Last updated: Apr 28, 2024

pKa Calculator

Q: What is the difference between pKa and Ka?

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. Read more

Q: Are pH and pKa the same?

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. Read more

Q: How do I calculate pKa from pH?

Use the Henderson-Hasselbalch equation to calculate pK a from pH: pH = pK a + log 10 [A - ]/[HA] where [A - ] is the conjugate base and [HA] is the weak acid. The pH and molar concentrations of the acid and base must be known to calculate pK a . Read more

Q: How do I calculate pKa from Ka?

Use the relationship between pK a and the acid dissociation constant (K a ): pK a = -log 10 [K a ] . The equation can also be reverted in case pK a is given to calculate K a : K a = 10 -pK a . Read more

Q: How do I calculate pKa of 0.58 M sodium acetate and 1.0 M acetic acid?

To calculate pK a from pH: Apply the Henderson-Hasselbalch equation: pH = pK a + log 10 [A - ]/[HA] The conjugate base, sodium acetate, is [A - ]: [C 2 H 3 NaO 2 ] The weak acid, acetic acid, is [HA]: [CH 3 COOH]. Thus, considering pH = 4.5, we can calculate the pK a as follows: pH = pK a + log[C 2 H 3 NaO 2 ] / [CH 3 COOH] 4.5 = pK a + log[0.58]/[1.0] pK a = -0.236 − 4.5 = 4.737 Congratulations! Now, you can check your answer using Omni Calculator's pK a calculator. Read more

Q: How much is pKa of acetic acid if the Ka is 1.8×10⁻⁵?

The pK a of acetic acid is 4.745 . Given the K a of 1.8 × 10 -5 , pK a can be calculated as follows: pK a = -log 10 [K a ] pK a = -(-4.745) = 4.745 Read more

Table of contents

What is pKa?pKa table pKa and pH — How to calculate pKa from pH?pKa and Ka — How to calculate pKa from Ka?Examples of pKa calculation FAQs

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!

What is pKa?

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⁺.

💡 If you are looking for a way to calculate the pH of a solution in your experiment ⚗️, knowing the concentration of the acid or base is very important. You might find Omni's pH calculator and concentration calculator helpful!

pKa table

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:

Functional group	Formula	pK_a
Hydroiodic acid	HI	-10
Hydrobromic acid	HBr	-9
Hydrochloric acid	HCl	-6
Sulfuric acid	H₂SO₄	-3
Hydronium ion	H₃O⁺	-1.7
Sulfonic acids	R–SO₃H	-1
Hydrofluoric acid	HF	3.2
Carboxylic acid	R–COOH	4
Protonated amines	R–NH₂	9-11
Thiols	R–SH	13
Malonates	CH₂(COOH)₂	13
Water	H₂O	14
Alcohol	CH₃CH₂OH	17
Ketone / Aldehyde	R–CH₂O / R–CHO	20-24
Nitrile	R–C≡N	25
Ester	R–COO–R'	25
Alkyne	R–C≡C–R'	25
Amines	R–NH₂	~35
Hydrogen	H	36
Alkene	R–C=C–R'	~43
Alkane	C_nH_2n+2	~50

For more visualization, check out Master Organic Chemistry website, which provides a clear structural illustration of each functional group and the conjugate base.

pKa and pH — How to calculate pKa from pH?

Henderson-Hasselbalch equation

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

\text{pH} = \mathrm{pK_a} + \mathrm{log_{10}\frac{[A^-]}{[HA]}}

where:

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

Relationship of pK_a and pH

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

When molar concentrations of weak acid and conjugate base are the same, the logarithm is exactly 0. This means that $\small\mathrm{pH = pK_a}$ .
If $\small\mathrm{[HA] > [A^-]}$ , then $\mathrm{log\large\frac{[A^-]}{[HA]}\small < 0}$ :

When the molar concentration of the weak acid is higher than that of the conjugate base, the logarithm is negative. Thus, $\small\mathrm{pH < pK_a}$ .
If $\small\mathrm{[HA] < [A^-]}$ , then $\mathrm{log\large\frac{[A^-]}{[HA]}\small > 0}$ :

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

💡 Did you know that the buffer capacity increases as the value of pH and pK_a are closer together? This allows a buffer to maintain its pH range despite the addition of a stronger acid or base. Check our buffer capacity calculator for more information! 💬

pKa and Ka — How to calculate pKa from Ka?

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:

\rm {pK_a} = {-log_{10}}[{K_a}]

You can also write the equation as follows in case you want to calculate an unknown K_a from pK_a:

\rm {K_a} = 10^{-pK_a}

If you are interested in knowing more about what is a logarithm and how to solve the log of a value, the log calculator will be very helpful!

💡 Unlike pH, K_a does not vary with the concentration, but it does vary with temperature changes. That means K_a values of acids are usually fixed.

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.

