About pH Calculator
Our pH Calculator helps you calculate pH, pOH, hydrogen ion concentration [H⁺], and hydroxide ion concentration [OH⁻] with instant results. Perfect for chemistry students, laboratory technicians, water quality specialists, and anyone working with acid-base chemistry. Get accurate calculations with detailed explanations and real-world examples.
How to Use the pH Calculator
- Select your calculation type (pH → Concentrations or Concentration → pH)
- For pH to concentrations: Enter the pH value (0-14)
- For concentration to pH: Select ion type (H⁺ or OH⁻) and enter concentration in molarity (M)
- Click "Calculate" to get instant results
- View pH, pOH, and both ion concentrations
- See acid-base classification and common examples
What is pH?
pH is a measure of the acidity or alkalinity of a solution. It is defined as the negative logarithm (base 10) of the hydrogen ion concentration. The pH scale ranges from 0 to 14, with 7 being neutral. Values below 7 are acidic, and values above 7 are alkaline (basic).
The term "pH" stands for "potential of hydrogen" or "power of hydrogen." It was introduced by Danish chemist Søren Peder Lauritz Sørensen in 1909 and has become one of the most important measurements in chemistry, biology, and environmental science.
Key Formulas
- pH = -log[H⁺] (pH from hydrogen ion concentration)
- pOH = -log[OH⁻] (pOH from hydroxide ion concentration)
- pH + pOH = 14 (at 25°C)
- [H⁺] × [OH⁻] = 10⁻¹⁴ (ion product of water at 25°C)
- [H⁺] = 10⁻ᵖᴴ (hydrogen ion concentration from pH)
- [OH⁻] = 10⁻ᵖᴼᴴ (hydroxide ion concentration from pOH)
pH Scale Classification
- 0-3: Strongly Acidic - Battery acid, stomach acid, hydrochloric acid
- 3-7: Weakly Acidic - Lemon juice, vinegar, coffee, rain water
- 7: Neutral - Pure water, blood (7.4)
- 7-11: Weakly Alkaline - Baking soda, seawater, hand soap
- 11-14: Strongly Alkaline - Bleach, drain cleaner, lye
Common pH Values
Strongly Acidic (pH 0-3)
- Battery Acid: pH 0-1
- Stomach Acid: pH 1.5-2.0
- Lemon Juice: pH 2.0-2.6
- Vinegar: pH 2.4-3.4
Weakly Acidic (pH 3-7)
- Orange Juice: pH 3.3-4.2
- Tomato Juice: pH 4.0-4.6
- Coffee: pH 4.85-5.10
- Milk: pH 6.5-6.7
- Rain Water: pH 5.6 (slightly acidic due to CO₂)
Neutral (pH 7)
- Pure Water: pH 7.0 (at 25°C)
- Human Blood: pH 7.35-7.45 (slightly alkaline)
Weakly Alkaline (pH 7-11)
- Seawater: pH 7.5-8.4
- Baking Soda: pH 8.3
- Hand Soap: pH 9.0-10.0
- Milk of Magnesia: pH 10.5
Strongly Alkaline (pH 11-14)
- Ammonia Solution: pH 11.0-11.5
- Bleach: pH 12.5
- Drain Cleaner: pH 14
- Sodium Hydroxide (Lye): pH 13-14
Practical Applications
- Water Quality Testing: Municipal water treatment, drinking water safety, wastewater management
- Agriculture & Soil Management: Soil pH affects nutrient availability and plant growth
- Food & Beverage Production: pH control in fermentation, preservation, and quality
- Swimming Pool Maintenance: Optimal pH 7.2-7.8 for comfort and sanitation
- Medical & Biological Research: Blood pH, cell culture, enzyme activity
- Chemical Manufacturing: Process control, product quality, safety
- Aquarium & Fishkeeping: Freshwater (6.5-7.5), Saltwater (8.0-8.4)
- Environmental Monitoring: Acid rain, ocean acidification, ecosystem health
- Cosmetics & Personal Care: Skin pH (~5.5), hair care products
Understanding pOH
pOH is the negative logarithm of the hydroxide ion concentration, similar to how pH relates to hydrogen ions. The relationship pH + pOH = 14 (at 25°C) allows you to calculate one from the other. pOH is particularly useful when working with basic solutions where hydroxide concentration is more relevant.
Temperature Effects on pH
The pH scale is temperature-dependent because the ion product of water (Kw = [H⁺][OH⁻]) changes with temperature:
- At 0°C: Kw = 1.14 × 10⁻¹⁵, neutral pH = 7.47
- At 25°C: Kw = 1.00 × 10⁻¹⁴, neutral pH = 7.00
- At 50°C: Kw = 5.47 × 10⁻¹⁴, neutral pH = 6.63
- At 100°C: Kw = 5.13 × 10⁻¹³, neutral pH = 6.14
This calculator assumes 25°C (room temperature). For precise measurements at other temperatures, temperature compensation is required.
