pH Calculator

pH Chemistry Guide

What's the key thing to understand about The pH Scale?

pH = −log₁₀[H⁺], where [H⁺] is the hydrogen ion concentration in mol/L. The scale runs from 0 (most acidic) to 14 (most basic), with 7 being neutral at 25°C. Each unit change represents a 10-fold difference in H⁺ concentration: pH 4 is 10× more acidic than pH 5, and 100× more acidic than pH 6. Common pH values: battery acid 0, stomach acid 1.5–3.5, lemon juice 2, coffee 5, rain water 5.6, pure water 7, blood 7.4, baking soda 8.3, bleach 12.5, drain cleaner 14.

What should I know about Calculating pH?

From [H⁺]: pH = −log([H⁺]). If [H⁺] = 0.0001 mol/L = 10⁻⁴, pH = 4. From pH to [H⁺]: [H⁺] = 10⁻ᵖᴴ. The pOH relationship: pH + pOH = 14 (at 25°C). For a weak acid: using the simplified formula when Ka << acid concentration, [H⁺] ≈ √(Ka × C). For example, acetic acid (Ka = 1.8×10⁻⁵) at 0.1 mol/L: [H⁺] = √(1.8×10⁻⁵ × 0.1) = 1.34×10⁻³, pH = 2.87.

What should I know about Buffers and pH Stability?

A buffer solution resists pH change when acid or base is added. Buffers are mixtures of weak acids and their conjugate bases (or weak bases and their conjugate acids). Blood (pH 7.35–7.45) is maintained by the bicarbonate buffer system. Even small deviations in blood pH (±0.2–0.3) cause serious physiological problems — acidosis (pH under 7.35) or alkalosis (pH over 7.45) can be life-threatening. Buffers are used extensively in biology labs, medicine (IV solutions), food (preservation), and photography (developing chemistry) wherever a stable pH is essential to the process.

What should I know about pH in Everyday Life?

Aquarium keeping: most freshwater fish need pH 6.5–7.5; some cichlids need 8.0+. Soil pH: most plants prefer 6.0–7.0; acidic plants (blueberries, rhododendrons) prefer 4.5–5.5; lawns prefer 6.0–7.0. Swimming pools: target pH 7.2–7.6 for sanitiser effectiveness and bather comfort. Skin: normal skin pH is 4.5–5.5 — alkaline soaps disrupt this barrier. Wine: typically pH 3–4. Bread dough: typically pH 5.5–6.0. Sourdough: pH 3.5–4.5.

What's the difference between Strong and Weak Acids and Bases?

A strong acid or base dissociates completely in water, while a weak one only partially dissociates, existing mostly as an equilibrium between the undissociated and dissociated forms. Strong acids (HCl, HNO3, H2SO4) and strong bases (NaOH, KOH) can be treated as fully ionised, so [H+] or [OH-] equals the stated concentration directly (adjusted for extra H+ or OH- per formula unit). Weak acids (ethanoic acid, carbonic acid) and weak bases (ammonia) require the dissociation constant (Ka or Kb) and an equilibrium calculation, because only a small fraction of molecules actually ionise. This is why 0.1 mol/L HCl has pH 1 (fully ionised) while 0.1 mol/L ethanoic acid (Ka = 1.8x10^-5) has pH around 2.87, nowhere near as acidic despite the same concentration. Strength (how much a substance ionises) is an entirely different property from concentration (how much substance is dissolved) — a concentrated weak acid can still have a higher pH than a dilute strong acid.

What should I know about pOH and the pH of Bases?

For bases, it is often easier to work via pOH and then convert. pOH = -log10[OH-], exactly analogous to pH. At 25°C, pH + pOH = 14 always holds, because the ionic product of water Kw = [H+][OH-] = 1x10^-14 is fixed at that temperature. For a strong base like NaOH, [OH-] equals the stated concentration directly. Example: 0.01 mol/L NaOH: [OH-] = 0.01, pOH = 2, so pH = 14 - 2 = 12. For a weak base like ammonia (Kb = 1.8x10^-5), [OH-] is approx sqrt(Kb x C). At 0.1 mol/L: [OH-] = sqrt(1.8x10^-5 x 0.1) = 1.34x10^-3, pOH = 2.87, pH = 11.13. Note that Kw itself is temperature-dependent — it increases as temperature rises, so pure water's pH is actually slightly below 7 above 25°C and above 7 below it, even though it remains perfectly neutral (pH = pOH) at every temperature.

What should I know about pH of Salts and Hydrolysis?

