pH and H⁺ Concentration
Last reviewed: January 2026
Convert between pH values and hydrogen ion concentration [H⁺]. Calculate pOH and hydroxide concentration. This calculator runs entirely in your browser — your data stays private, and no account is required.
pH = −log₁₀[H⁺], where [H⁺] is the molar concentration of hydrogen ions. The scale runs 0–14. pH 7 = neutral (pure water at 25°C). Below 7 = acidic (excess H⁺ ions). Above 7 = basic or alkaline (excess OH⁻ ions). The scale is logarithmic: each whole number change represents a 10-fold change in acidity. pH 5 is 10× more acidic than pH 6, and 100× more acidic than pH 7. Reference values: battery acid ~1, stomach acid 1.5–3.5, lemon juice ~2.5, black coffee ~5, blood 7.35–7.45, bleach ~12.
Swimming pools: maintain pH 7.2–7.6. Below this, chlorine becomes irritating; above 8, chlorine loses effectiveness rapidly. Soil pH for plants: most plants thrive at 6.0–7.0. Below 5.5, aluminum and manganese become toxic to roots; above 7.5, iron and manganese become unavailable. Fermentation: beer at pH 4.2–4.6, wine 3.0–3.5, sourdough bread 3.5–4.5. Human blood has the tightest regulation: deviations beyond ±0.4 pH units from 7.4 are life-threatening.
| pH | H⁺ Concentration | Common Substance | Category |
|---|---|---|---|
| 1 | 10⁻¹ M | Battery acid | Strong acid |
| 3 | 10⁻³ M | Vinegar, citrus juice | Acid |
| 5 | 10⁻⁵ M | Black coffee | Weak acid |
| 7 | 10⁻⁷ M | Pure water | Neutral |
| 9 | 10⁻⁹ M | Baking soda solution | Weak base |
| 13 | 10⁻¹³ M | Bleach, oven cleaner | Strong base |
The pH scale measures how acidic or basic (alkaline) a solution is, running from 0 (most acidic) to 14 (most basic), with 7 representing neutrality (pure water at 25°C). What many people do not realize is that pH is a logarithmic scale — each whole number change represents a tenfold change in hydrogen ion (H⁺) concentration. A solution with pH 4 is ten times more acidic than pH 5, one hundred times more acidic than pH 6, and one thousand times more acidic than pH 7. This logarithmic nature means that seemingly small pH changes can represent enormous differences in chemical activity, which is why precise pH measurement and control are critical in many applications.
The term "pH" was introduced by Danish chemist Søren Peder Lauritz Sørensen in 1909 while working at the Carlsberg Laboratory (yes, the beer company). Mathematically, pH = -log₁₀[H⁺], where [H⁺] is the molar concentration of hydrogen ions in solution. Pure water undergoes a slight self-ionization: H₂O ⇌ H⁺ + OH⁻, producing equal concentrations of H⁺ and OH⁻ ions (each at 10⁻⁷ M at 25°C), giving a pH of 7. Adding an acid increases [H⁺] and decreases pH, while adding a base decreases [H⁺] (by consuming H⁺ ions or producing OH⁻ ions) and increases pH.
pH plays a critical role in numerous everyday situations that most people never consider. Human blood is maintained within a remarkably narrow pH range of 7.35-7.45 — even slight deviations can be life-threatening. Blood pH below 7.35 (acidosis) impairs enzyme function, cardiac output, and consciousness; above 7.45 (alkalosis) causes muscle spasms, nerve irritability, and cardiac arrhythmias. The body uses multiple buffer systems (primarily the bicarbonate system: CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻), respiratory rate adjustments, and kidney function to maintain this critical balance.
