Ionization of Water

What is Ionization?

Ionisation is defined as the process by which an atom or molecule gains or loses a positive or negative charge as a result of chemical changes. An ion is an electrically charged atom or molecule that results. If the ion has a negative charge, it is called an anion; if it has a positive charge, it is called a cation.

The basic ionisation reaction can be represented as follows:

M → M⁺ + e^–

Ionisation can occur as a result of the loss of an electron in collisions with subatomic particles, collisions with other atoms, molecules, and ions, or interactions with electromagnetic radiation.

Table of Contents

• Amphiprotic Nature of Water
• Self Ionization of Water
• Pure Water’s Ion-Product Constant
• Relationship among pH, pOH, and pKw
• Frequently Asked Questions – FAQs

Amphiprotic Nature of Water

Due to its highly polar structure, liquid water can either act as an acid (by donating a proton to a base) or a base (by using a lone pair of electrons to accept a proton).

As bases:

When a strong acid like HCl dissolves in water, it separates into chloride ions (Cl^–) and protons (H⁺). In turn, the proton reacts with a water molecule to form the hydronium ion (H₃O⁺):

HCl (aq) + H₂O(l) → H₃O⁺(aq) + Cl^–(aq)

The acid in this reaction is HCl, and the base is water, which accepts an H⁺ ion.

As acids:

Water can also act as an acid. H₂O donates a proton to NH₃, which acts as a base, in this equilibrium reaction:

H₂O(aq) + NH₃(aq) ⇆ (NH₄)⁺ + OH^–

Water is thus called amphiprotic because it can act as an acid or a base depending on the nature of the other reactant.

Self Ionization of Water

Water will self-ionize to a very small extent under normal conditions. The reaction in which a water molecule donates one of its protons to a neighbouring water molecule, either in pure water or in an aqueous solution, is called the self-ionization of water.

In an autoionization process, one water molecule can react with another to form an OH^– ion and an H₃O⁺ ion:

2H₂O(l) ⇆ H₃O⁺(aq) + OH^–(aq)

Pure Water’s Ion-Product Constant

This reaction’s equilibrium constant K can be written as follows:

As a result, the number of dissociated water molecules is very small.

The autoionization reaction’s equilibrium is far to the left, with few water molecules dissociate. As a result, the autoionization reaction leaves [H₂O] essentially unchanged and can be treated as a constant. By incorporating this constant into the equilibrium expression, we can rearrange the equilibrium constant equation to define a new equilibrium constant, the liquid water ion-product constant (K_w):

The concentrations of the hydronium ion and the hydroxide ion are equal when pure liquid water is in equilibrium with hydronium and hydroxide ions at 25°C:

H₃O⁺(aq) = OH^–(aq) = 1.003 × 10^-7 mol/L

Therefore,

At 25°C, the equilibrium constant K_w applies to any aqueous solution, not just pure water.

• The hydronium ion and the hydroxide ion concentrations are equal in pure water, so the solution is neutral.
• If [H₃O⁺] > [OH^–], the solution is acidic.
• If [H₃O⁺] < [OH^–], the solution is basic.

The H₃O⁺ concentration in an aqueous solution is a quantitative measure of acidity: the higher the H₃O⁺ concentration, the more acidic the solution. In contrast, the higher the concentration of OH^–, the more basic the solution.

Relationship among pH, pOH, and pK_w

The pH scale is a concise way of describing the concentration of H₃O⁺ in a solution, and thus its acidity or basicity.

The equilibrium constant K_w is expressed as:

Taking the negative logarithm of both sides of the equation:

pK_w = pH + pOH

Since, pK_w = 14

pH + pOH = 14.

At 25°C,

pH + pOH = 14.00 for any neutral solution with pH = pOH = 7.