pH Calculator

How to use this calculator

Four input mode tabs let you calculate from whichever value you know. Select the tab, enter your value, and press Calculate.

From [H+]: Enter hydrogen ion concentration in mol/L. Scientific notation is supported: type 1e-7 for 1×10⁻⁷ mol/L (neutral water at 25°C).

From [OH-]: Enter hydroxide ion concentration in mol/L. The calculator computes [H+] via Kw = 1×10⁻¹⁴.

From pOH: Enter the pOH value. pH = 14 − pOH at 25°C.

From pH: Enter a known pH and the calculator returns [H+], [OH-], pOH, and the acid/base classification.

The Mathematical Definition

pH is defined as the negative base-10 logarithm of hydrogen ion concentration:

where [H⁺] is the molar concentration of hydrogen ions (mol/L). The negative sign ensures that low concentrations (acidic solutions with many H⁺ ions per liter) give high numbers in terms of acidity representation: lower pH means more acidic.

The base-10 logarithm means that each pH unit represents a tenfold change in H⁺ concentration. A solution at pH 3 has ten times more hydrogen ions than one at pH 4, and one hundred times more than one at pH 5. This logarithmic compression is what makes the 0-14 pH scale practical; the actual H⁺ concentrations range from 1 mol/L (pH 0) to 10⁻¹⁴ mol/L (pH 14), spanning 14 orders of magnitude.

Understanding Hydrogen Ion Concentration

In aqueous solution, water undergoes autoionization:

At 25°C, the ion product of water (Kw) is:

In pure water, [H⁺] = [OH⁻] = 1 × 10⁻⁷ mol/L, giving pH = 7.

When an acid is added, it donates H⁺ ions to the solution, increasing [H⁺] above 10⁻⁷ mol/L and pushing pH below 7. When a base is added, it accepts H⁺ ions (or donates OH⁻), reducing [H⁺] and raising pH above 7.

The relationship between pH and pOH:

where pOH = -log₁₀([OH⁻]). This holds because log₁₀(Kw) = log₁₀(10⁻¹⁴) = -14, and the negative logs of [H⁺] and [OH⁻] must sum to 14.

Strong vs Weak Acids: Complete vs Partial Dissociation

A fundamental distinction in acid chemistry is between strong and weak acids.

Strong acids dissociate completely in water. Every molecule releases its H⁺:

• Hydrochloric acid (HCl): HCl → H⁺ + Cl⁻
• Sulfuric acid (H₂SO₄): H₂SO₄ → 2H⁺ + SO₄²⁻ (first dissociation complete)
• Nitric acid (HNO₃): HNO₃ → H⁺ + NO₃⁻

For a strong acid, pH is simply -log₁₀(C) where C is the acid concentration. A 0.1 mol/L HCl solution has pH = -log₁₀(0.1) = 1.

Weak acids only partially dissociate. They reach an equilibrium between dissociated and undissociated forms, described by the acid dissociation constant Ka:

For acetic acid (Ka = 1.8 × 10⁻⁵), a 0.1 mol/L solution reaches pH approximately 2.87, which is much higher than the pH 1 of an equal concentration strong acid. Only about 1.3% of acetic acid molecules have dissociated. This partial dissociation is what makes weak acids less corrosive and why they buffer solutions.

Temperature Dependence of Neutral pH

One of the most frequently misunderstood aspects of pH is that neutral pH is only 7 at exactly 25°C. At other temperatures, the Kw changes, and so does the neutral pH point:

At body temperature (37°C), neutral pH is 6.81. Blood pH of 7.4 is therefore more basic relative to neutral at body temperature than it might appear compared to the 25°C reference point. Pure water at 37°C has pH about 6.8: it is not acidic, just neutral at that temperature.

This means pH meters require temperature compensation to give accurate readings. Most laboratory pH meters have automatic temperature compensation (ATC) built in.

Biological pH Regulation

The human body operates within tightly controlled pH ranges. Blood pH must stay between 7.35 and 7.45 for normal function. Deviations cause serious problems:

• Below 7.35: acidosis. Impairs enzyme function, reduces hemoglobin’s oxygen affinity, causes irregular heart rhythms. Severe acidosis below pH 7.0 can be fatal.
• Above 7.45: alkalosis. Causes muscle spasms, tingling, tetany. Severe alkalosis above pH 7.8 is also dangerous.

Blood is buffered primarily by the carbonate/bicarbonate system. CO₂ dissolves in blood to form carbonic acid (H₂CO₃), which equilibrates with bicarbonate (HCO₃⁻):

The lungs control CO₂ concentration (and therefore carbonic acid). The kidneys control bicarbonate concentration. Together they maintain pH within the narrow safe range.

Different compartments in the body have different pH values. Stomach gastric acid: pH 1.5-3.5 (hydrochloric acid secreted by parietal cells). Small intestine: pH 6-7.5. Lysosomal interior in cells: pH 4.5-5.0. Mitochondrial matrix: pH 7.5-8. Urine: pH 4.5-8.5 depending on diet and hydration.

