pH Calculator
Calculate pH from hydrogen ion concentration, or find [H+] from pH. Includes pOH and acid/base determination.
Calculate pH
Calculate From:
= 1.00e-6 M
Common Substances:
pH Value
6.0000
Weakly Acidic
pH Scale:
Key Formulas:
pH = -log₁₀[H⁺]
pOH = -log₁₀[OH⁻]
pH + pOH = 14 (at 25°C)
[H⁺][OH⁻] = 10⁻¹⁴ (Kw)
pH of Common Substances
0
Battery Acid
1.5
Stomach Acid
2.4
Lemon Juice
2.9
Vinegar
3.5
Orange Juice
5
Coffee
6.5
Milk
7
Pure Water
7.4
Blood
8.3
Baking Soda
11
Ammonia
12.5
Bleach
14
Drain Cleaner
Understanding pH
The pH scale measures how acidic or basic a solution is. It ranges from 0 to 14, with 7 being neutral. Solutions with a pH less than 7 are acidic, while those greater than 7 are basic (alkaline). The scale is logarithmic, meaning each unit change represents a 10-fold change in hydrogen ion concentration. For example, pH 4 is 10 times more acidic than pH 5, and 100 times more acidic than pH 6.
What Is pH?
pH is a logarithmic scale measuring the acidity or basicity of an aqueous solution, ranging from 0 to 14. It quantifies the hydrogen ion (H⁺) concentration, with each unit representing a tenfold difference in acidity. pH is fundamental to chemistry, biology, medicine, and environmental science.
| pH Range | Classification | H⁺ Concentration | Examples |
|---|---|---|---|
| 0-1 | Strongly acidic | 1-0.1 M | Battery acid, stomach acid |
| 2-3 | Acidic | 10⁻²-10⁻³ M | Lemon juice, vinegar |
| 4-6 | Weakly acidic | 10⁻⁴-10⁻⁶ M | Coffee, rain, milk |
| 7 | Neutral | 10⁻⁷ M | Pure water at 25°C |
| 8-10 | Weakly basic | 10⁻⁸-10⁻¹⁰ M | Sea water, baking soda |
| 11-14 | Strongly basic | 10⁻¹¹-10⁻¹⁴ M | Ammonia, bleach, lye |
pH Definition
Where:
- pH= Measure of acidity (0-14 scale)
- [H⁺]= Hydrogen ion concentration (mol/L)
- log₁₀= Logarithm base 10
pH, pOH, and the Water Equilibrium
Water undergoes autoionization: H₂O ⇌ H⁺ + OH⁻. At 25°C, the ion product of water (Kw) is 10⁻¹⁴, establishing the relationship between pH and pOH.
| Relationship | Formula | At 25°C |
|---|---|---|
| Ion product of water | Kw = [H⁺][OH⁻] | Kw = 10⁻¹⁴ |
| pH + pOH | pH + pOH = pKw | pH + pOH = 14 |
| pOH definition | pOH = -log₁₀[OH⁻] | — |
| Neutral solution | [H⁺] = [OH⁻] | pH = pOH = 7 |
pH-pOH Relationships
Where:
- pOH= Measure of basicity
- [OH⁻]= Hydroxide ion concentration (mol/L)
- Kw= Ion product of water (10⁻¹⁴ at 25°C)
pH of Strong Acids and Bases
Strong acids and bases dissociate completely in water, making pH calculations straightforward—the H⁺ or OH⁻ concentration equals the acid or base concentration.
| Strong Acid | Formula | Strong Base | Formula |
|---|---|---|---|
| Hydrochloric acid | HCl → H⁺ + Cl⁻ | Sodium hydroxide | NaOH → Na⁺ + OH⁻ |
| Nitric acid | HNO₃ → H⁺ + NO₃⁻ | Potassium hydroxide | KOH → K⁺ + OH⁻ |
| Sulfuric acid* | H₂SO₄ → 2H⁺ + SO₄²⁻ | Calcium hydroxide | Ca(OH)₂ → Ca²⁺ + 2OH⁻ |
| Hydrobromic acid | HBr → H⁺ + Br⁻ | Barium hydroxide | Ba(OH)₂ → Ba²⁺ + 2OH⁻ |
Note: For strong acids, pH = -log[acid]. For strong bases, pOH = -log[base], then pH = 14 - pOH.
pH of Weak Acids and Bases
Weak acids and bases only partially dissociate. Their pH depends on both concentration and dissociation constant (Ka or Kb).
| Weak Acid | pKa | Weak Base | pKb |
|---|---|---|---|
| Acetic acid (CH₃COOH) | 4.76 | Ammonia (NH₃) | 4.75 |
| Carbonic acid (H₂CO₃) | 6.35 | Pyridine | 8.77 |
| Formic acid (HCOOH) | 3.75 | Methylamine | 3.36 |
| Hydrofluoric acid (HF) | 3.17 | Trimethylamine | 4.19 |
Weak Acid pH Formula
Where:
- Ka= Acid dissociation constant
- Kb= Base dissociation constant
- C= Initial concentration
- pKa= -log₁₀(Ka)
pH of Buffer Solutions
Buffers resist pH changes by containing both a weak acid and its conjugate base (or weak base and conjugate acid). The Henderson-Hasselbalch equation calculates buffer pH.
