Acid Dissociation Constant

Q: What is Ka of an acid?

Ka of an acid tells us if the acid is a strong acid or a weak acid. Strong acids have high Ka value while weak acids have low Ka value.

Chemical Analysis

Chemical Reactions

Chemistry Branches

History of Chemistry

Inorganic Chemistry

Ionic and Molecular Compounds

Kinetics

Making Measurements

Nuclear Chemistry

Organic Chemistry

Physical Chemistry

The Earths Atmosphere

TABLE OF CONTENTS :

TABLE OF CONTENTS

Lerne mit deinen Freunden und bleibe auf dem richtigen Kurs mit deinen persönlichen Lernstatistiken

Jetzt kostenlos anmelden

Nie wieder prokastinieren mit unseren Lernerinnerungen.

Jetzt kostenlos anmelden

Did you know that most food items like apples and lemons contain naturally occurring acids? Have you ever wondered what makes an acid 'strong' or 'weak'? You will learn that, here.

Acid dissociation constant definition

We know that an acid dissociates into an H⁺ ion and the corresponding anion. We also know that the reaction of acid and water is an ionisation reaction. These reactions are reversible and reach an equilibrium state where the rates of forward and backward reactions are the same. The equilibrium constant calculated for an acid solution at equilibrium is called the acid dissociation constant.

Acid dissociation constant is a term used to quantitatively measure the strength of an acid in a solution.

You must remember how to find the equilibrium constant of a chemical reaction. The dissociation constant is nothing else but the equilibrium constant of the ionisation reaction of any acid in a solution. For any acid HA, the equilibrium reaction for the solution of HA would be:

$HA ⇌ H^{+} + A^{-}$

Here, H⁺ is a Hydron, or Hydrogen ion, which gives any acid its acidic nature. The anion formed (A^-) after HA releases H⁺. This is called the conjugate base of the acid HA. As you might remember, a solution is said to be in a state of equilibrium when the rate of the forward reaction is equal to the rate of the backward reaction. The acid dissociation constant is calculated at equilibrium. It is denoted by K_a.

$K_{e q} = K_{a} = \frac{[H^{+}] [A^{-}]}{[H A]}$

Remember that square brackets refer to the concentration of the molecule/ion.

Looking at the formula for K_a, it might seem that it depends on the concentration of the acid/ions in the solution, but that is not true. The value of K_afor an acid is fixed, and does not depend on the concentration of the acid. K_a varies only with temperature.

Usually, the value of K_a is very large, so a logarithm of K_a is taken. This is called pK_a.

The acid dissociation constant is sometimes also called acid ionisation constant, or acidity constant.

pH of a solution

You might have come across the term pH while studying Chemistry. pH is a measure of the acidic strength of a solution. For the sake of understanding, it can be considered as a scaled quantity of concentration of H⁺ ions in a solution.

$p H = - \log_{10} ([H^{+}])$ $= - \log_{10} [H^{+}]$

Find out the pH of a solution with H⁺ ion concentration equal to 10^-7.

Solution

Given:

[H⁺] = 10^-7

we know that:

pH = -log₁₀[H⁺]

∴ pH = -log₁₀(10^-7)

pH = 7

Common exam questions might give you information about the pH of a solution and ask you to find the K_a of a particular acid. It may be the other way round; they'll give you information about K_a of an acid, and ask you to find the pH of a solution of that acid.

The pH of a solution can range from 0 to 14. Solutions with pH less than 7 are acidic, while those with a pH above 7 are basic. The characteristic sour taste of lemons is due to the presence of citric acid, giving it a pH of ≈ 2. The black coffee you have in the morning has a pH of ≈ 5. Saliva has a pH of ≈ 6.6. On the basic side, the baking soda you have in the kitchen has a pH of 9.5. Liquid drain cleaners have a pH of ≈ 14 (very strong, which is why we use them with care; they can hurt skin).

Did you know that the pH of human blood ranges between 7.35 and 7.45; i.e. it is slightly on the basic side. This is necessary to balance out the human body's various metabolic processes which produce acids.

It's easy to confuse pH and pK_a when you're learning about them for the first time. For the sake of clarity, pK_a is for an acid, and pH is for a solution. pK_a depends on the temperature of the acid/solution, while pH depends on the concentration of H⁺ ions in the solution.

