pH Calculator | pH, pOH & Ions
Calculate pH, pOH, [H⁺], and [OH⁻] concentrations for acids and bases.
Common solutions
What Is the pH Calculator | pH, pOH & Ions?
This pH calculator converts between all four acid-base measures, pH, pOH, [H⁺] concentration, and [OH⁻] concentration, in any direction. Enter any one value and get all others instantly. A colour-coded scale shows where the solution falls from strongly acidic (red) to strongly basic (violet), with 8 preset common solutions for quick reference.
- ›4 input modes: Enter pH, pOH, [H⁺] in mol/L, or [OH⁻] in mol/L, the other three are derived.
- ›pH scale bar: Visual gradient from acidic red to basic violet with a marker at the calculated pH.
- ›Classification: Strongly Acidic → Acidic → Weakly Acidic → Neutral → Weakly Basic → Basic → Strongly Basic.
- ›8 presets: Pure water, stomach acid, blood, coffee, lemon juice, ammonia, baking soda, bleach.
- ›6 decimal precision: pH, pOH, [H⁺], and [OH⁻] all shown to 4–6 significant figures.
Formula
| pH range | Classification | Examples |
|---|---|---|
| 0–2 | Strongly Acidic | Battery acid, gastric acid (pH ~1) |
| 2–5 | Acidic | Lemon juice (2.5), vinegar (3), coffee (5) |
| 5–7 | Weakly Acidic | Milk (6.5), rainwater (5.6), skin (5.5) |
| 7 | Neutral | Pure water at 25°C |
| 7–9 | Weakly Basic | Seawater (8), baking soda (9) |
| 9–12 | Basic | Ammonia (11), milk of magnesia (10) |
| 12–14 | Strongly Basic | Bleach (12.5), drain cleaner (14) |
How to Use
- 1Select input mode: pH, pOH, [H⁺], or [OH⁻].
- 2Click a preset (e.g. Blood, Coffee) to load a known pH, or type your value.
- 3Press Enter or click Calculate.
- 4Read the pH scale bar, classification, and all four derived values.
- 5Click Clear to reset.
Example Calculation
Finding all values from pH = 3
pH and temperature
Understanding pH | pH, pOH & Ions
The pH Scale and Logarithms
pH = −log₁₀[H⁺], introduced by Søren Sørensen in 1909. The logarithmic scale compresses the enormous range of hydrogen ion concentrations (10⁻¹⁵ to 10⁰ mol/L) into the convenient 0–14 range. Each pH unit represents a 10-fold change in [H⁺]: pH 3 is 10× more acidic than pH 4, and 100× more acidic than pH 5. This exponential relationship means small pH changes can represent large chemical differences.
pH in Biology and Medicine
Human blood is maintained between pH 7.35 and 7.45 by carbonate and phosphate buffer systems. Deviations outside this range cause acidosis (pH < 7.35) or alkalosis (pH > 7.45), both life-threatening. The seemingly tiny 0.1 unit range actually represents a 26% change in [H⁺], illustrating how biologically sensitive this parameter is. Enzyme activity, protein folding, oxygen transport by hemoglobin, and membrane potential all depend critically on pH.
pH in Industry and Environment
- ›Agriculture: most crops grow best at pH 6–7; acidic soils (pH < 5.5) need liming
- ›Water treatment: municipal water maintained near pH 7–7.5 to minimize pipe corrosion
- ›Swimming pools: pH 7.2–7.8 for comfort and disinfection efficacy
- ›Food preservation: low pH (vinegar pickling, fermentation) inhibits bacterial growth
- ›Electroplating: pH controls metal ion speciation and deposit quality
Frequently Asked Questions
What does pH mean and where does the term come from?
pH stands for "potential of hydrogen" (from the German "Potenz" meaning power/exponent). It measures the molar concentration of hydrogen ions (H⁺) in aqueous solution on a negative base-10 logarithmic scale: pH = −log₁₀[H⁺]. Danish biochemist Søren Sørensen introduced the notation in 1909 to replace unwieldy scientific notation, writing pH 7 rather than [H⁺] = 0.0000001 mol/L. The scale from 0 to 14 covers all common solutions at 25°C.
- ›pH < 7: acidic, more H⁺ ions than OH⁻ ions
- ›pH = 7: neutral, equal H⁺ and OH⁻ concentrations (pure water at 25°C)
- ›pH > 7: basic (alkaline), more OH⁻ ions than H⁺ ions
- ›Each pH unit = exactly 10× change in [H⁺]: pH 3 has 100× more H⁺ than pH 5
- ›The 0–14 range covers all biologically and environmentally relevant aqueous solutions
Why is the pH scale logarithmic?
Hydrogen ion concentrations in natural systems span about 15 orders of magnitude, from 1 mol/L (strongly acidic, pH 0) to 10⁻¹⁵ mol/L (strongly basic, pH 15). A linear scale would require writing numbers like 0.000000001 mol/L, making comparisons impractical. A logarithmic scale compresses this enormous range into a convenient 0–14 interval where each unit represents a consistent 10× change in actual concentration. The same approach is used in the Richter scale for earthquakes and decibels for sound.