Not every salt dissolved in water gives a neutral pH of 7. A salt from a strong acid and strong base (like NaCl) is neutral, because neither ion reacts with water. A salt from a weak acid and strong base (like sodium acetate, CH₃COONa) gives a basic solution, because the conjugate base (acetate) hydrolyses, pulling H⁺ from water and leaving excess OH⁻. A salt from a strong acid and weak base (like ammonium chloride, NH₄Cl) gives an acidic solution, because the conjugate acid (ammonium) donates H⁺ to water instead. Predicting a salt's effect on pH is really just identifying which ion, if any, is a significant conjugate acid or base.

What should I know about Temperature's Effect on pH and Kw?

The ion product of water, Kw, is not a true constant — it increases with temperature, rising from about 1.0 × 10⁻¹⁴ at 25°C to roughly 2.9 × 10⁻¹⁴ at 40°C, because the dissociation of water into H⁺ and OH⁻ is itself an endothermic process that is pushed further forward as temperature rises. Since neutral water always has equal concentrations of H⁺ and OH⁻ by definition, this means the pH of neutral water is only exactly 7.00 at 25°C — at higher temperatures, neutral water has a pH slightly below 7, around 6.14 at 40°C, even though it is still perfectly neutral, because both ion concentrations have risen equally. This is a common source of confusion: a measured pH below 7 does not automatically mean a solution is acidic unless the measurement temperature is known, and it's why precise pH work always records temperature alongside the reading, and why pH meters include temperature compensation. The practical acid/base character of a solution is best judged by comparing its pH to the neutral pH at that specific temperature, not against a fixed value of 7.

What do I need to know about Calculating the pH of a Weak Acid?

Weak acids don't fully ionise, so you can't find pH directly from concentration the way you can for strong acids. Instead, use the acid dissociation constant Ka together with the approximation [H⁺] ≈ √(Ka × C), valid when the acid is weak enough that less than about 5% of it ionises. Worked example: find the pH of 0.10 M acetic acid (Ka = 1.8 × 10⁻⁵). [H⁺] = √(Ka × C) = √(1.8×10⁻⁵ × 0.10) = √(1.8×10⁻⁶) = 1.342×10⁻³ M pH = −log₁₀(1.342×10⁻³) = 2.87 Compare that to 0.10 M hydrochloric acid, essentially 100% ionised, giving [H⁺] = 0.10 M and pH = 1.00. The weak acid's pH is nearly two units higher despite the identical starting concentration.

The smaller the Ka (and so the larger the pKa), the weaker the acid. The same square-root approximation works for weak bases using Kb and [OH⁻] ≈ √(Kb × C), followed by pOH = −log₁₀[OH⁻] and pH = 14 − pOH at 25°C. For 0.10 M ammonia (Kb = 1.8×10⁻⁵, numerically identical to acetic acid's Ka): [OH⁻] = √(1.8×10⁻⁵×0.10) = 1.342×10⁻³ M, pOH = 2.87, so pH = 14 − 2.87 = 11.13 — exactly as far above neutral (pH 7) as acetic acid sits below it.

Worked Example: pH of a Weak Acid From Its Ka

Find the pH of a 0.10 M solution of acetic acid (Ka = 1.8 × 10^-5). Setting up the equilibrium HA ⇌ H+ + A-, with x as the concentration of H+ formed: Ka = x^2 / (0.10 - x) ≈ x^2 / 0.10 (since x is expected to be small relative to 0.10). Solving, x^2 = 1.8×10^-5 × 0.10 = 1.8×10^-6, so x = sqrt(1.8×10^-6) ≈ 1.34×10^-3 M. Checking the small-x assumption: 1.34×10^-3 is about 1.3% of 0.10, well under the usual 5% threshold for the approximation to be considered valid. So [H+] ≈ 1.34×10^-3 M, and pH = -log(1.34×10^-3) ≈ 2.87. Compare this to a strong acid at the same 0.10 M concentration, which would fully dissociate and give pH = -log(0.10) = 1.00 exactly -- the weak acid's pH is almost two full units higher, demonstrating how incomplete dissociation makes a weak acid solution far less acidic than a strong acid of the same nominal concentration.

pH of Common Household and Biological Substances

Approximate pH values of familiar substances help build intuition for where a calculated result should fall.

Human blood's pH is tightly regulated between about 7.35 and 7.45 by the bicarbonate buffer system (closely related to the buffer chemistry covered in the dedicated buffer calculator); even a shift of half a pH unit outside this narrow range is medically serious, which is why blood pH is one of the most tightly homeostatically controlled values in the entire human body, far tighter than the swings tolerated by most laboratory or industrial pH-controlled processes.