For the same amount of acids added to 2 solutions, the pH of the solution with the stronger acid (higher pK_a) will be lower (more acidic).

Acid dissociation constant of HCl

A common example of a strong acid is hydrochloric acid (HCl). For HCl, the ionisation reaction would be written like this:

$HCl + H_{2} O ⇌ H_{3} O^{+} + {Cl}^{-}$

The equilibrium constant for the above reaction would be calculated as:

$K_{a} = \frac{[H_{3} O^{+}] [C l^{-}]}{[H C l] [H_{2} O]}$

The H⁺ ions are transferred to H₂O and form hydronium ions (H₃O⁺). This is because free H⁺ ions do not exist in an aqueous solution. For simplicity, we can replace [H₃O⁺] with [H⁺]. Also, [H₂O] is equal to 1 as the concentration of water is 1. Hence, the equation for K_a becomes:

$K_{a} = \frac{[H^{+}] [C l^{-}]}{[H C l]}$

The addition of H⁺ ion to a water molecule is called protonation of water. Can you guess why? (Hint: it's in the name, 'protonation').

HCl is a strong acid. This means that it dissociates completely into H⁺ and Cl^- ions in solution. For example, if 1 mol of HCl is added to 1 L of water, it will give out 1 mol of H⁺ ions and 1 mol of Cl^- ions.

This can be represented with the help of an ICE table (Initial, Change, Equilibrium), where we can write the concentrations of the reactants and the products of the reaction at the initial stage (Initial), the change in their concentration from initial stage till equilibrium (Change), and their final concentrations at equilibrium (Equilibrium).

	[HCl]	[H⁺]	[Cl^-]
Initial	1	0	0
Change	-1	1	1
Equilibrium	negligible	1	1

In the ICE table for the ionisation reaction of HCl, we have written the concentrations of the reactants and the products at the initial stage, and the final stage. We can now substitute these values in the formula for K_a for HCl.

$K_{a} = \frac{1 \cdot 1}{(v e r y s m a l l n u m b e r)}$

There is a very small number in the denominator, as there is always a very small amount of undissociated acid at equilibrium. Therefore, the value of K_a for strong acids is very large. K_a for HCl is very high, at 1.3 x 10⁶.

Examples of other strong acids include:

Perchloric acid (Superacid) - HClO₄ (K_a > 1.0 x 10⁹)
Nitric acid - HNO₃ (K_a = 2.4 x 10¹)
Hydrobromic acid - HBr (K_a = 1.0 x 10⁹)
Hydroiodic acid - HI (K_a = 3.2 x 10⁹)
Sulfuric acid - H₂SO₄ (K_a = 1.0 x 10³)

In this subsection, we learnt that strong acids are so-called because they dissociate completely into H⁺ ions and the corresponding conjugate bases.

But what about acids that cannot dissociate completely in a solution? Those acids are called weak acids.

Determination of the dissociation constant of a weak acid

Let us consider an acid HB. The ionisation reaction equation will be written as:

$H B ⇌ H^{+} + B^{-}$

Consider HB as a weak acid. Weak acids do not dissociate completely, i.e. if 1 mol of HB is added to water, at equilibrium the solution will have less than 1 mol of H⁺ and B^-, and will also contain a certain concentration of undissociated HB at equilibrium (unlike the case of strong acid, which only contained H⁺ and the conjugate base at equilibrium). Due to incomplete dissociation/ionisation of the weak acid, the solution is not as acidic as it would be if an equal quantity of a strong acid was added to the solution.

Let us take the example of acetic acid (CH₃COOH), a weak acid. Consider adding 1 mol of acetic acid to 1 L of water. The concentration of acetic acid would be 1 mol/L or 1M. The equation of the reaction can be written as:

$C H_{3} C O O H + H_{2} O ⇌ H_{3} O^{+} + C H_{3} C O O^{-}$

At equilibrium, the K_a for this equation can be written as:

$K_{a} = \frac{[H^{+}] [C H_{3} C O O^{-}]}{[C H_{3} C O O H]}$

Acetic acid is an organic acid that is the primary acid in vinegar. Besides that, it is also found in apples, strawberries, and grapes.