Industrial pH Control

Water treatment: Drinking water is typically adjusted to pH 7.5-8.5 before distribution to minimize corrosion of pipes and distribution infrastructure. Low pH water dissolves lead and copper from plumbing. High pH reduces corrosion but can cause scaling. Lime (calcium hydroxide) or sodium hydroxide is used to raise pH; carbon dioxide or sulfuric acid is used to lower it.

Food preservation: pH below 4.6 inhibits the growth of Clostridium botulinum, the organism responsible for botulism. This is why acidified foods (pickles, canned tomatoes, hot sauces) have pH well below 4.6. The FDA defines a low-acid food as having equilibrium pH above 4.6.

Pharmaceutical manufacturing: Most drug molecules are either weak acids or weak bases. Their solubility, stability, and membrane permeability depend heavily on pH. Tablets are often buffered to specific pH ranges. Injectable formulations require pH close to physiological (7.4) to minimize irritation at injection sites.

Fermentation: Yeast for beer and wine fermentation works optimally at pH 4.0-4.5. Bacteria in yogurt production lower pH from about 6.5 in milk to below 4.6 as they produce lactic acid. pH control in bioreactors is critical for maintaining optimal enzyme and organism activity.

How pH Meters Work

A pH meter measures the electrical potential (voltage) of a solution using a glass electrode. The glass electrode contains an internal reference solution (often 0.1 mol/L HCl) separated from the test solution by a thin, specially formulated glass membrane. Hydrogen ions can migrate through this glass membrane, creating a potential difference proportional to the pH difference across it.

The Nernst equation governs this potential:

At 25°C, this simplifies to approximately 59.16 mV per pH unit (the Nernst slope). So a change of 1 pH unit corresponds to a 59.16 mV change in electrode potential. The meter electronics read this voltage and convert it to pH.

Calibration using two or three buffer solutions of known pH is essential. Single-point calibration (pH 7 buffer only) sets the offset; two-point calibration (usually pH 4 and 7, or 7 and 10) corrects both offset and slope. The slope should be between 95% and 105% of the theoretical 59.16 mV/pH for a functioning electrode.

Titration and pH Curves

Acid-base titrations track pH as one solution is gradually added to another. The resulting pH curve has a characteristic S-shape with a steep inflection point at the equivalence point, where moles of acid equal moles of base.

For a strong acid titrated with a strong base, the equivalence point is at pH 7. For a weak acid titrated with strong base, the equivalence point pH is above 7 (the conjugate base of a weak acid is itself weakly basic). The buffer region of the titration curve, roughly 1 pH unit above and below the half-equivalence point, is where the weak acid/conjugate base pair provides buffering capacity.

The Henderson-Hasselbalch equation describes pH in the buffer region:

where [A⁻] is the concentration of conjugate base and [HA] is the concentration of weak acid. When [A⁻] = [HA], the log term is zero and pH = pKa. This is the optimal buffer pH, the point at which the buffer has maximum capacity to resist change when acid or base is added.

Common laboratory buffers and their useful pH ranges: phosphate buffer (pH 5.8-8.0), HEPES (pH 6.8-8.2), TRIS (pH 7.0-9.0), acetate buffer (pH 3.6-5.6), citrate buffer (pH 3.0-6.2).

pH in industrial processes

Industrial processes depend on pH control to optimize reaction rates, prevent corrosion, ensure product quality, and meet discharge regulations.

Water treatment: Municipal water treatment adjusts pH to 7.5-8.5 during disinfection to balance microbial inactivation efficiency with minimizing disinfection byproduct formation. After chlorination, pH is re-adjusted. Lead and copper corrosion from distribution pipes is strongly pH-dependent: lower pH increases metal leaching.

Food production: Yogurt, cheese, and fermented beverages depend on pH for flavor and safety. Lactic acid fermentation lowers pH to 4-4.5, which inhibits pathogens. Canning processes lower pH below 4.6 (the threshold for Clostridium botulinum growth) with added acid.

Electroplating and surface finishing: Metal deposition baths operate at specific pH ranges. Nickel plating typically uses pH 3.5-4.5. Chrome plating operates in strongly acidic conditions (pH 0-2). pH control directly affects deposit quality and bath stability.

Pharmaceutical manufacturing: API synthesis reactions often require precise pH control for yield and selectivity. Enzyme-catalyzed reactions must be carried out at the enzyme’s optimal pH. Drug stability is pH-dependent: many drugs are stable only within a narrow pH range.

Paper and pulp: Cellulose hydrolysis rates are strongly pH-dependent. Acidic conditions favor certain bleaching reactions. Mill wastewater must be neutralized to pH 6-9 before discharge.

Temperature dependence of the pH scale

The pH scale is not fixed at 0-14 for all temperatures. The autoionization constant of water (Kw) is temperature-dependent, which means the neutral pH changes with temperature.

At body temperature (37°C), neutral pH is 6.81, not 7.00. Blood at pH 7.4 is slightly basic relative to the neutral point at body temperature.

This also means pH meters must be temperature-compensated. Modern meters include a temperature sensor and automatically adjust the neutral reference. Manual temperature compensation is required for precision measurements with basic pH equipment.