pH is a logarithmic measure related to the activity of hydrogen ions in a solution and is commonly expressed approximately as the negative base ten logarithm of hydrogen ion activity. In dilute aqueous solutions, this is often simplified to pH equals negative logarithm base ten of the hydrogen ion concentration, although the rigorous thermodynamic definition uses activity rather than concentration. The pH scale is used to describe how acidic or alkaline an aqueous solution is.
A lower pH corresponds to greater hydrogen ion activity, while a higher pH corresponds to lower hydrogen ion activity under the same general conditions. Because the scale is logarithmic, a change of one pH unit represents a tenfold change in hydrogen ion activity in idealised dilute conditions. Thus pH three is not merely slightly more acidic than pH four; its hydrogen ion activity is approximately ten times greater. The concept is especially useful because hydrogen ion concentrations can span many orders of magnitude in chemical and biological systems.
The mathematical definition of pH is pH equals negative log ten of a H plus, where a H plus represents the activity of hydrogen ions. Activity incorporates nonideal interactions between ions and solvent and is related to concentration through an activity coefficient. In dilute solutions, activity coefficients can approach one, making concentration a useful approximation. In concentrated solutions, however, the distinction becomes important. A free hydrogen ion does not exist as an isolated bare proton in ordinary liquid water.
It is strongly solvated, with hydronium and larger hydrated proton structures participating in proton transfer. The notation H plus is nevertheless retained because it provides a convenient thermodynamic representation. pH is therefore not literally a direct count of isolated free protons. It is a thermodynamic quantity associated with proton activity. Standard pH measurements depend on reference conventions and electrochemical measurement principles, which become increasingly important when high accuracy is required.
The pH of a solution depends on its chemical composition and equilibrium processes. Strong acids can produce high hydrogen ion activities when sufficiently concentrated because they dissociate extensively in water. Weak acids establish equilibria between protonated and deprotonated forms, so their hydrogen ion activity depends on both acid concentration and acid dissociation constant. Bases generally reduce hydrogen ion activity by accepting protons or increasing hydroxide concentration.
Water itself undergoes autoionisation according to the equilibrium two water molecules reversible reaction hydronium plus hydroxide. At a specified temperature, the product of hydrogen ion and hydroxide ion activities is related to the ionic product of water, Kw. This relationship connects pH and pOH. In dilute aqueous systems at approximately room temperature, pH seven is often used as a convenient reference for neutrality, but the exact neutral pH changes with temperature because Kw changes with temperature.
Because pH is logarithmic, numerical differences have multiplicative chemical meaning. If one idealised aqueous solution has pH three and another has pH five, the first has approximately one hundred times the hydrogen ion activity of the second under comparable conditions. This logarithmic representation makes pH convenient for laboratory work because hydrogen ion activity can vary over many powers of ten. However, the pH scale should not be treated as an absolute scale extending without qualification from zero to fourteen.
Extremely concentrated acids or bases can have conventional pH values outside the familiar zero to fourteen range, and the meaningful range depends on the solvent, activity convention and measurement conditions. In nonaqueous solvents, alternative acidity scales and definitions may be more appropriate. pH is therefore most naturally understood as a thermodynamic quantity associated with a particular solvent system rather than as a universal dimensionless ruler independent of context.
pH is measured using several techniques. Indicators are dyes whose colours change over particular pH ranges because their protonated and deprotonated forms have different electronic structures. Litmus, methyl orange, phenolphthalein and universal indicator are common examples, although each has a limited transition range. A pH meter generally uses a glass electrode and a reference electrode to measure an electrochemical potential related to hydrogen ion activity.
The electrode response is described approximately by the Nernst equation and depends on temperature, which is why accurate pH meters often include temperature compensation. Calibration is normally performed using standard buffer solutions with known reference pH values. Measurement accuracy can be affected by temperature, ionic strength, electrode condition, contamination, junction potentials and calibration quality. Consequently, a pH value is a measured thermodynamic quantity rather than a number that can always be obtained accurately from a simple colour comparison.
pH has enormous practical importance in biology, medicine, environmental science, agriculture, food chemistry and industry. Enzymes often operate effectively only within particular pH ranges because protonation states influence protein structure and catalytic activity. Human blood is maintained within a narrow physiological pH range because large deviations can disrupt biochemical reactions and cellular function. Soil pH influences nutrient availability, mineral solubility and microbial activity. Freshwater and marine pH affect carbonate equilibria and aquatic organisms.
Food production uses pH to control flavour, preservation, fermentation and microbial growth. Industrial processes monitor pH to control corrosion, precipitation, chemical reactions and wastewater treatment. In each case, pH is useful because proton activity influences the equilibrium and reaction behaviour of many chemical systems. The practical significance of pH therefore extends far beyond the simple classification of substances as acidic or alkaline.
pH must be distinguished from several related quantities. Acid concentration describes how much acid substance is present, whereas pH describes hydrogen ion activity in the resulting solution. Acid strength describes the thermodynamic tendency of an acid to donate protons and is commonly characterised by Ka or pKa.
A dilute strong acid can have a higher pH than a concentrated weak acid, depending on their concentrations and equilibria. pOH measures hydroxide ion activity and is related to pH through the ionic product of water under specified conditions. Buffer capacity describes the ability of a solution to resist pH change when acid or base is added and is not the same thing as pH. Alkalinity measures the capacity of a solution to neutralise acid and is especially important in natural water chemistry. These distinctions prevent the common error of treating pH as a direct universal measure of how much acid or base a substance contains.
The modern concept of pH was introduced by the Danish chemist Soren Sorensen in 1909 while working on biochemical measurements. The notation is commonly associated with the expression power of hydrogen, although the exact historical interpretation of the letter p and the symbol H has been discussed in the history of chemistry. The concept provided a convenient logarithmic representation of hydrogen ion concentration for analytical and biochemical work.
Modern physical chemistry has refined the definition so that activity, reference states and measurement conventions are central to accurate pH determination. The scale is therefore more rigorous than the simple school level rule that acids have pH below seven and bases have pH above seven. That rule is useful for ordinary dilute aqueous solutions near room temperature, but pH depends on temperature, solvent and thermodynamic conventions. A scientifically precise understanding treats pH as a logarithmic measure of proton activity and recognises that neutrality and numerical pH seven are not universally identical under all conditions.