Just as we drew an ICE table for the ionisation of HCl, we can draw the ICE table for the ionisation of acetic acid. This will help us better understand the concentrations of the reactants and the products at different stages of the reaction.

	[CH₃COOH]	[H⁺]	[CH₃COOH^-]
Initial	1	0	0
Change	-x	+x	+x
Equilibrium	1-x	x	x

Weak acids do not dissociate completely, which is why the change in concentration of CH₃COOH is -x. We use the variable x because we do not know how much the acid will dissociate, and hence how many H⁺ ions or CH₃COO^- ions there will be in the solution.

You know that acetic acid is a weak acid. You have also understood that weak acids do not ionise completely in a solution. Therefore, in the equation of K_afor acetic acid, the concentration of the H⁺ ions and the conjugate base is less than 1M, while the concentration of the reactants is a significant value, as we know there is still undissociated acid at equilibrium (unlike the case of strong acids). So, for weak acids, the numerator while calculating K_a is small, whilst the denominator is big. This implies that the value of K_a for weak acids is small. K_a for acetic acid is 1.8 × 10⁻⁵.

Examples of other weak acids are:

Methanoic acid - HCOOH (K_a = 1.78 × 10⁻⁴)
Benzoic acid - C₆H₅COOH (K_a = 6.3 × 10⁻⁵)
Hypochlorous acid - HClO (K_a = 2.9 × 10⁻⁸)
Hydrocyanic acid - HCN (K_a = 6.2 × 10⁻¹⁰)

1 mol of acetic acid is added to 1 L of water. Calculate the pH of the solution at equilibrium.

Solution

Given:

CH₃COOH = 1 mol

Water = 1L

∴ [CH₃COOH ] = $\frac{1 m o l}{1 L}$ = 1M

We know that,

$p H = - \log_{10} [H^{+}]$

To calculate pH, we need to know the concentration of H⁺ ions in the solution at equilibrium. To do this, we shall draw an ICE table for the reaction.

$C H_{3} C O O H ⇌ H^{+} + C H_{3} C O O^{-}$

	[CH₃COOH]	[H⁺]	[CH₃COO^-]
Initial	1	0	0
Change	-x	+x	+x
Equilibrium	1-x	x	x

$K_{a} = \frac{[H^{+}] [C H_{3} C O O^{-}]}{[C H_{3} C O O H]}$

Since acetic acid is a weak acid and so it wouldn't have dissociated very much, we can assume that its equilibrium concentration is approximately that of the original concentration.

$K a = \frac{[H^{+}] [C H_{3} C O O^{-}]}{[C H_{3} C O O H]} ~ \frac{{[H^{+}]}^{2}}{{[C H_{3} C O O H]}_{i n i t i a l}}$

Note that we can replace [CH₃COO^-] with [H⁺] since both are equal.

$\Rightarrow [H^{+}] = \sqrt{K a . [C H_{3}} C O O H] = 0.004242 m o l / l$

Now that we know the concentration of H⁺ ions in solution, we can calculate the pH of the solution by taking its log.

$[H^{+}] = 0.004242 ∴ p H = - \log_{10} (0.004242) \Rightarrow p H = 2.37$

Acid Dissociation Constant - Key takeaways

Acid dissociation constant (K_a) is used to quantitatively measure the strength of an acid.
Acid dissociation constant is nothing but the equilibrium constant calculated for the ionisation of an acid.
An acid HA dissociates into the hydrogen ion (H⁺) and the conjugate base of that acid (A^-).
The value of K_a is usually very large or small. That is why it is scaled down by taking a negative log; pK_a = -log₁₀K_a.
pH determines how acidic a solution is. pH = -log₁₀[H⁺].
pH ranges from 0-14. 7 is neutral. Below 7 is acidic, above 7 is basic.
Strong acids are those acids that dissociate completely. Weak acids are those that do not dissociate completely.
Strong acids have a high K_a value. Weak acids have a small K_a value.
An acid dissociation constant table (or ICE table) can help understand the concentrations of the products and reactants at the initial stage of the reaction, and at equilibrium.

Frequently Asked Questions about Acid Dissociation Constant

Acid dissociation constant is a term used to quantitatively measure the strength of an acid.

Acid dissociation constant is denoted by Ka. It is the equilibrium constant of the ionisation reaction of an acid.