\rm HA \rightleftharpoons H^{+} + A^{-}

Hence, K_a can be represented by the concentration of products to the concentration of the reactant:

\rm K_a = \small\frac{[H^+][A^-]}{HA}

Examples of pKa calculation

✏️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 $\small0.1\ \mathrm{M}$ of acetic acid and $\small0.01\ \mathrm{M}$ of acetate ion. Note that $\small\mathrm{pH} = 4.8$ .

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

\footnotesize \begin{align*} \rm HA\ &\rm\rightleftharpoons H^{+} + A^{-} \\ \rm CH_3COOH\ &\rm\rightleftharpoons H^{+}_3 + CH_3COO^{-} \end{align*}

2. Calculate the pK_a using the Henderson-Hasselbalch equation:

\footnotesize \begin{align*} \rm \text{pH}\ &\rm= \mathrm{pK_a} + \mathrm{log_{10}\frac{[CH_3COO^-]}{[CH_3COOH]}} \\[1.5em] -\rm{pK_a}\ &\rm= \mathrm{log_{10}\frac{[CH_3COO^-]}{[CH_3COOH]}} - \mathrm{pH} \\[1.5em] -\rm{pK_a}\ &\rm= \mathrm{log_{10}\frac{[0.01]}{[0.1]}} - 4.8 \\[1.5em] -\rm{pK_a}\ &\rm= -1 - 4.8 \\[.7em] -\rm{pK_a}\ &\rm= -5.8 \\[.7em] \rm{pK_a}\ &\rm= 5.8 \end{align*}

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

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

\footnotesize \begin{align*} \rm HA\ &\rm\rightleftharpoons H^{+} + A^{-} \\ \rm HC_3H_5O_3\ &\rm\rightleftharpoons H^{+} + C_3H_5O^{-}_3 \end{align*}

2. Calculate the pK_a using the Henderson-Hasselbalch equation:

\footnotesize \begin{align*} \text{pH}\ &\rm= \rm{pK_a} + \mathrm{log_{10}\frac{[C_3H_5O_3^-]}{[HC_3H_5O_3]}} \\[1.5em] -\rm{pK_a}\ &\rm= \rm{log_{10}\frac{[C_3H_5O_3^-]}{[HC_3H_5O_3]}} - \rm{pH} \\[1.5em] -\rm{pK_a}\ &\rm= \rm{log_{10}\frac{[0.25]}{[0.75]}} - 3.38 \\[1.5em] -\rm{pK_a}\ &\rm= -0.047 - 3.38 \\[.7em] -\rm{pK_a}\ &\rm= -3.85 \\[.7em] \rm{pK_a}\ &\rm= 3.85 \end{align*}

Phew! 😮‍💨 That was a bit tiring, wasn't it? Let Omni Calculator's pK_a calculator make things a lot easier for you — easily fill in your given pH, conjugate base concentration and weak acid concentration in the section pK_a from pH.

Calculating pK_a from K_a

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

\footnotesize \begin{align*} \mathrm{pK_a} &= \mathrm{-log_{10}}[\mathrm{K_a}] \\ &= \mathrm{-log_{10}}[\mathrm1.5 × 10^{-5}] \\ &= 3.824 \end{align*}

Example 4 — If $\small\mathrm{K_a} = 6.8 × 10^{-10}$ , how much is pK_a?

\footnotesize \begin{align*} \mathrm{pK_a} &= \mathrm{-log_{10}}[\mathrm{K_a}] \\ &= \mathrm{-log_{10}}[\mathrm 6.8 × 10^{-10}] \\ &= 9.162 \end{align*}

Otherwise, feel free to input the K_a value in the section pK_a from K_a in our pK_a calculator and get your result in one second!

FAQs

What is the difference between pKa and Ka?

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.

Are pH and pKa the same?

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.

How do I calculate pKa from pH?

Use the Henderson-Hasselbalch equation to calculate pK_a from pH:

pH = pK_a + log₁₀[A^-]/[HA]

where [A^-] is the conjugate base and [HA] is the weak acid. The pH and molar concentrations of the acid and base must be known to calculate pK_a.

How do I calculate pKa from Ka?

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.

How do I calculate pKa of 0.58 M sodium acetate and 1.0 M acetic acid?

To calculate pK_a from pH:

Apply the Henderson-Hasselbalch equation:

pH = pK_a + log₁₀[A^-]/[HA]
The conjugate base, sodium acetate, is [A^-]: [C₂H₃NaO₂]
The weak acid, acetic acid, is [HA]: [CH₃COOH].
Thus, considering pH = 4.5, we can calculate the pK_a as follows:

pH = pK_a + log[C₂H₃NaO₂] / [CH₃COOH]

4.5 = pK_a + log[0.58]/[1.0]

pK_a = -0.236 − 4.5 = 4.737
Congratulations! Now, you can check your answer using Omni Calculator's pK_a calculator.

How much is pKa of acetic acid if the Ka is 1.8×10⁻⁵?

The pK_a of acetic acid is 4.745. Given the K_a of 1.8 × 10^-5, pK_a can be calculated as follows:

pK_a = -log₁₀[K_a]
pK_a = -(-4.745) = 4.745