What Is pH?
pH is a dimensionless number that quantifies the acidity or basicity of an aqueous solution. The term "pH" stands for "potential of hydrogen" (or "power of hydrogen") and was first introduced by Danish chemist Søren Peder Lauritz Sørensen in 1909. It provides a convenient way to express hydrogen ion concentrations that would otherwise require cumbersome scientific notation.
The pH scale runs from 0 to 14 under standard conditions (25°C, 1 atm). A pH of 7 is considered neutral — this is the pH of pure water at room temperature. Values below 7 indicate acidic solutions where hydrogen ions [H⁺] dominate, while values above 7 indicate basic (alkaline) solutions where hydroxide ions [OH⁻] are more prevalent.
What makes pH particularly powerful is its logarithmic nature. Each whole number change in pH represents a tenfold (10×) change in hydrogen ion concentration. This means a solution at pH 3 is ten times more acidic than one at pH 4, and one hundred times more acidic than one at pH 5. This logarithmic compression allows us to express the enormous range of hydrogen ion concentrations found in nature — from concentrated acids to strong bases — on a manageable 0–14 scale.
The pH Formula and Related Equations
The fundamental equation for pH is deceptively simple but tremendously useful:
pH = −log₁₀[H⁺]Where [H⁺] is the molar concentration of hydrogen ions in mol/L (M). Related formulas include:
- pOH = −log₁₀[OH⁻] — the hydroxide equivalent
- pH + pOH = 14 — at 25°C (the water autoionization constant pKw)
- [H⁺] = 10^(−pH) — to find concentration from pH
- [OH⁻] = 10^(−pOH) — to find hydroxide concentration
- [H⁺] × [OH⁻] = Kw = 1×10⁻¹⁴ — the ion product of water at 25°C
For example, if you have a hydrochloric acid solution with [H⁺] = 0.001 M (1×10⁻³ M), then pH = −log(10⁻³) = 3. The corresponding pOH would be 14 − 3 = 11, and [OH⁻] = 10⁻¹¹ M.
For weak acids and bases, the Henderson-Hasselbalch equation is essential:
pH = pKa + log₁₀([A⁻]/[HA])This equation is invaluable for buffer calculations, where [A⁻] is the conjugate base concentration and [HA] is the weak acid concentration.
The pH Scale: Acids, Neutral, and Bases
Understanding where common substances fall on the pH scale provides practical context for chemistry, biology, and everyday life. Below is a comprehensive reference table of pH values for familiar substances:
| Substance | Approximate pH | Classification |
|---|---|---|
| Battery acid (H₂SO₄) | 0–1 | Strong acid |
| Gastric acid (stomach) | 1.5–3.5 | Strong acid |
| Lemon juice / Vinegar | 2.0–3.0 | Acid |
| Orange juice | 3.5–4.5 | Acid |
| Coffee | 4.5–5.5 | Weak acid |
| Milk | 6.5–6.8 | Slightly acidic |
| Pure water | 7.0 | Neutral |
| Human blood | 7.35–7.45 | Slightly basic |
| Baking soda | 8.3 | Weak base |
| Ammonia solution | 11.0 | Base |
| Bleach (NaClO) | 12.5 | Strong base |
| Drain cleaner (NaOH) | 13–14 | Strong base |
Buffers and Buffer Capacity
Buffer solutions are one of the most important practical applications of pH chemistry. A buffer resists changes in pH when small amounts of acid or base are added. This is critical in biological systems — for instance, human blood maintains a remarkably stable pH of 7.35–7.45 through the bicarbonate buffer system (H₂CO₃/HCO₃⁻).
A buffer typically consists of a weak acid paired with its conjugate base (like acetic acid + sodium acetate) or a weak base paired with its conjugate acid (like ammonia + ammonium chloride). The buffer works because the weak acid can neutralize added bases, while the conjugate base neutralizes added acids.