K_a of an acid tells us if the acid is a strong acid or a weak acid. Strong acids have high K_a value while weak acids have low K_a value.

Yes. Strong acids dissociate completely, which results in a higher concentration of products than reactants at equilibrium, which is why they have a high dissociation constant.

Yes. Equilibrium constant (or acid dissociation constant) for weak acids applies to weak acids.

Final Acid Dissociation Constant Quiz

Acid Dissociation Constant Quiz - Teste dein Wissen

Question

What kind of chemical reaction is an acid with water?

Show answer

Answer

Ionization reaction.

Show question

Question

What is the range of pH?

Show answer

Answer

From 0 (most acidic) to 14 (most basic).

Show question

Question

What is the pH of drinking water?

Show answer

Answer

6.5 - 8.5

Show question

Question

When an acid dissociates it gives a hydron and an anion. What is that anion called?

Show answer

Answer

The conjugate base of that acid.

Show question

Question

What is the source of acetic acid in our daily lives?

Show answer

Answer

Vinegar, apples, strawberries, grapes.

Show question

Question

What acid is present in lemons and oranges?

Show answer

Answer

Citric acid.

Show question

Question

Are washing detergents acidic or basic?

Show answer

Answer

Acidic

Show question

Question

What is an ICE table used for?

Show answer

Answer

ICE table is useful for determining the concentrations of the products and reactants at the initial stage, the change in their concentrations over the course of reaction, and their concentrations at equilibrium.

Show question

Question

What is a compound with pH lower than 7 called?

Show answer

Answer

Acidic

Show question

Question

What is a compound with pH higher than 7 called?

Show answer

Answer

Basic or alkali.

Show question

Question

Which of the following is NOT a name for an H⁺ ion?

Show answer

Answer

Hydrogen ion

Show question

Question

What is a hydronium ion?

Show answer

Answer

A protonated water molecule forms a hydronium ion, H₃O⁺.

Show question

Question

Does K_a depend on the concentration of the acid?

Show answer

Answer

No. It only depends on the temperature.

Show question

Question

How is pH calculated?

Show answer

Answer

pH = -log([H⁺])

Show question

Question

Does pH depend on the concentration of the acid?

Show answer

Answer

Yes. Higher the concentration of the acid, more H⁺ ions will be present in the solution at equilibrium, more acidic will be the solution and lower the pH will be.

Show question

More about Acid Dissociation Constant

How would you like to learn this content?

Creating flashcards

Studying with content from your peer

Taking a short quiz

94% of StudySmarter users achieve better grades.

How would you like to learn this content?

Creating flashcards

Studying with content from your peer

Taking a short quiz

Free chemistry cheat sheet!

Everything you need to know on . A perfect summary so you can easily remember everything.

Email Address*

Select your language

Acid Dissociation Constant

Acid Dissociation Constant

Acid Dissociation Constant

Acid dissociation constant definition

pH of a solution

Acid dissociation constant of HCl

Determination of the dissociation constant of a weak acid

Acid Dissociation Constant - Key takeaways

Frequently Asked Questions about Acid Dissociation Constant

What is acid dissociation constant?

How do you write an acid dissociation constant?

What is Ka of an acid?

Do strong acids posses a high acid dissociation constant?

Does equilibrium constant apply to dissociation of weak acid?

Final Acid Dissociation Constant Quiz

Acid Dissociation Constant Quiz - Teste dein Wissen

More explanations about Physical Chemistry

Discover the right content for your subjects

Biology

Business Studies

Combined Science

Computer Science

Economics

English

English Literature

Environmental Science

Geography

History

Human Geography

Law

Macroeconomics

Marketing

Math

Microeconomics

Physics

Politics

Psychology

Sociology

No need to cheat if you have everything you need to succeed! Packed into one app!

Study Plan

Quizzes

Flashcards

Notes

Study Sets

Documents

Study Analytics

Weekly Goals

Smart Reminders

Rewards

Magic Marker

Smart Formatting

Join millions of people in learning anywhere, anytime - every day

This is still free to read, it's not a paywall.

You need to register to keep reading

This is still free to read, it's not a paywall.

You need to register to keep reading

Start learning with Vaia, the only learning app you need.

StudySmarter bietet alles, was du für deinen Lernerfolg brauchst - in einer App!