Henderson-Hasselbalch Calculator
Calculate buffer pH using the Henderson-Hasselbalch equation. Determine pH from pKa and concentration ratio, or solve for any variable. Includes buffer capacity and common buffer recipes.
pH Calculator
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.
Enter value 0-14
Dilute by 10× to see pH change
pH 3.50
0 (Acidic)7 (Neutral)14 (Basic)
pH⧉
3.5000
Acidic
pOH⧉
10.5000
pH + pOH = 14
[H⁺]⧉
3.162e-4 M
Hydrogen ion concentration
[OH⁻]⧉
3.162e-11 M
Hydroxide ion concentration
pH after 10× dilution⧉
4.500
Shift: 1.000 pH units
Common Substances pH
| Substance | pH | Nature | Scale |
|---|---|---|---|
| Battery acid | 0 | Acidic | |
| Stomach acid | 1.5 | Acidic | |
| Lemon juice | 2 | Acidic | |
| Vinegar | 2.4 | Acidic | |
| Orange juice | 3.5 | Acidic | |
| Coffee | 5 | Acidic | |
| Milk | 6.5 | Acidic | |
| Pure water | 7 | Neutral | |
| Blood | 7.4 | Basic | |
| Seawater | 8.1 | Basic | |
| Baking soda | 8.3 | Basic | |
| Ammonia | 11.5 | Basic | |
| Bleach | 12.5 | Basic | |
| Drain cleaner | 14 | Basic |
Planning notes, formulas, and examples
About the pH Calculator
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.
When This Page Helps
Quickly convert between pH, pOH, [H⁺], and [OH⁻]. Calculate pH for strong and weak acids/bases. Visualize where solutions fall on the pH scale.
How to Use the Inputs
- Enter any one of: pH, pOH, [H⁺], or [OH⁻], and all others are calculated.
- Switch mode to calculate pH of a strong acid, strong base, or weak acid/base.
- For weak acid/base mode, enter concentration and Ka/Kb.
- View the pH visual scale with common substance markers.
- Compare with everyday substances in the reference table.
- Use the dilution section to see how adding water affects pH.
Formula used
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 = 0Example Calculation
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).
Tips & Best Practices
- For strong monoprotic acids, pH = -log(concentration) exactly.
- For weak acids, the approximation pH ≈ ½(pKa - log C) works when C >> Ka.
- Diluting a strong acid by 10× raises pH by exactly 1 unit (until very dilute).
- pH paper has ±0.5 accuracy; a calibrated meter gives ±0.01.
- Blood pH outside 6.8-7.8 is generally fatal; normal is 7.35-7.45.
- Most enzymes have narrow optimal pH ranges (e.g., pepsin: pH 2, trypsin: pH 8).
The pH Scale in Everyday Life
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).
Weak Acid and Weak Base pH Calculations
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.
Biological pH Regulation
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.
Sources & Methodology
Last updated:
Frequently Asked Questions
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.
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