pH & Buffer Solution Calculator

pH & Buffer Guide

What's the key thing to understand about pH Basics?

pH = −log₁₀[H⁺]. Where [H⁺] is the hydrogen ion concentration in mol/L. Scale: 0-14 in water. Below 7: acidic. Equal 7: neutral (pure water at 25°C). Above 7: alkaline/basic. Each pH unit = 10× change in [H⁺]. pH 4 has 10× more H⁺ than pH 5, 100× more than pH 6. Strong acids (HCl, HNO₃, H₂SO₄) fully dissociate: pH = −log(concentration). 0.1 M HCl: pH = 1. 0.01 M HCl: pH = 2. Strong bases (NaOH, KOH) fully dissociate: pOH = −log(concentration). pH = 14 − pOH. 0.1 M NaOH: pH = 13. Weak acids/bases only partially dissociate, so their pH cannot be calculated directly from concentration alone — it depends on the acid dissociation constant Ka (or base constant Kb), which describes how far the equilibrium lies toward the dissociated ions.

What's the key thing to understand about Weak Acids and Ka?

Weak acids only partially ionise: HA ⇌ H⁺ + A⁻. Acid dissociation constant: Ka = [H⁺][A⁻] / [HA]. pKa = −log(Ka). Common pKa values: acetic acid (vinegar): 4.76. Carbonic acid: 6.35 (first dissociation). Phosphoric acid: 2.15, 7.20, 12.35 (three protons). Citric acid: 3.13, 4.76, 6.40. For weak acid pH: pH ≈ ½(pKa − log[HA]) — approximation works when [HA] >> Ka. Example: 0.1 M acetic acid: pH ≈ ½(4.76 − log 0.1) = ½(4.76 + 1) = 2.88. Verify: 0.1 M acetic acid pH measured ≈ 2.87 ✓ confirming the approximation is valid here since [HA] (0.1 M) is far larger than Ka (1.8×10⁻⁵ for acetic acid), so the amount dissociated is small enough not to meaningfully change the equilibrium concentration of the undissociated acid used in the calculation.

What's the key thing to understand about Buffers and Henderson-Hasselbalch?

Buffer: solution that resists pH change when small amounts of acid or base are added. Consists of weak acid + its conjugate base (or weak base + its conjugate acid). Henderson-Hasselbalch equation: pH = pKa + log([A⁻]/[HA]). Where [A⁻] = conjugate base concentration, [HA] = weak acid concentration. Buffer pH = pKa when [A⁻] = [HA]. Most effective buffering range: pKa ± 1. Examples: acetate buffer (acetic acid + sodium acetate): effective pH 3.76-5.76. Phosphate buffer (H₂PO₄⁻ + HPO₄²⁻): effective pH 6.1-8.1 — chosen in biology and medicine because it closely brackets physiological pH (blood pH is tightly held near 7.4). Choosing a buffer whose pKa sits close to the desired working pH gives the widest, most stable buffering capacity against small additions of acid or base.

What do I need to know about Biological pH Regulation?

Blood pH: maintained at 7.35-7.45. Tiny range — outside it causes serious illness or death. Below 7.35: acidosis. Above 7.45: alkalosis. Buffer systems in blood: carbonic acid/bicarbonate: HCO₃⁻ + H⁺ ⇌ H₂CO₃ ⇌ CO₂ + H₂O. Lungs control CO₂ removal. Kidneys regulate bicarbonate. Phosphate buffer: minor blood role, important in intracellular fluids and urine. Protein buffer: amino acid side chains (especially histidine) buffer within proteins and haemoglobin. Cellular pH: tightly regulated. Lysosomes maintain an acidic internal pH (around 4.5-5) distinct from the near-neutral cytoplasm, since the digestive enzymes they contain (hydrolases) only function effectively in that acidic environment — a leaked lysosome would have its enzymes largely inactive at cytoplasmic pH, which is itself a protective mechanism against accidental self-digestion of the cell.