Buffer capacity depends on two factors: the total concentration of the buffer components and how close the pH is to the pKa of the weak acid. Maximum buffering occurs when pH = pKa, where the concentrations of acid and conjugate base are equal. As a rule of thumb, a buffer is effective within ±1 pH unit of its pKa value.
pH Indicators: Visual Chemistry
pH indicators are substances that change color at specific pH ranges, providing a visual method to estimate the acidity of a solution. They are themselves weak acids or bases whose protonated and deprotonated forms have different colors.
| Indicator | pH Range | Color Change |
|---|---|---|
| Methyl violet | 0.0–1.6 | Yellow → Blue |
| Methyl orange | 3.1–4.4 | Red → Yellow |
| Litmus | 4.5–8.3 | Red → Blue |
| Bromothymol blue | 6.0–7.6 | Yellow → Blue |
| Phenolphthalein | 8.2–10.0 | Colorless → Pink |
| Alizarin yellow | 10.1–12.0 | Yellow → Red |
Universal indicator is a mixture of several indicators that produces a continuous spectrum of colors across the entire pH range, making it ideal for quick pH estimation in educational and laboratory settings.
Real-World Applications of pH
Water Treatment
Municipal water treatment plants carefully control pH between 6.5 and 8.5 to prevent pipe corrosion, ensure disinfection effectiveness, and meet drinking water standards. Acidic water dissolves lead and copper from pipes, while overly basic water causes scale buildup.
Agriculture & Soil Science
Soil pH affects nutrient availability for plants. Most crops thrive in pH 6.0–7.0. Blueberries prefer acidic soil (pH 4.5–5.5), while alfalfa needs slightly alkaline conditions (pH 6.8–7.5). Farmers use lime to raise pH or sulfur to lower it.
Medicine & Biology
Blood pH is maintained at 7.35–7.45 through respiratory and renal buffering. Deviation outside this range (acidosis or alkalosis) can be life-threatening. Enzyme activity is pH-dependent — pepsin works optimally at pH 2 in the stomach, while trypsin prefers pH 8 in the intestine.
Food & Beverage Industry
pH control is essential for food safety and taste. Canned foods must be below pH 4.6 to prevent botulism. Cheese making, fermentation, and bread baking all depend on precise pH management. Wine quality is influenced by must pH during fermentation.
Strong vs. Weak Acids and Bases
Strong acids (HCl, H₂SO₄, HNO₃) completely dissociate in water, meaning their pH can be calculated directly from concentration: a 0.01 M HCl solution has pH = 2. Strong bases (NaOH, KOH) similarly fully dissociate.
Weak acids (CH₃COOH, H₂CO₃) and weak bases (NH₃) only partially dissociate, requiring equilibrium calculations involving their Ka or Kb values. For a weak acid with concentration C and dissociation constant Ka, the pH is approximately:
pH ≈ ½(pKa − log C)This approximation holds when the degree of dissociation is small (less than ~5%), which is usually the case for weak acids at moderate concentrations.
Frequently Asked Questions
What is pH?
pH is a logarithmic scale measuring the acidity or basicity of an aqueous solution, ranging from 0 (most acidic) to 14 (most basic). It is defined as pH = −log₁₀[H⁺], where [H⁺] is the hydrogen ion concentration in moles per liter.
How do I calculate pH from hydrogen ion concentration?
Use pH = −log₁₀[H⁺]. For example, if [H⁺] = 1×10⁻⁵ M, then pH = −log(10⁻⁵) = 5. Our calculator handles this computation instantly for any concentration you provide.
What is the relationship between pH and pOH?
At 25°C, pH + pOH always equals 14. This comes from the water autoionization constant Kw = [H⁺][OH⁻] = 10⁻¹⁴. Knowing one value immediately gives you the other.
What pH values indicate acids, neutral, and bases?
pH < 7 is acidic, pH = 7 is neutral, and pH > 7 is basic (alkaline). Each unit represents a 10× change in H⁺ concentration, so pH 3 is 10,000 times more acidic than pH 7.
What is a buffer solution?
A buffer resists pH changes when small amounts of acid or base are added. It consists of a weak acid and its conjugate base (or vice versa). Blood uses the bicarbonate buffer (H₂CO₃/HCO₃⁻) to stay at pH 7.4.
How do pH indicators work?
pH indicators are weak acids/bases whose protonated and deprotonated forms have different colors. They change color over a specific pH range. Litmus turns red in acid and blue in base; phenolphthalein turns pink above pH 8.2.
What is the pH of common substances?
Battery acid: ~0, lemon juice: ~2, vinegar: ~2.5, coffee: ~5, pure water: 7, blood: 7.4, baking soda: 8.3, ammonia: 11, bleach: 12.5, drain cleaner: ~14.
Is this pH calculator free to use?
Yes! Our pH calculator is completely free, requires no registration, and runs entirely in your browser. Enter any value and get instant pH, pOH, [H⁺], and [OH⁻] calculations.