Buffer Solutions
Buffer solutions resist pH changes when small amounts of acid or base are added. They consist of a weak acid and its conjugate base (or weak base and conjugate acid). Common buffers include:
- Acetate Buffer: pH 3.6-5.6 (acetic acid/sodium acetate)
- Phosphate Buffer: pH 5.8-8.0 (dihydrogen phosphate/hydrogen phosphate)
- Carbonate Buffer: pH 9.2-10.8 (bicarbonate/carbonate)
- Tris Buffer: pH 7.0-9.0 (common in biological research)
pH Measurement Methods
- pH Meters: Electronic devices with glass electrodes, accurate to ±0.01 pH units
- pH Paper/Strips: Color-changing indicators, accurate to ±0.5-1.0 pH units
- pH Indicators: Chemical dyes that change color at specific pH ranges
- Universal Indicator: Mixture showing full pH range with different colors
Acids and Bases
Strong Acids (Completely Ionize)
- Hydrochloric acid (HCl)
- Sulfuric acid (H₂SO₄)
- Nitric acid (HNO₃)
- Perchloric acid (HClO₄)
Weak Acids (Partially Ionize)
- Acetic acid (CH₃COOH)
- Citric acid (C₆H₈O₇)
- Carbonic acid (H₂CO₃)
- Phosphoric acid (H₃PO₄)
Strong Bases (Completely Ionize)
- Sodium hydroxide (NaOH)
- Potassium hydroxide (KOH)
- Calcium hydroxide (Ca(OH)₂)
- Barium hydroxide (Ba(OH)₂)
Weak Bases (Partially Ionize)
- Ammonia (NH₃)
- Sodium bicarbonate (NaHCO₃)
- Aluminum hydroxide (Al(OH)₃)
pH in Biological Systems
- Human Blood: pH 7.35-7.45 (tightly regulated)
- Stomach: pH 1.5-3.5 (aids digestion)
- Small Intestine: pH 7.5-8.0 (neutralizes stomach acid)
- Saliva: pH 6.5-7.5
- Skin: pH 4.5-6.0 (acid mantle protects against bacteria)
- Urine: pH 4.5-8.0 (varies with diet)
Safety Considerations
- Extreme pH values (below 2 or above 12) are corrosive and dangerous
- Always wear appropriate PPE when handling strong acids or bases
- Never add water to concentrated acid; always add acid to water
- Calibrate pH meters regularly for accurate measurements
- Store pH electrodes properly in storage solution
- Dispose of chemical waste according to regulations
Frequently Asked Questions
Why is pH measured on a logarithmic scale?
The logarithmic scale compresses the wide range of hydrogen ion concentrations (from 1 M to 10⁻¹⁴ M) into a manageable scale (0-14). Each pH unit represents a 10-fold difference in hydrogen ion concentration. For example, pH 4 is 10 times more acidic than pH 5, and 100 times more acidic than pH 6.
Can pH be negative or greater than 14?
Yes, in very concentrated solutions. For example, 10 M HCl has pH = -1, and 10 M NaOH has pH = 15. However, the pH scale typically refers to dilute aqueous solutions where pH ranges from 0 to 14.
Why is pure water pH 7?
Pure water undergoes self-ionization: H₂O ⇌ H⁺ + OH⁻. At 25°C, [H⁺] = [OH⁻] = 10⁻⁷ M, giving pH = -log(10⁻⁷) = 7. This is neutral because hydrogen and hydroxide ions are present in equal concentrations.
How does pH affect enzyme activity?
Enzymes have optimal pH ranges where they function best. Changes in pH can alter the enzyme's shape (denaturation) and the ionization state of amino acids in the active site, affecting substrate binding and catalytic activity. For example, pepsin works best at pH 2 (stomach), while trypsin prefers pH 8 (small intestine).
What causes ocean acidification?
Increased atmospheric CO₂ dissolves in seawater, forming carbonic acid (H₂CO₃), which lowers pH. Ocean pH has decreased from 8.2 to 8.1 since pre-industrial times (a 30% increase in acidity). This affects marine life, especially organisms with calcium carbonate shells or skeletons.
Calculation Examples
Example 1: pH to Concentrations
Given: pH = 3.0
[H⁺] = 10⁻³ = 0.001 M = 1.0 × 10⁻³ M
pOH = 14 - 3 = 11.0
[OH⁻] = 10⁻¹¹ = 1.0 × 10⁻¹¹ M
Classification: Weakly Acidic
Example 2: H⁺ Concentration to pH
Given: [H⁺] = 1.0 × 10⁻⁵ M
pH = -log(1.0 × 10⁻⁵) = 5.0
pOH = 14 - 5 = 9.0
[OH⁻] = 10⁻⁹ = 1.0 × 10⁻⁹ M
Classification: Weakly Acidic
Example 3: OH⁻ Concentration to pH
Given: [OH⁻] = 1.0 × 10⁻⁴ M
pOH = -log(1.0 × 10⁻⁴) = 4.0
pH = 14 - 4 = 10.0
[H⁺] = 10⁻¹⁰ = 1.0 × 10⁻¹⁰ M
Classification: Weakly Alkaline
Example 4: Neutral Solution
Given: pH = 7.0 (pure water at 25°C)
[H⁺] = 10⁻⁷ = 1.0 × 10⁻⁷ M
pOH = 14 - 7 = 7.0
[OH⁻] = 10⁻⁷ = 1.0 × 10⁻⁷ M
Classification: Neutral
Advanced Topics
- Henderson-Hasselbalch Equation: pH = pKa + log([A⁻]/[HA]) for buffer calculations
- Activity vs. Concentration: pH technically measures H⁺ activity, not concentration
- Junction Potential: Voltage at the liquid junction in pH electrodes
- Nernst Equation: Relates electrode potential to ion concentration
- Polyprotic Acids: Acids with multiple ionizable protons (H₂SO₄, H₃PO₄)
- Amphoteric Substances: Can act as both acid and base (water, amino acids)