Stomach acid has a pH of 1.5-3.5, strong enough to dissolve metals, yet the stomach lining protects itself with a mucus barrier and bicarbonate secretion. Skin's acid mantle maintains a pH of 4.5-5.5, which inhibits bacterial and fungal growth — this is why soaps with high pH (9-10) can disrupt skin health by stripping this protective barrier. Swimming pool water is maintained at pH 7.2-7.6, the range where chlorine disinfection is most effective and swimmer comfort is optimized. Drinking water typically ranges from pH 6.5-8.5, with most municipal water systems targeting 7.0-7.5.
Soil pH is one of the most important factors determining which plants will thrive in a given location and whether nutrients are available for plant uptake. Most garden plants grow best in slightly acidic to neutral soil (pH 6.0-7.0), but optimal ranges vary by species. Blueberries, azaleas, and rhododendrons prefer acidic soils (pH 4.5-5.5), while lavender, clematis, and asparagus prefer slightly alkaline conditions (pH 7.0-7.5). The reason pH matters so much for plant growth is nutrient availability — at very low pH, aluminum and manganese become toxic to plants; at very high pH, iron, zinc, and phosphorus become chemically locked in insoluble forms that plants cannot absorb.
Soil pH can be modified through amendments. Lime (calcium carbonate) raises pH by neutralizing soil acidity — the amount needed depends on both the current pH and the soil's buffering capacity (clay and organic-matter-rich soils require more lime than sandy soils). Sulfur and aluminum sulfate lower pH for acid-loving plants. These amendments work gradually over weeks to months and should be applied based on soil test results rather than guesswork. For related calculations, see our Soil Amendment Calculator.
Industrial processes depend heavily on pH control. Water treatment plants adjust pH at multiple stages — coagulation works best at pH 5.5-7.0, disinfection efficiency varies with pH, and final distribution water must meet pH standards to prevent pipe corrosion (low pH) or scale formation (high pH). Brewing and winemaking require precise pH management for optimal fermentation, flavor development, and microbial stability. Pharmaceutical manufacturing controls pH to ensure drug stability, solubility, and bioavailability — many drug compounds change form at different pH values, affecting whether they dissolve in the stomach (pH 1.5-3.5) or intestine (pH 6-7.4).
Environmental monitoring uses pH as a key indicator of ecosystem health. Acid rain (pH below 5.6, the natural pH of rain in equilibrium with atmospheric CO₂) has caused widespread damage to lakes, forests, and buildings, particularly in regions downwind of industrial sulfur and nitrogen oxide emissions. Healthy lakes typically have pH 6.5-8.5; as pH drops below 6.0, sensitive aquatic species like trout begin to decline, and below pH 5.0, most fish species cannot survive. Ocean acidification — the ongoing decrease in ocean pH caused by absorption of atmospheric CO₂ — has reduced average ocean surface pH from approximately 8.2 to 8.1 since pre-industrial times, a 26% increase in acidity that threatens coral reefs, shellfish, and marine food chains. For related chemistry tools, try our Molar Mass Calculator and Chemical Equation Balancer.
pH can be measured using several methods with varying levels of accuracy. pH indicator papers (litmus paper and universal indicator strips) change color in response to pH and provide quick, approximate measurements (typically accurate to ±0.5-1.0 pH units). Liquid pH indicators like phenolphthalein, methyl orange, and bromothymol blue change color at specific pH ranges and are used in titration procedures. Electronic pH meters, the gold standard for accurate measurement, use a glass electrode that generates a voltage proportional to the hydrogen ion concentration. Quality benchtop pH meters provide accuracy of ±0.01 pH units, while portable field meters typically achieve ±0.02-0.05 units. Proper pH meter use requires regular calibration with standard buffer solutions (typically pH 4.00, 7.00, and 10.00), as electrode response drifts over time due to aging, contamination, and temperature changes. Temperature compensation is essential because pH values shift with temperature — even pure water's pH changes from 7.47 at 0°C to 6.14 at 100°C due to changes in the water ionization equilibrium constant.
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See also: Molar Mass Calculator · Pool Chemical Calculator · Density Calculator