pKa Calculator | Find pKa from pH or Ka (2024)

This pK_a calculator will help you determine the pK_a value in two ways: from a specific pH with the Henderson-Hasselbalch equation or from the acid dissociation constant (K_a).

In this article, you will find information on:

• The definition of pK_a;
• How to use the pK_a table;
• Relationship between pK_a and pH;
• Relationship between pK_a and K_a; and
• Useful examples are also provided to help you calculate pK_a values like a pro!

You have probably seen the term pKa before from your chemistry class in high school🧑‍🔬​. What is it exactly? Do you recall?

It's very simple! The pK_a determines how weak or strong an acid is. To be more precise, pK_a tells you how strongly the Brønsted acid holds on a given proton(H⁺). It allows you to predict how each acid and base solution will react in a specific experimental setting.

Interpretation of pK_a

The lower the pK_a, the stronger the acid. This means:

• The H⁺ is held more loosely by the acid; and
• The acid can give up on H⁺ more easily.

The higher the pK_a, the weaker the acid. This means:

• The H⁺ is held more tightly by the acid; and
• The acid does not easily donate a H⁺.

Before going through equations and calculations of pK_a, here is the easiest way to find the pK_a of a compound — by using the pK_a table.

This table can be used when you are trying to make a buffer or carry out a reaction from scratch. It guides the selection of acids and bases in a reaction, since knowing the pK_a values of each compound will predict their reactive behavior 🧪​. So, if you don't know the pH or the K_a of your compounds, check it out!

You can find a shortened version of the pK_a table with the most common functional groups used in basic chemistry below:

Henderson-Hasselbalch equation

The pK_a calculator is based on the well-known Henderson-Hasselbalch equation, providing the relationship between pH and pK_a.

where:

• $\mathrm{A^-}$A− — Molar concentration of the conjugate base; and
• $\mathrm{HA}$HA — Molar concentration of the weak acid.

Relationship of pK_a and pH

• If $\small\mathrm{[HA] = [A^-]}$[HA]=[A−], then $\mathrm{log\large\frac{[A^-]}{[HA]} \small = 0}$log[HA][A−]​=0:

  When molar concentrations of weak acid and conjugate base are the same, the logarithm is exactly 0. This means that $\small\mathrm{pH = pK_a}$pH=pKa​.

• If $\small\mathrm{[HA] > [A^-]}$[HA]>[A−], then $\mathrm{log\large\frac{[A^-]}{[HA]}\small < 0}$log[HA][A−]​<0:

  When the molar concentration of the weak acid is higher than that of the conjugate base, the logarithm is negative. Thus, $\small\mathrm{pH < pK_a}$pH<pKa​.

• If $\small\mathrm{[HA] < [A^-]}$[HA]<[A−], then $\mathrm{log\large\frac{[A^-]}{[HA]}\small > 0}$log[HA][A−]​>0:

  When the molar concentration of the conjugate base is higher than that of the weak acid, the logarithm is positive. Thus, $\small\mathrm{pH > pK_a}$pH>pKa​.

Here is another terminology to recall from your chemistry lecture🤓​ — the acid dissociation constant (K_a), also known as acid ionization constant.

K_a is a constant value measured at equilibrium, indicating how acids dissociate in a solution. The higher K_a values, the stronger the acid and the easier the dissociation (H⁺ donating) from the other components.

Relationship of pK_a and K_a

pK_a is negatively correlated to K_a, meaning that if one value increases⬆️, the other value decreases​⬇️. ​ Basically, K_a is simply the logarithm of pK_a:

Acid dissociation equation

This equation explains the dissociation of an acid at equilibrium, giving hydronium ions (H⁺) and conjugate base (A^-) as products. Knowing how to write the reaction in this form is very important to make sure you know which compound is the nominator or denominator when using the Henderson-Hasselbalch equation to calculate pK_a or pH.

✏️​Practice makes progress! Explore the provided examples below to improve your knowledge of pK_a even more.

Calculating pK_a from pH

Example 1 — Calculate the pK_a of a solution containing $\small0.1\ \mathrm{M}$0.1M of acetic acid and $\small0.01\ \mathrm{M}$0.01M of acetate ion. Note that $\small\mathrm{pH} = 4.8$pH=4.8.

1. Write the acid dissociation equation of this reaction to identify the weak acid and conjugate base:

Example 2 — Calculate the pK_a of a solution containing $\small0.75\ \mathrm{M}$0.75M of lactic acid and $\small0.25\ \mathrm{M}$0.25M of sodium lactate. Note that $\small\mathrm{pH} = 3.38$pH=3.38.

1. Write the acid dissociation equation of this reaction to identify the weak acid and conjugate base:

Calculating pK_a from K_a

Example 3 — If $\small\mathrm{K_a} = 1.5 × 10^{-5}$Ka​=1.5×10−5, how much is pK_a?

K_a is the acid dissociation constant, which determines how strong an acid is by its ability to dissociate in a solution. pK_a, on the other hand, is basically the negative log of K_a. Both of these values can determine how strong or weak an acid is.

No, pH and pK_a are two different things. pH is a scale that measures the presence of H⁺ ions in a solution, making it acidic, neutral, or basic. As for pK_a, it tells us how strong an acid is.

Use the relationship between pK_a and the acid dissociation constant (K_a): pK_a = -log₁₀[K_a]. The equation can also be reverted in case pK_a is given to calculate K_a: K_a = 10^-pK_a.