The concentration of hydronium ions in an aqueous solution is usually written with negative exponents and can be as small as 1 × 10⁻¹⁴ M. The pH scale was developed by chemist Soren Sorenson in 1909 as a more convenient way to quickly compare the acidity of different solutions.

The pH of a solution is the negative logarithm of its hydronium ion concentration. For example, a solution with 1 × 10⁻⁵ M hydronium concentration has a pH of 5. The greater the hydronium ion concentration of a solution, the lower its pH.

As pH is expressed on a logarithmic scale, a single-unit change corresponds to a 10-fold increase or decrease of the hydronium ion concentration. A solution with a pH of 3 will have ten times more hydronium ions than a solution with a pH 4 and one hundred times more than a solution with a pH of 5.

An acidic solution has a higher concentration of hydronium ions than hydroxide ions and a pH less than 7, whereas a basic solution has a lower concentration of hydronium ions than hydroxide ions and a pH greater than 7.

A neutral solution with an equal concentration of hydronium and hydroxide ions has a pH of 7.

The concentration of hydroxide ions can also be expressed as pOH. pOH is the negative logarithm of the hydroxide ion concentration. The higher the hydroxide ion concentration, the lower its pOH value.

A pH or pOH value of an aqueous solution ranges from 0 – 14. This is because K_W, the equilibrium constant for the autoionization of water, is equal to 1 × 10⁻¹⁴. 

Taking the negative log of both sides of the equation results in an equation where pK_W is equal to the sum of pH and pOH. Since the negative log of 1 × 10⁻¹⁴  is 14, the sum of the pH and pOH of an aqueous solution will always be 14. This can be used to determine the pH value when the pOH value is known and vice versa.

For example, a solution with a pOH of 10 has a pH of 4.