pH Calculator
How pH Works
pH is a logarithmic measure of the hydrogen ion concentration in an aqueous solution, expressing how acidic or basic that solution is on a scale that typically ranges from 0 to 14. The concept was introduced in 1909 by Danish chemist Soren Sorensen at the Carlsberg Laboratory to simplify the expression of hydrogen ion activity in solutions. According to the IUPAC Compendium of Chemical Terminology, pH is formally defined as the negative base-10 logarithm of the hydrogen ion activity in a solution.
Because pH is logarithmic, each whole-number change represents a tenfold change in hydrogen ion concentration. A solution with pH 3 contains 10 times more H+ ions than one at pH 4, and 100 times more than pH 5. Pure water at 25 degrees Celsius has a pH of exactly 7.0, which is considered neutral. Solutions below pH 7 are acidic, while those above pH 7 are basic (alkaline). The U.S. Environmental Protection Agency notes that most natural water bodies have a pH between 6.5 and 8.5, and deviations outside this range can harm aquatic ecosystems.
pH measurement is critical across many disciplines. In medicine, human blood pH must remain between 7.35 and 7.45 for normal physiological function. In agriculture, soil pH between 6.0 and 7.0 is optimal for most crops because it maximizes nutrient availability. Industrial applications include water treatment, food processing, pharmaceutical manufacturing, and chemical synthesis, where precise pH control can determine product quality and safety. This calculator converts between pH, pOH, hydrogen ion concentration [H+], and hydroxide ion concentration [OH-], giving you a complete picture of any aqueous solution's acid-base chemistry.
The pH Formula and How It Is Calculated
The fundamental pH equation is: pH = -log10[H+], where [H+] is the molar concentration of hydrogen ions (in mol/L). The complementary measure, pOH, is calculated as pOH = -log10[OH-]. At 25 degrees Celsius, these values are related by the water autoionization constant (Kw = 1.0 x 10^-14): pH + pOH = 14, and [H+] x [OH-] = 1.0 x 10^-14.
Worked example: A solution has a hydrogen ion concentration of 0.001 mol/L (1 x 10^-3 M). pH = -log10(1 x 10^-3) = 3.00. Since pH + pOH = 14, pOH = 14 - 3 = 11.00. The hydroxide concentration is [OH-] = 10^-11 = 1 x 10^-11 M. This solution is strongly acidic. Conversely, if you know pH = 8.5, then [H+] = 10^-8.5 = 3.16 x 10^-9 M, and pOH = 5.5.
Key Terms You Should Know
- Hydrogen Ion Concentration [H+] — The molar concentration of hydrogen ions (protons) in a solution, measured in moles per liter (mol/L). Higher [H+] means lower pH and greater acidity.
- Hydroxide Ion Concentration [OH-] — The molar concentration of hydroxide ions in a solution. Higher [OH-] means higher pH and greater basicity (alkalinity).
- pOH — The negative logarithm of hydroxide ion concentration. It is the complementary measure to pH: pH + pOH = 14 at standard temperature.
- Buffer Solution — A solution that resists changes in pH when small amounts of acid or base are added, typically made from a weak acid and its conjugate base (or vice versa). Blood uses a bicarbonate buffer system to maintain pH near 7.4.
- Water Dissociation Constant (Kw) — The equilibrium constant for water's autoionization: H2O ⇌ H+ + OH-. At 25 degrees C, Kw = 1.0 x 10^-14. This value increases with temperature, meaning pure water's pH drops below 7 at higher temperatures.
- Indicator — A substance that changes color at a specific pH range. Litmus paper turns red below pH 4.5 and blue above pH 8.3. Universal indicator gives a continuous color gradient across the full pH scale.
pH Values of Common Substances
The following reference table shows the approximate pH of everyday substances, from strongly acidic to strongly basic. These values are widely cited in chemistry education and verified against data from the CRC Handbook of Chemistry and Physics. Understanding where common substances fall on the pH scale helps contextualize calculated results.
| Substance | Approximate pH | Classification |
|---|---|---|
| Battery acid (H2SO4) | 0.5 - 1.0 | Strongly acidic |
| Gastric acid (stomach) | 1.5 - 2.0 | Strongly acidic |
| Lemon juice | 2.0 - 2.6 | Acidic |
| Vinegar | 2.4 - 3.4 | Acidic |
| Orange juice | 3.3 - 4.2 | Acidic |
| Coffee | 4.8 - 5.1 | Mildly acidic |
| Milk | 6.3 - 6.6 | Slightly acidic |
| Pure water (25 C) | 7.0 | Neutral |
| Human blood | 7.35 - 7.45 | Slightly basic |
| Seawater | 7.8 - 8.3 | Mildly basic |
| Baking soda solution | 8.3 - 8.5 | Basic |
| Household ammonia | 11.0 - 12.0 | Strongly basic |
| Household bleach (NaOCl) | 12.5 - 13.5 | Strongly basic |
| Drain cleaner (NaOH) | 13.0 - 14.0 | Strongly basic |
Practical Examples
Example 1 — Pool Water Testing: You measure [H+] = 6.31 x 10^-8 M in your swimming pool. pH = -log10(6.31 x 10^-8) = 7.2. The CDC recommends pool water pH between 7.2 and 7.8, so this reading is at the lower acceptable limit. You may want to add a small amount of sodium carbonate (soda ash) to raise the pH slightly.
Example 2 — Soil Acidity for Blueberries: A soil test shows pH 5.5. Blueberries thrive at pH 4.5-5.5, so this soil is at the upper edge of the ideal range. The [H+] = 10^-5.5 = 3.16 x 10^-6 M. To lower pH further, you could add elemental sulfur, which soil bacteria convert to sulfuric acid over several weeks. Use our Molarity Calculator to determine the sulfur concentration needed.