- ›pH 1 (stomach acid) vs pH 7 (pure water): stomach acid has 10⁶ times more H⁺
- ›pH 3 (vinegar) vs pH 5 (coffee): vinegar has 100× the hydrogen ion concentration
- ›The "small" change from blood pH 7.40 to 7.35 = a 12% increase in [H⁺]
- ›Without logarithms, you would need to compare 0.0000004 vs 0.00000004, logarithms make this 6.4 vs 7.4
Why does pH + pOH = 14?
At 25°C, water partially self-ionises: H₂O ⇌ H⁺ + OH⁻. The equilibrium constant for this reaction is Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴. Taking −log₁₀ of both sides: −log[H⁺] + (−log[OH⁻]) = 14, which is exactly pH + pOH = 14. Increasing [H⁺] forces a proportional decrease in [OH⁻] to maintain the constant product Kw. This relationship only holds exactly at 25°C, at other temperatures, Kw changes and so does the sum.
- ›At 37°C (body temperature): Kw ≈ 2.4×10⁻¹⁴, so pH + pOH ≈ 13.62
- ›At 60°C: Kw ≈ 9.6×10⁻¹⁴, so pH + pOH ≈ 13.0
- ›At 100°C: Kw ≈ 10⁻¹², so pH + pOH = 12, neutral pH is about 6, not 7
- ›Only at 25°C does "neutral" exactly equal pH = pOH = 7
What is [H⁺] for a given pH and how are they related?
[H⁺] = 10⁻ᵖᴴ mol/L, this is the direct inverse of the pH definition. Going from pH to concentration involves raising 10 to the negative pH power; going the other direction, pH = −log₁₀[H⁺]. Because the relationship is exponential, a small change in pH corresponds to a large change in actual H⁺ concentration, a fact that is easy to underestimate when looking only at pH numbers.
What is pOH and when do I use it?
pOH = −log₁₀[OH⁻] is the hydroxide-ion analogue of pH. While pH is the more commonly reported measure, pOH is the natural description for strongly basic solutions, specifying pOH = 3 (meaning [OH⁻] = 0.001 mol/L, strongly basic) is more intuitive in some laboratory contexts than pH = 11. Since pH + pOH = 14 at 25°C, knowing one gives the other instantly. Chemists preparing alkaline reagents often find it easier to work in pOH directly.
- ›pOH 1: very strongly basic, [OH⁻] = 0.1 mol/L, pH = 13
- ›pOH 3: strongly basic, [OH⁻] = 0.001 mol/L, pH = 11 (ammonia region)
- ›pOH 7: neutral, same as pH 7 at 25°C (pure water)
- ›pOH 11: weakly acidic, [OH⁻] = 10⁻¹¹ mol/L, pH = 3
- ›pOH is most useful when titrating strong bases or computing [OH⁻] from Kw
What is the pH of blood and why must it stay constant?
Normal arterial blood pH is 7.35–7.45, slightly alkaline. This narrow range is maintained by three buffer systems working simultaneously: the carbonate/bicarbonate system (the primary rapid buffer), plasma proteins, and the phosphate buffer. The lungs provide minute-to-minute regulation by adjusting CO₂ excretion rate; the kidneys provide slower but more sustained correction over hours to days. The seemingly small 0.1-unit range actually represents a 26% change in [H⁺], illustrating the biological precision required.
- ›Acidosis (pH < 7.35): CO₂ retention (respiratory), lactic acid or ketoacid buildup (metabolic)
- ›Alkalosis (pH > 7.45): hyperventilation (respiratory) or excess HCO₃⁻ (metabolic)
- ›pH < 7.0 or > 7.7: typically life-threatening without rapid medical intervention
- ›Enzyme active sites are pH-sensitive, even a 0.1 unit shift alters enzyme kinetics significantly
- ›Hemoglobin releases O₂ more readily at lower pH (Bohr effect), critical for tissue oxygen delivery
Can pH be negative or greater than 14?
Yes, pH = −log₁₀[H⁺] has no mathematical restriction to the 0–14 range. The conventional 0–14 scale is simply the range relevant for dilute aqueous solutions at 25°C, covering almost all biological and environmental situations. At very high acid or base concentrations, the effective hydrogen ion activity (measured by the Hammett acidity function H₀) extends far beyond these bounds, and ion activity corrections become important.
- ›Concentrated HCl (12 mol/L): pH ≈ −1.1 (activity correction needed)
- ›Concentrated H₂SO₄ (18 mol/L): pH ≈ −1.4
- ›Concentrated NaOH (10 mol/L): pH ≈ 15
- ›Superacids such as fluorosulfuric acid (HSO₃F): Hammett H₀ values reaching −20 or lower
- ›Note: at high concentrations, activity ≠ concentration, the Debye-Hückel correction applies