Relative Strength of Acids and Bases

The ionisation or dissociation of all acids and bases is not the same. This implies that not all acids and bases are equally effective in producing of H+ and OH– ions in an aqueous solution. 

In addition, the strength of acids and bases varies widely. When acids and bases dissolve in water, they create variable amounts of hydronium or hydroxyl ions. ‘Strong’ and ‘weak’ refer to the acid or base’s strength. Acid and base solutions are categorised as ‘strong’ or ‘weak’ based on their electrical conductivity.

The Bronsted-Lowery definition

Johannes Bronsted and Martin Lowry separately defined acids and bases in 1923 based on a compound’s ability to donate or receive protons (H+ ions). Acids are described as having the ability to contribute protons in the form of hydrogen ions, while bases are characterised as having the ability to accept protons. 

Before and after the reaction, count the number of hydrogen present in each substance to identify whether it is acidic or basic. If the amount of hydrogen atoms in a substance decreases, the substance is said to be acidic (donates hydrogen ions). If the quantity of hydrogen atoms in a substance has increased, that substance is referred to as basic (accepts hydrogen ions). These definitions are often used for the reactant molecules on the left-hand side of the equation. Reversing the process identifies a distinct acid and base. Compared to the left side of the equation, the elements on the right are referred to as conjugate base and conjugate acid.

Strength of Acids or Bases

Acids and bases are categorised according to the number of H3O+ (or) OH– molecules generated for every mole of the chemical dissolved in H2O. In general, acids and bases are classified as either strong or weak. A weak acid partially dissociates in water, and a strong acid dissociates in water completely.

Ionisation Constants in Acidic Solutions

The equilibrium constants of acids in aqueous solutions can determine the relative strengths of different acids in a solution. For the same amount, stronger acids ionise more, producing more hydronium ions than weaker acids. The acid-ionisation constant, abbreviated Ka, is the acidic solution equilibrium constant for an acidic solution.

HA is used in the reactivity of an acid:

HA (aq) +H2O (l) ⇌ H3O+ (aq) +A− (aq)

The acid ionisation constant is denoted by:

Ka= [H3O+] [A−]/ [HA]

We do not add [H2O] in the formula because water acts as a reactant and serves as a solvent. Therefore we will not include it. The higher the Ka value of an acid, the higher the concentrations of H3O+ and A− are compared to the amount of the non-ionised acidic solution, HA. As a result, the ionisation constant of a stronger acid is higher than that of a weaker acid. The ionisation constant increases under the increasing strength of the acid.

An acid’s ionisation percentage is another way of evaluating its strength. The weak acid’s per cent ionisation is defined as the ratio of the ionised acid’s concentration to the initial acid’s concentration multiplied by 100:

Percent ionisation = [H3O+] eq/ [HA] 0 × 100%

Ionisation Constants in Basic Solutions

In aqueous solutions, the amount of base-ionisation constant (Kb) reflects the relative strength of the base. Stronger bases ionise more readily than weaker bases, resulting in larger hydroxide ion concentrations in similar solutions. The ionisation constant of a stronger base is greater than that of a weaker base.

B is used in the reactivity of a base:

B (aq) +H2O (l) ⇌ HB+ (aq) +OH− (aq)

The basic ionisation constant is denoted by:

Kb= [HB+] [OH−]/ [B]

The per cent ionisation of a base, a measure of its relative strength, can be calculated as:

Per cent ionisation = [OH−]eq/[B]0 × 100%

Conjugate Acid-Base Pairs’ Relative Strengths

The transport of protons in Bronsted-Lowry acid-base chemistry shows a correlation between the relative strengths of conjugate acid-base pairs; this is consistent with logic. An acid or base’s ionisation constant (Ka or Kb) measures how much the acid or base has been ionised.

Equilibrium constants for additional reactions are mathematically multiplied to get a summed reaction’s equilibrium constant.

Ka × Kb= [H3O+] [A−] / [HA] × [HA] [OH−] / [A−] = [H3O+] [OH−] = Kw

This equation expresses the relationship between the ionisation constants of any conjugate acid-base pairing, namely that their mathematical product equals the water ion product, Kw. Rearranging this equation reveals a reciprocal relationship between conjugate acid-base strengths:

Ka=Kw/Kb or Kb=Kw/Ka

This inverse proportional relationship means that the acid or base has a greater effect on the conjugate’s strength.

Weak Acids and Bases

Weak acids and bases’ hydronium or hydroxyl ionic strength is lower than their overall concentration.

Calculating the concentrations of hydronium or hydroxyl ions formed using weak acids and bases is difficult than calculating the concentrations produced by strong acids and bases. It involves leveraging the acid-base equilibrium constant to produce the hydronium or hydroxyl ion.

Conclusion

Strong acids and bases have almost complete ionisation; weak acids and bases have only partial ionisation. According to the dissociation constants of weak acids and bases, the relative strength of these substances can be determined. Temperature is the only factor that affects equilibrium constants. Higher the acid’s dissociation constant value, the stronger the acid. Similarly, the greater the base’s dissociation constant value, the stronger the base.