
7. Inductive Effect:

• Electronegative atoms attract electron pairs, creating bond polarisation.
• A higher inductive effect facilitates proton release, strengthening the acid.