pH and Acid–Base Chemistry
The pH scale, the equilibrium chemistry of acids and bases, and the role of buffers in maintaining stable conditions.
pH is a measure of the acidity of a solution, defined as the negative base-10 logarithm of the hydrogen ion activity, conventionally written as concentration:
pH = −log10[H+]
Pure water at 25 °C has a hydrogen ion concentration of 10−7 mol/L, giving pH 7. Because the product of hydrogen and hydroxide concentrations is constant (Kw = 10−14 at 25 °C), pH 7 is neutral: solutions with pH below 7 are acidic, and solutions above 7 are basic (alkaline). The scale is logarithmic, so a solution at pH 3 is ten times more acidic than one at pH 4 and a hundred times more acidic than one at pH 5.
Acids and bases are understood through proton transfer (Brønsted–Lowry theory). An acid is a proton donor and a base is a proton acceptor; when an acid donates a proton it leaves behind its conjugate base, and when a base accepts one it forms its conjugate acid. Strong acids such as HCl dissociate completely in water, so their pH follows directly from concentration. Weak acids such as acetic acid dissociate only partially; the equilibrium is characterized by the acid dissociation constant Ka, and the pH of a weak acid solution depends on both Ka and concentration through the Henderson–Hasselbalch equation:
pH = pKa + log10([A−]/[HA])
This equation also describes buffers — solutions containing a weak acid and its conjugate base (or a weak base and its conjugate acid). A buffer resists pH change when small amounts of acid or base are added, because the added protons are absorbed by the conjugate base or released by the acid. Blood is buffered near pH 7.4 by the carbonic acid/bicarbonate system, and most enzymes work within narrow pH ranges; deviations of a few tenths of a pH unit can be life-threatening.
Neutralization reactions between acids and bases produce water and a salt, and titration uses measured neutralization to determine unknown concentrations. pH is measured with indicators (color changes at characteristic pH values), with pH paper, or most precisely with a glass electrode pH meter. The concept was introduced by the Danish chemist Søren Sørensen in 1909 and has become one of the most used quantities in chemistry, biology, medicine, agriculture, and environmental science.