Water has a unique, bent structure with an uneven charge distribution, making it polar. This polarity allows water molecules to interact strongly, leading to dissociation, where a small fraction naturally breaks apart. This process establishes a constant, dynamic equilibrium that forms electrically charged particles, or ions, even in pure water.
The Ions That Water Creates
When a water molecule dissociates, it forms two specific ions: the Hydronium ion and the Hydroxide ion. One water molecule transfers a proton, which is a positively charged hydrogen nucleus, to a neighboring water molecule. This transfer is what forms the two products of the dissociation.
The molecule that accepts the proton becomes the positively charged Hydronium ion (\(\text{H}_3\text{O}^+\)). Although the simplified reaction is often shown as the formation of a bare Hydrogen ion (\(\text{H}^+\)), this proton is never found free in an aqueous solution. Instead, it immediately associates with a water molecule to create the Hydronium ion, the stable positive product.
The molecule that loses the proton becomes the negatively charged Hydroxide ion, which has the chemical formula \(\text{OH}^-\). This ion consists of an oxygen atom bonded to a hydrogen atom and carries a single negative charge. The Hydroxide ion is an important component in chemical reactions and plays a significant role in determining a solution’s basicity.
The Reversible Process of Self-Ionization
The process by which water molecules spontaneously form these ions is called self-ionization or autoionization. This is a reversible reaction where two water molecules interact, with one acting as a proton donor and the other as a proton acceptor. The simplified chemical equation for this dynamic equilibrium is \(\text{2H}_2\text{O} \rightleftharpoons \text{H}_3\text{O}^+ + \text{OH}^-\).
Water molecules are simultaneously breaking apart into ions and reforming back into neutral water molecules. The vast majority of water molecules remain intact, as only about two molecules in every billion dissociate into ions. This slight ionization is quantified by the ion-product constant for water, known as \(K_w\).
At a standard temperature of 25 degrees Celsius, the value of \(K_w\) is \(1.0 \times 10^{-14}\). This extremely small constant confirms that the equilibrium lies heavily on the side of the intact water molecules. In pure water, the concentration of the Hydronium ions and the Hydroxide ions must be equal, meaning both are present at a concentration of \(1.0 \times 10^{-7}\) moles per liter.
The Role of Dissociated Ions in pH
The concentrations of Hydronium and Hydroxide ions created by water’s self-ionization form the baseline for the \(\text{pH}\) scale. The term \(\text{pH}\) is a measure of the concentration of the Hydronium ion (\(\text{H}_3\text{O}^+\)) in a solution. Mathematically, \(\text{pH}\) is the negative logarithm of the Hydronium ion concentration, making it a convenient scale to express very small numbers.
In pure water, where the Hydronium and Hydroxide concentrations are equal at \(10^{-7}\) moles per liter, the \(\text{pH}\) calculates to 7, which is defined as a neutral solution. When an acid is introduced to water, it increases the concentration of \(\text{H}_3\text{O}^+\) ions, pushing the \(\text{pH}\) below 7 and making the solution acidic. Conversely, when a base is added, it increases the concentration of \(\text{OH}^-\) ions.
Because the product of the two ion concentrations, \(K_w\), must remain constant, an increase in Hydroxide ions forces the Hydronium ion concentration to decrease. This decrease in \(\text{H}_3\text{O}^+\) concentration results in a \(\text{pH}\) greater than 7, which defines a basic or alkaline solution.