| Buffer System | pKa | Useful pH Range | Application |
|---|---|---|---|
| Acetate (CH₃COOH/CH₃COO⁻) | 4.76 | 3.8-5.8 | Food preservation |
| Phosphate (H₂PO₄⁻/HPO₄²⁻) | 7.20 | 6.2-8.2 | Biological systems |
| Tris (Tris base/Tris-HCl) | 8.06 | 7.0-9.0 | Biochemistry |
| Bicarbonate (H₂CO₃/HCO₃⁻) | 6.35 | 5.4-7.4 | Blood buffering |
Henderson-Hasselbalch Equation
Where:
- [A⁻]= Conjugate base concentration
- [HA]= Weak acid concentration
- pKa= Acid dissociation constant (-log Ka)
Methods of Measuring pH
Several techniques exist for measuring pH, each with different accuracy, range, and applications.
| Method | Accuracy | Range | Best For |
|---|---|---|---|
| pH meter (electrode) | ±0.01 pH | 0-14 | Precise lab measurements |
| pH paper (strips) | ±0.5 pH | 0-14 | Quick estimates, fieldwork |
| Universal indicator | ±1 pH | 0-14 | Visual demonstrations |
| Specific indicators | ±0.2 pH | Narrow range | Titration endpoints |
| Colorimetric kits | ±0.2 pH | Varies | Pool/aquarium testing |
Calibration: pH meters require calibration with buffer solutions (typically pH 4, 7, and 10) before use for accurate readings.
Importance of pH in Different Fields
pH affects nearly every chemical and biological process. Maintaining proper pH is critical across many fields.
| Field | Optimal pH | Why It Matters |
|---|---|---|
| Human blood | 7.35-7.45 | Enzyme function, oxygen transport |
| Stomach | 1.5-3.5 | Protein digestion, killing pathogens |
| Drinking water | 6.5-8.5 | Taste, pipe corrosion prevention |
| Swimming pools | 7.2-7.8 | Chlorine effectiveness, comfort |
| Soil for plants | 5.5-7.5 | Nutrient availability |
| Aquariums | 6.5-8.5 | Fish health, species-specific |
| Wine making | 3.0-4.0 | Flavor, microbial stability |
| Skincare products | 4.5-6.5 | Skin barrier protection |
Worked Examples
pH of a Strong Acid
Problem:
Calculate the pH of 0.01 M hydrochloric acid (HCl).
Solution Steps:
- 1HCl is a strong acid—it dissociates completely: HCl → H⁺ + Cl⁻
- 2Therefore, [H⁺] = [HCl] = 0.01 M = 10⁻² M
- 3Apply pH formula: pH = -log₁₀[H⁺] = -log₁₀(10⁻²)
- 4pH = -(-2) = 2
Result:
pH = 2. This is a moderately acidic solution. Note: For very dilute strong acids (< 10⁻⁶ M), you must also consider H⁺ from water autoionization.
pH of a Strong Base
Problem:
Calculate the pH of 0.001 M sodium hydroxide (NaOH).
Solution Steps:
- 1NaOH is a strong base—it dissociates completely: NaOH → Na⁺ + OH⁻
- 2[OH⁻] = [NaOH] = 0.001 M = 10⁻³ M
- 3Calculate pOH: pOH = -log₁₀(10⁻³) = 3
- 4Use pH + pOH = 14: pH = 14 - 3 = 11
Result:
pH = 11. This is a moderately basic solution. Alternatively, calculate [H⁺] = Kw/[OH⁻] = 10⁻¹⁴/10⁻³ = 10⁻¹¹, then pH = 11.
pH of a Weak Acid
Problem:
Calculate the pH of 0.1 M acetic acid (Ka = 1.8 × 10⁻⁵).
Solution Steps:
- 1Set up equilibrium: CH₃COOH ⇌ H⁺ + CH₃COO⁻
- 2Use approximation: [H⁺] = √(Ka × C) = √(1.8 × 10⁻⁵ × 0.1)
- 3[H⁺] = √(1.8 × 10⁻⁶) = 1.34 × 10⁻³ M
- 4pH = -log(1.34 × 10⁻³) = 2.87
Result:
pH ≈ 2.87. Acetic acid is a weak acid (only ~1.3% dissociated at this concentration), so the pH is higher than a strong acid at the same concentration would give (pH = 1).
Tips & Best Practices
- ✓Each pH unit represents a 10× change in H⁺ concentration—pH 3 is 100× more acidic than pH 5.
- ✓For strong acids: pH = -log[acid]. For strong bases: pOH = -log[base], then pH = 14 - pOH.
- ✓Weak acids require Ka for pH calculation: [H⁺] ≈ √(Ka × C) when C >> Ka.
- ✓Buffer pH is closest to pKa—choose an acid with pKa near your target pH.
- ✓Always calibrate pH meters with at least two buffer solutions before measuring.
- ✓Temperature affects pH readings—specify temperature for precise work (usually 25°C).
- ✓The Henderson-Hasselbalch equation only applies to buffers, not to pure acids or bases.
Frequently Asked Questions
Sources & References
Last updated: 2026-01-22