pH and pOH Guide

Understand pH, pOH, and the pH scale. Learn how to calculate pH from hydrogen ion concentration, the relationship between pH and pOH, buffers, and real-world applications.

What Is pH?

pH is a logarithmic scale that measures the acidity or basicity of an aqueous solution. It is defined as pH = -log10[H+], where [H+] is the molar concentration of hydrogen ions (H+, more precisely hydronium ions H3O+) in the solution. The pH scale typically ranges from 0 to 14 for most common solutions. A pH of 7 is neutral (pure water at 25 degrees Celsius), values below 7 are acidic, and values above 7 are basic (alkaline). Each whole number change in pH represents a tenfold change in hydrogen ion concentration: a solution with pH 3 is ten times more acidic than one with pH 4, and one hundred times more acidic than one with pH 5.

What Is pOH?

pOH is the analogous measure for hydroxide ion (OH-) concentration: pOH = -log10[OH-]. In any aqueous solution at 25 degrees Celsius, the product of hydrogen ion and hydroxide ion concentrations is a constant: [H+] * [OH-] = 1.0 * 10^-14 (the ion product of water, Kw). Taking the negative logarithm of both sides gives the fundamental relationship pH + pOH = 14. This means that if you know either pH or pOH, you can immediately find the other. A strongly acidic solution has a low pH and a high pOH, while a strongly basic solution has a high pH and a low pOH. At neutrality, pH = pOH = 7.

Calculating pH from Ion Concentration

To find the pH of a solution, you need to know the hydrogen ion concentration [H+]. For a 0.001 M HCl solution (a strong acid that dissociates completely), [H+] = 0.001 = 10^-3, so pH = -log10(10^-3) = 3. For a 0.01 M NaOH solution (a strong base), [OH-] = 0.01, pOH = -log10(0.01) = 2, and pH = 14 - 2 = 12. For weak acids and bases that only partially dissociate, you need to use the equilibrium constant (Ka for acids, Kb for bases) to find the actual [H+] or [OH-] at equilibrium. The ICE table method (Initial, Change, Equilibrium) is the standard approach for these calculations.

Strong vs. Weak Acids and Bases

Strong acids (HCl, HNO3, H2SO4, HBr, HI, HClO4) dissociate completely in water, so the [H+] equals the acid concentration. Strong bases (NaOH, KOH, Ca(OH)2, Ba(OH)2) also dissociate completely. Weak acids (acetic acid, citric acid, carbonic acid) and weak bases (ammonia, methylamine) only partially dissociate, establishing an equilibrium between the undissociated form and the ions. The acid dissociation constant Ka quantifies the extent of dissociation: larger Ka values indicate stronger weak acids. For a weak acid HA with concentration C and dissociation constant Ka, the [H+] at equilibrium is approximately sqrt(Ka * C) when the percent dissociation is small (less than about 5%).

The pH Scale in Context

Everyday substances span the entire pH scale. Battery acid has a pH around 0-1. Stomach acid (gastric juice) has a pH of about 1.5-3.5. Lemon juice is around pH 2, vinegar around pH 2.5, and coffee around pH 5. Pure water is pH 7. Human blood is tightly regulated at pH 7.35-7.45. Baking soda solution is about pH 8.5, milk of magnesia about pH 10.5, household ammonia about pH 11.5, and drain cleaner (lye) about pH 13-14. Understanding where common substances fall on the pH scale helps build intuition for acidity and basicity, and illustrates why pH regulation is critical in biological systems, food science, and environmental monitoring.

Buffers: Resisting pH Change

A buffer is a solution that resists changes in pH when small amounts of acid or base are added. Buffers are made from a weak acid and its conjugate base (or a weak base and its conjugate acid). The Henderson-Hasselbalch equation describes buffer pH: pH = pKa + log10([A-]/[HA]), where pKa = -log10(Ka), [A-] is the concentration of the conjugate base, and [HA] is the concentration of the weak acid. Buffers are most effective when pH is within one unit of pKa, and when the concentrations of the weak acid and conjugate base are both substantial. Biological systems rely heavily on buffers; for example, the carbonic acid-bicarbonate buffer system maintains blood pH near 7.4.

Measuring pH

pH can be measured using several methods. pH indicator paper (litmus paper or universal indicator strips) changes color in response to pH and provides a rough estimate. Liquid pH indicators, such as phenolphthalein or bromothymol blue, change color over specific pH ranges and are used in titrations to detect the equivalence point. For accurate, quantitative measurements, a digital pH meter with a glass electrode is the standard tool. pH meters should be calibrated regularly using buffer solutions of known pH (typically pH 4, 7, and 10). Temperature affects pH readings because the ion product of water (Kw) changes with temperature, so precise measurements should note the temperature.

Applications of pH in Science and Industry

pH control is critical in a vast number of fields. In water treatment, pH is adjusted to optimize disinfection and prevent pipe corrosion. In agriculture, soil pH determines nutrient availability and crop suitability; most crops grow best at pH 6.0-7.0. In food science, pH affects taste, preservation, and safety; low pH inhibits bacterial growth, which is why pickling in vinegar is an effective preservation method. In medicine, blood pH outside the range of 7.35-7.45 constitutes a medical emergency (acidosis or alkalosis). In industrial chemistry, pH controls reaction rates, product purity, and catalyst activity. Swimming pool water is maintained at pH 7.2-7.8 to maximize chlorine disinfection effectiveness and minimize skin and eye irritation.

Try These Calculators

Put what you learned into practice with these free calculators.

pH Calculator pOH Calculator Buffer pH Calculator