What is Acid Dissociation in Water?

Ques: What is Acid Dissociation in Water?

Sol: Solvated hydroxide and solvated protons are in equilibrium with liquid water. To define the concentration of these solvated ions in water, we usually use a value related to the equilibrium constant. The Kw constant for water dissociation is 1 x 10-14.

For acidic and basic solutions, pH and pOH are critical values. (The log base 10 of the hydrogen ion concentration and the log base 10 of the hydroxide ion concentration, respectively.)

pH = -log [H+]

pOH = -log [HO]

14 = pH + pOH

The LeChateliers principle states that adding one of the products to an equilibrium system causes the equilibrium to shift towards reactants. Acids and bases dissolve in water and suppress water dissociation by increasing the concentration of one of the products of water self-ionization, either protons or hydroxide ions.

Acid and base solutions in water are typically described using pH and pOH. The concentration of solvated protons equals that of solvated hydroxide anions in pure water, and the pH is 7. The pH of acidic solutions is lower, whereas the pH of basic solutions is higher.

Definitions of Acid Dissociation Constant (Ka)

The concept of pKa is based on the thermodynamics of the dissociation reaction. The acid dissociation constant, Ka, is related to the standard Gibbs free energy change for the reaction. The pKa value is the negative logarithm of the Ka value and is a measure of the acidity of a compound.

The pKa value changes with temperature, which can be explained using Le Châtelier’s principle. For an endothermic reaction, an increase in temperature increases Ka and lowers pKa, while for an exothermic reaction, an increase in temperature increases Ka and raises pKa.

In the Arrhenius definition, an acid is a substance that dissociates in water to produce H+ ions. This can be represented by the reaction:

HA ⇌ A- + H+

However, because H+ ions are hydrated to form H3O+ ions, the reaction is more accurately represented as:

HA + H2O ⇌ A- + H3O+

This reaction can also be viewed as a proton transfer reaction, where the acid donates a proton to the base to form the conjugate base and conjugate acid. This is known as the Bronsted-Lowry definition of acids and bases.

The pKa value is a measure of the acidity of a compound, which is related to the acid dissociation constant and the thermodynamics of the dissociation reaction. The Arrhenius and Bronsted-Lowry definitions describe the behavior of acids and bases in aqueous solutions.

Strong Acids and Water Dissociation

Strong acids are acids that completely dissociate in water, meaning that all of their acid molecules react with water to produce hydronium ions (H3O+) and the conjugate base. Examples of strong acids include hydrochloric acid (HCl), sulfuric acid (H2SO4), and nitric acid (HNO3).

Water itself also undergoes a dissociation reaction in which a small fraction of water molecules dissociate into hydronium ions and hydroxide ions (OH-):

H2O ⇌ H+ + OH-

The equilibrium constant for this reaction is the ion product constant for water, Kw, which is equal to the product of the concentrations of H+ and OH- ions in water at equilibrium:

Kw = [H+][OH-]

At 25°C, Kw has a value of 1.0 x 10^-14 mol^2/L^2. This means that if the concentration of H+ ions in water is increased, the concentration of OH- ions must decrease in order to maintain a constant value of Kw.

The pH of a solution is a measure of the concentration of H+ ions in the solution. It is defined as the negative logarithm of the H+ ion concentration:

pH = -log[H+]

Similarly, the pOH of a solution is a measure of the concentration of OH- ions in the solution and is defined as the negative logarithm of the OH- ion concentration:

pOH = -log[OH-]

The pH and pOH of a solution are related by the equation:

pH + pOH = 14

This means that as the pH of a solution decreases (i.e. as the H+ ion concentration increases), the pOH must increase (i.e. the OH- ion concentration must decrease) in order to maintain a constant value of 14 for pH + pOH.

Weak Acids and Water Dissociation

For weak acids, there is an equivalent situation where the [H+] due to acid dissociation is comparable to the [H+] due to water dissociation (1.0×10^-7 M). Water dissociation can contribute to the pH of a weak acid solution if the acid is extremely dilute, very weak, or dilute and weak, unlike strong acids, where it is important only when very dilute acid is involved.

