Calculate pH, pOH, [H⁺], and [OH⁻] from any one value. Includes strong/weak acid-base calculations, dilution, and a visual pH scale with common substance comparisons.
The pH scale quantifies how acidic or basic an aqueous solution is. Defined as pH = -log₁₀[H⁺], the scale typically runs from 0 to 14 at 25 °C, with 7 being neutral. Every one-unit change in pH represents a tenfold change in hydrogen ion concentration.
Understanding pH is essential in chemistry, biology, medicine, environmental science, and everyday life. Blood pH is tightly regulated between 7.35 and 7.45; swimming pools are maintained at 7.2–7.8; soil pH affects nutrient availability for plants. The four interrelated values — pH, pOH, [H⁺], and [OH⁻] — are all connected through the water autoionization constant Kw = 1.0 × 10⁻¹⁴ at 25 °C.
This calculator converts between pH, pOH, [H⁺], and [OH⁻] from any one input. It also calculates pH for strong acids/bases at any concentration, weak acids/bases from Ka or Kb, and shows how dilution affects pH. A visual pH scale with everyday substance comparisons makes the results intuitive.
Instantly convert between pH, pOH, [H⁺], and [OH⁻]. Calculate pH for strong and weak acids/bases. Visualize where solutions fall on the pH scale. This ph calculator helps you compare outcomes quickly and reduce avoidable mistakes when making day-to-day care decisions. Use the estimate as a planning baseline and confirm final decisions with a qualified professional when risk is high.
pH = -log₁₀[H⁺] pOH = -log₁₀[OH⁻] pH + pOH = 14 (at 25°C) Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ Strong acid: pH = -log₁₀(C_acid × n) Weak acid: [H⁺] = √(Ka × C) when Ka << C or quadratic: [H⁺]² + Ka[H⁺] − Ka·C = 0
Result: pOH = 10.5, [H⁺] = 3.16 × 10⁻⁴ M, [OH⁻] = 3.16 × 10⁻¹¹ M
From pH = 3.5: [H⁺] = 10⁻³·⁵ = 3.16 × 10⁻⁴ M. pOH = 14 - 3.5 = 10.5. [OH⁻] = 10⁻¹⁰·⁵ = 3.16 × 10⁻¹¹ M. The solution is acidic (pH < 7).
The pH of common substances spans the entire scale: battery acid (pH ~0), stomach acid (1-2), lemon juice (~2), vinegar (~2.4), coffee (~5), milk (~6.5), pure water (7), blood (7.4), seawater (8.1), baking soda solution (~8.3), ammonia cleaner (~11.5), bleach (~12.5), and drain cleaner (~14).
For a weak acid HA with dissociation constant Ka and initial concentration C, the equilibrium concentration of H⁺ is found by solving Ka = x²/(C-x). When Ka << C, this simplifies to x ≈ √(Ka×C), giving pH = ½(pKa - logC). For weak bases, replace Ka with Kb and calculate pOH first, then pH = 14 - pOH.
Living systems use buffering, respiration, and kidney function to maintain pH. The bicarbonate buffer system (CO₂/HCO₃⁻) is the primary blood buffer. Respiratory compensation adjusts CO₂ levels within minutes, while renal compensation takes hours to days. Acidosis (pH < 7.35) and alkalosis (pH > 7.45) have distinct clinical presentations and treatments.
Yes. Concentrated strong acids can have pH < 0 (e.g., 10 M HCl has pH ≈ -1). Concentrated strong bases can exceed pH 14. The 0-14 range applies to dilute aqueous solutions at 25°C.
At 25°C, water autoionizes: Kw = [H⁺][OH⁻] = 10⁻¹⁴. When [H⁺] = [OH⁻], both equal 10⁻⁷, so pH = pOH = 7. At different temperatures, the neutral pH changes (e.g., ~6.14 at 100°C).
pH measures acidity (H⁺ concentration), pOH measures basicity (OH⁻ concentration). They always sum to pKw, which is 14.00 at 25°C.
Use the ICE table or the approximation pH = ½(pKa - log C). For the exact solution, solve the quadratic: x² + Ka·x - Ka·C = 0, where x = [H⁺].
Yes. Kw increases with temperature (2.4 × 10⁻¹⁴ at 37°C). Pure water at body temperature has pH = 6.81, not 7.00. The pH of buffers also shifts with temperature.
Digital pH meters use a glass electrode with a reference electrode. pH paper and indicators give approximate values based on color changes.