pH Calculator
Calculate pH values from H+ concentration, pOH, or acid/base solutions. Includes visual pH scale and detailed results.
Frequently Asked Questions
pH is a measure of the acidity or basicity of a solution. It is calculated as the negative logarithm of the hydrogen ion concentration: pH = -log[H+]. A pH of 7 is neutral, values below 7 are acidic, and values above 7 are basic.
pH and pOH are related by the equation: pH + pOH = 14. To calculate pH from pOH, simply subtract the pOH value from 14. For example, if pOH is 5, then pH = 14 - 5 = 9.
Ka (acid dissociation constant) is a measure of the strength of an acid. For weak acids, we use Ka to calculate the concentration of H+ ions in solution, which then allows us to calculate the pH. The larger the Ka value, the stronger the acid.
For buffer solutions, pH can be calculated using the Henderson-Hasselbalch equation: pH = pKa + log([A-]/[HA]), where [A-] is the concentration of the conjugate base and [HA] is the concentration of the weak acid. This equation is particularly useful for solutions containing a weak acid and its conjugate base.
For strong acids, pH calculation is straightforward as they completely dissociate in water. The pH is simply -log[H+]. For weak acids, we need to use the Ka value and solve a quadratic equation to find the H+ concentration, as they only partially dissociate. This is why our calculator includes separate options for acid solutions.
Our pH calculator provides accurate results based on standard chemical equations and constants. However, for very precise laboratory work, we recommend verifying results with experimental measurements, as factors like temperature and ionic strength can affect pH values.
Let's calculate the pH of a 0.1M acetic acid solution (Ka = 1.8 × 10^-5). Using our calculator, input the concentration as 0.1 and Ka as 1.8e-5. The calculator will solve the quadratic equation to find the H+ concentration and then calculate the pH. The result should be approximately 2.87, showing that this is a weak acid solution.