In an aqueous solution, weak acids dissociate only partially. The table below shows some of the weak acids:


Ka of Weak Acids
Name Formula Ka pKa
acetic HC2H3O2 1.8 x 10-5 4.7
ascorbic (I) H2C6H6O6 7.9 x 10-5 4.1
ascorbic (II) HC6H6O6 1.6 x 10-12 11.8
benzoic HC7H5O2 6.4 x 10-5 4.2
boric (I) H3BO3 5.4 x 10-10 9.3
boric (II) H2BO3 1.8 x 10-13 12.7
boric (III) HBO32- 1.6 x 10-14 13.8
carbonic (I) H2CO3 4.5 x 10-7 6.3
carbonic (II) HCO3 4.7 x 10-11 10.3
citric (I) H3C6H5O7 3.2 x 10-7 6.5
citric (II) H2C6H5O7 1.7 x 105 4.8
citric (III) HC6H5O72- 4.1 x 10-7 6.4
formic HCHO2 1.8 x 10-4 3.7
hydrazidic HN3 1.9 x 10-5 4.7
hydrocyanic HCN 6.2 x 10-10 9.2
hydrofluoric HF 6.3 x 10-4 3.2
hydrogen peroxide H2O2 2.4 x 10-12 11.6
hydrogen sulfate ion HSO4 1.2 x 10-2 1.9
hypochlorous HOCl 3.5 x 10-8 7.5
lactic HC3H5O3 8.3 x 10-4 3.1
nitrous HNO2 4.0 x 10-4 3.4
oxalic (I) H2C2O4 5.8 x 10-2 1.2
oxalic (II) HC2O4 6.5 x 10-5 4.2
phenol HOC6H5 1.6 x 10-10 9.8
propanic HC3H5O2 1.3 x 10-5 4.9
sulfurous (I) H2SO3 1.4 x 10-2 1.85
sulfurous (II) HSO3 6.3 x 10-8 7.2
uric HC5H3N4O3 1.3 x 10-4 3.9


Acidity in nonaqueous solutions

In an acidic solvent, the ionization of an acidic molecule is typically less than in water because the solvent molecules can compete with the acid molecules for the available hydrogen ions. However, in a protic solvent, the solvent molecules can form hydrogen bonds with the acid molecules, stabilizing the resulting ions and promoting acid ionization. The strength of the solvent’s Lewis base character can also affect the ionization of an acidic molecule, as it can either stabilize or destabilize the resulting ions. And finally, the high dielectric constant of some solvents can increase the solubility of ionic species and promote the ionization of acidic molecules.


Acid Dissociation in Water FAQs

Acid dissociation in water refers to the process in which an acid molecule breaks apart or dissociates into ions when dissolved in water. This dissociation results in the release of hydrogen ions (H+) into the solution.
When an acid is dissolved in water, the polar water molecules surround the acid molecules. The water molecules effectively pull the acidic hydrogen atoms away from the acid molecule, resulting in the formation of hydronium ions (H3O+) or hydrogen ions (H+) and the corresponding conjugate base. This dissociation occurs due to the transfer of protons from the acid to water.
The acid dissociation constant, also known as Ka, is a quantitative measure of the extent to which an acid dissociates or ionizes in water. It represents the equilibrium constant for the dissociation reaction of an acid in water and provides information about the strength of the acid.
The acid dissociation constant, Ka, is calculated by dividing the concentrations of the products (hydrogen ions and the conjugate base) by the concentration of the undissociated acid. The equation for the dissociation of a generic acid, HA, can be represented as follows: HA ⇌ H+ + A-. The acid dissociation constant is given by Ka = [H+][A-] / [HA].
A strong acid completely dissociates in water, meaning that nearly all of its acid molecules dissociate into ions. Examples include hydrochloric acid (HCl) and sulfuric acid (H2SO4). In contrast, a weak acid only partially dissociates in water, with a small fraction of its acid molecules forming ions. Examples include acetic acid (CH3COOH) and carbonic acid (H2CO3).
The acid dissociation constant is influenced by factors such as temperature, concentration of the acid, and the nature of the acid and its conjugate base. Higher temperatures generally increase the acid dissociation. Additionally, higher acid concentrations tend to enhance dissociation, while the presence of a stronger conjugate base weakens the acid's dissociation.
The pH of a solution is related to the concentration of hydrogen ions (H+). A higher concentration of hydrogen ions leads to a lower pH, indicating a more acidic solution. The degree of acid dissociation affects the concentration of hydrogen ions, thus impacting the pH. Greater acid dissociation results in higher hydrogen ion concentrations and lower pH values.
Yes, the acid dissociation constant provides valuable information about the strength of an acid. A larger Ka value indicates a stronger acid that dissociates more readily, while a smaller Ka value corresponds to a weaker acid that dissociates to a lesser extent.
The acid dissociation constant can be measured experimentally using various methods, including pH measurements, conductivity measurements, and spectrophotometric techniques. These methods involve analyzing the changes in the concentration of ions or the properties of the solution as the acid dissociates.
Scroll to Top