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Calculate pH from [H⁺], pOH, or Ka. Convert between pH, pOH, [H⁺], and [OH⁻]. Essential for chemistry, biology, medicine, and environmental science acid-base calculations.
pH measures hydrogen ion concentration on a log scale from 0-14. pH 7 = neutral, below 7 = acidic, above 7 = basic. Each unit = 10-fold change in [H⁺]. Blood must stay pH 7.35-7.45 for survival.
| pH | Category | Examples |
|---|---|---|
| 0-2 | Strong acids | Battery acid, HCl |
| 2-4 | Weak acids | Lemon juice (2.2), vinegar (2.9) |
| 4-6 | Mild acids | Coffee (5.0), tomatoes (4.5) |
| 7 | Neutral | Pure water at 25°C |
| 7-8 | Slightly basic | Blood (7.4), seawater (8.1) |
| 10-12 | Bases | Ammonia (11) |
| 12-14 | Strong bases | Bleach (12.5), NaOH |
Blood pH 7.35-7.45 is critical — pH 7.0 causes fatal arrhythmias. Pool water needs pH 7.2-7.8 for effective chlorination. Most enzymes function at pH 7.4. Soil pH 6.0-7.0 ideal for plants.
Each pH unit = 10× change in [H⁺]. pH 3 is 100× more acidic than pH 5. At pH 7: [H⁺] = 10⁻⁷ mol/L.
Ka = acid dissociation constant. pKa = -log(Ka). For weak acid concentration C: pH ≈ ½(pKa - log C). Acetic acid: Ka = 1.8×10⁻⁵, pKa = 4.75.
pH is one of the most important parameters in chemistry and biology. It affects reaction rates, protein structure, enzyme activity, drug ionization, and almost every biological process. Understanding pH calculation and its implications is fundamental to chemistry education and laboratory practice.
Strong acids (HCl, H₂SO₄, HNO₃) dissociate completely: [H⁺] = initial concentration. 0.01 M HCl: [H⁺] = 0.01, pH = 2.0. Weak acids only partially dissociate — use Ka and the Henderson-Hasselbalch or quadratic approach. Acetic acid (Ka = 1.8×10⁻⁵): 0.1 M acetic acid has pH = ½(pKa - log 0.1) = ½(4.74 + 1) = 2.87, not 1.0 as for a strong acid.
pH paper: quick, semi-quantitative (±0.5 pH units). Universal indicator: visual, colorimetric. pH electrode: most accurate (±0.01 pH units). Always calibrate pH electrodes with two standard buffers (pH 4.00 and pH 7.00, or pH 7.00 and pH 10.00 depending on range) before use. Rinse electrode with distilled water between measurements. Store electrode in storage solution, never in distilled water.
The body maintains blood pH 7.35-7.45 through three systems: chemical buffers (bicarbonate, hemoglobin, phosphate), respiratory compensation (CO₂ adjustment), and renal compensation (HCO₃⁻ excretion or retention). Metabolic acidosis: low pH, low HCO₃⁻ (causes: diabetic ketoacidosis, renal failure, lactic acidosis). Respiratory acidosis: low pH, high CO₂ (causes: hypoventilation, COPD). Understanding blood gas interpretation requires mastery of pH and acid-base chemistry.
pH is the negative base-10 logarithm of the hydrogen ion concentration: pH = -log₁₀[H⁺]. It measures acidity and basicity on a scale from 0 (most acidic) to 14 (most basic), with 7.0 representing neutral pure water at 25°C. The logarithmic scale means each unit represents a 10-fold difference in hydrogen ion concentration — pH 3 has 10 times more H⁺ than pH 4, and 100 times more than pH 5.
pH is one of the most important chemical parameters in science, medicine, and industry. It affects reaction rates, protein structure, enzyme activity, drug ionization, nutrient availability in soil, aquatic life, and virtually every biological process. Understanding pH calculation is fundamental to chemistry education and laboratory work.
Strong acids dissociate completely: [H⁺] = initial acid concentration. HCl (0.01 M): [H⁺] = 0.01 M, pH = -log(0.01) = 2.00. H₂SO₄ (0.005 M, 2 H⁺ per molecule): [H⁺] = 0.01 M, pH = 2.00. Strong base NaOH (0.001 M): [OH⁻] = 0.001 M, pOH = 3.00, pH = 14 - 3 = 11.00. Strong acids and bases are fully ionized at any concentration — use concentration directly to find [H⁺] or [OH⁻].
Weak acids partially dissociate according to Ka. For acetic acid (CH₃COOH, Ka = 1.8 × 10⁻⁵) at 0.1 M: [H⁺] = √(Ka × C) = √(1.8 × 10⁻⁵ × 0.1) = √(1.8 × 10⁻⁶) = 1.34 × 10⁻³ M, pH = 2.87. This approximation works when Ka << C (less than 5% dissociation). For more accurate results, solve the quadratic equation: Ka = x²/(C-x). Common weak acid pKa values: acetic acid 4.74, carbonic acid 6.35, phosphoric acid 7.20, ammonium 9.25.
Biological systems maintain precise pH for enzyme function and protein stability. Blood: pH 7.35-7.45 (critical range — death occurs outside pH 6.8-8.0). Stomach acid: pH 1.5-3.5 (HCl secretion for protein digestion and pathogen killing). Small intestine: pH 6-7.4 (pancreatic bicarbonate neutralizes stomach acid). Urine: pH 4.5-8.0 (varies with diet and kidney function). Cytoplasm: pH 7.0-7.4. Lysosomes (cellular waste processors): pH 4.5-5.0 (acidic to activate degradative enzymes).
pH indicator paper: quick, semi-quantitative (±0.5 pH units), inexpensive. Universal indicator solution: colorimetric, useful for teaching and rough estimates. Digital pH electrode: most accurate (±0.01 pH units), requires two-point calibration with pH 4.00 and pH 7.00 standard buffers, must be rinsed between samples, stored in 3M KCl or storage solution (never distilled water). Spectrophotometric pH indicators: precise and remotely readable, used in ocean chemistry monitoring (bromothymol blue, phenolphthalein).
Blood gas analysis interprets pH in context of CO₂ and HCO₃⁻. Normal: pH 7.40, pCO₂ 40 mmHg, HCO₃⁻ 24 mEq/L. Respiratory acidosis (high CO₂): pH <7.35, elevated pCO₂ — causes: hypoventilation, COPD, pneumonia. Metabolic acidosis (low HCO₃⁻): pH <7.35, low bicarbonate — causes: diabetic ketoacidosis, renal failure, lactic acidosis, diarrhea. Respiratory alkalosis (low CO₂): pH >7.45, low pCO₂ — cause: hyperventilation, anxiety, PE. Metabolic alkalosis (high HCO₃⁻): pH >7.45 — causes: vomiting (HCl loss), loop diuretics, excessive antacid use.