Example 3 — Laboratory Buffer Preparation: You need a phosphate buffer at pH 7.4 for a biological experiment. The pKa of the H2PO4-/HPO4^2- pair is 7.2. Using the Henderson-Hasselbalch equation (pH = pKa + log([A-]/[HA])), you calculate the ratio: 7.4 = 7.2 + log(ratio), so ratio = 10^0.2 = 1.58. You need 1.58 parts HPO4^2- for every 1 part H2PO4- to achieve pH 7.4.
Tips and Strategies for pH Measurement
- Calibrate pH meters regularly using at least two standard buffer solutions (pH 4.0 and 7.0, or 7.0 and 10.0) before each measurement session. Uncalibrated meters can drift by 0.5 pH units or more.
- Account for temperature when interpreting pH. The Kw of water changes with temperature: at 37 degrees C (body temperature), pure water has pH 6.81, not 7.0. Most modern pH meters include automatic temperature compensation.
- Use the right indicator for the range. Phenolphthalein works well for pH 8-10 (colorless to pink) but is useless for acidic solutions. Methyl orange covers pH 3.1-4.4. For broad-range testing, universal indicator paper covers the full 0-14 scale.
- Understand buffer capacity. A buffer works best within one pH unit of its pKa. Outside that range, the buffer is overwhelmed and pH changes rapidly. Choose a buffer system whose pKa is close to your target pH.
- Consider dilution effects carefully. Diluting a strong acid by a factor of 10 raises its pH by exactly 1 unit (e.g., pH 2 becomes pH 3). But diluting a weak acid or buffer changes pH less predictably — use the Henderson-Hasselbalch equation for accuracy. Our Dilution Calculator can help with concentration changes.
- Store pH electrodes properly. Keep glass pH electrodes in storage solution (3M KCl), never in distilled water or dry. Dry storage permanently damages the glass membrane and leads to slow, inaccurate readings.
This calculator is for informational purposes only and does not constitute financial, tax, or legal advice. Always consult a qualified professional for decisions specific to your situation.
Frequently Asked Questions
What is the pH of common household substances?
Common household substances span nearly the entire pH scale. Battery acid has a pH around 1, lemon juice is about pH 2, vinegar is pH 2.4-3.4, coffee ranges from pH 4.8 to 5.1, milk is approximately pH 6.5, and pure water is exactly pH 7.0. On the basic side, baking soda solution is about pH 8.3, household ammonia ranges from pH 11 to 12, and bleach is approximately pH 12.5-13.5. Drain cleaners containing sodium hydroxide can reach pH 14.
Why is pH important in everyday life?
pH is critical because it governs chemical reactions, biological processes, and material compatibility in nearly every field. Human blood pH must remain between 7.35 and 7.45 for survival; a deviation of just 0.3 units can be fatal. Swimming pool water requires pH 7.2-7.8 for effective chlorine disinfection. Soil pH determines which nutrients plants can absorb, with most vegetables growing best between pH 6.0 and 7.0. In food science, pH controls microbial growth, which is why pickling (low pH) preserves food. Even aquarium fish require specific pH ranges for their species.
What is the difference between pH and pOH?
pH measures the hydrogen ion concentration in a solution, while pOH measures the hydroxide ion concentration. They are mathematically complementary: pH + pOH = 14 at 25 degrees Celsius. An acidic solution has a low pH (high [H+]) and a correspondingly high pOH (low [OH-]). A basic solution has a high pH and low pOH. Neutral water has both pH and pOH equal to 7.0. While pH is far more commonly used in practice, pOH can be more convenient when working with bases, since it directly reflects the hydroxide concentration driving the reaction.
Can pH be negative or greater than 14?
Yes, pH values outside the 0-14 range are theoretically and practically possible. Concentrated strong acids can produce negative pH values because their hydrogen ion concentration exceeds 1 mol/L. For example, concentrated hydrochloric acid (about 12 M HCl) has a pH of approximately -1.1. Similarly, very concentrated sodium hydroxide solutions can exceed pH 14. However, the 0-14 range covers the vast majority of practical chemistry applications, and pH measurements at the extremes require specialized techniques because standard glass electrodes become unreliable.
How does temperature affect pH measurements?
Temperature significantly affects pH because it changes the water dissociation constant (Kw). At 25 degrees C, Kw = 1.0 x 10^-14 and pure water has pH 7.0. At 37 degrees C (body temperature), Kw increases to about 2.4 x 10^-14, making pure water's pH approximately 6.81 — still neutral, but not 7.0. At 60 degrees C, pure water has a pH of about 6.51. This means that a pH reading of 6.8 at 37 degrees C does not indicate acidity — it indicates neutrality at that temperature. Modern pH meters include automatic temperature compensation (ATC) to account for this effect.
What is a buffer solution and how does it maintain pH?
A buffer solution resists changes in pH when small amounts of acid or base are added. It consists of a weak acid and its conjugate base (e.g., acetic acid and sodium acetate), or a weak base and its conjugate acid. When acid is added, the conjugate base neutralizes it; when base is added, the weak acid neutralizes it. The most important biological buffer is the bicarbonate system in blood (H2CO3/HCO3-), which maintains blood pH near 7.4. Buffers work most effectively within one pH unit of the weak acid's pKa, and their capacity depends on the total concentration of the buffer components. You can use our Molarity Calculator to determine the concentrations needed for your buffer preparation.