E13 Acid-base equilibria in water : strong and weak acids

Aim: To compare the dissociation percentage of aqueous solutions of 1 mol/l in strong, medium strong and weak acids and to examine the relationship between the Ka of an acid, its concentration and its degree of dissociation.

The homogeneous equilibrium reactions thus far considered have been mainly reactions in the gas phase. In reality most reactions (equilibrium reactions and nonequilibriumn reactions) take place in water. If all the components present are completely dissolved (i.e. homogeneously distributed over the whole liquid phase), then such reactions are homogeneous reactions in a closed system.

In most cases the test tubes, beakers and reaction vessels remain open to the air, but the evaporation of water is fairly slow. However, “open” systems in which very volatile components such as gases are present, particularly if they are poorly soluble in water, can not be regarded as closed systems.

The illustrations H1 to H7 in DIDAC-1 form an introduction to this topic, covering the role of water in different processes such as dissolution, protonation and deprotonation of a molecule or ion. This species appears as a base or an acid.

An extremely important aspect of such acid-base behaviour is that H3O+ or OH- ions are to a lesser or greater extent to be found in the aqueous solution. Very many chemical properties depend upon the presence of these ions. The mostly small concentrations of these ions are usually expressed in terms of pH and pOH, where pH = -log [H3O+]e and pOH = -log [OH-]e.
This avoids the necessity of working with exponents.

In pure water the concentrations of H3O+ and OH- ions are very small: i.e. 10-7 mol/L. The very low concentrations of these ions are due to the dissociation of water.

For the equilibrium reaction :

For the equilibrium constants for the equilibria of acids, the symbol Ka(cid) is used rather than Kc. The concentration of water [H2O]e is, by international convention, not included in the expression for Ka. Generally, Ka values refer to a temperature of 25 °C.

In the following examples the dissociation of water will not be considered. It is negligible compared with the dissociation of acids or bases, which also provide H3O+ or OH- ions.

Virtually all acids are very soluble in water. They form homogeneous solutions. The gas hydrogen chloride is extremely soluble in water (more than 12 mol/L), its aqueous solution being known as hydrochloric acid. All HCl-molecules which dissolve in water dissociate forming hydrated protons H3O+ and chloride anions. This is completely analogous to the process occurring when a liquid such as HNO3 is added to water. This process is depicted in illustration H2 of DIDAC-1. The acids HCl, HBr, HI, HClO4, HNO3 and H2SO4 (at least the first deprotonation) exhibit this complete deprotonation reaction or dissociation under almost all conditions, provided that the dissociation takes place in water. Such reactions are not regarded as equilibrium reactions. Most other stable acids must be classified under the weak acids.


  An example of a weak acid found in a school or at home, is ethanoic acid (acetic acid). 8 % Household vinegar contains 1.3 mol/L of CH3COOH.
This acid also dissolves easily in water, but its dissociation reaction (protonation of water) is very limited. There is no connection between solubility and dissociation tendency.
Many acids dissolve well, while exhibiting a low dissociation tendency. The polarity of an acid plays an important rôle in the solubility of an acid. Such incomplete protonation, dissociation or acid-base reactions in water are dealt with in illustration H1 of DIDAC-1.

Illustration E13 shows the substantial differences in dissociation percentage between a strong acid (HCl) and two weak acids ( and CH3COOH). The solutions all have the same concentration : 1 mol/L.
is an anion with acidic properties. It is prepared by dissolving the white powder NaHSO4 in water. The Ka-value of these acids are 103 for HCl, 1.2 x 10-2 for and 1.8 x 10-5 for .

The pie-charts show the equilibrium percentage of the molecules dissociated in water:

=> light green areas represent non-dissociated acid molecules
=> red areas represent the fraction of H3O+-ions formed
=> purple areas represent the acid anions.
The Na+-ions from the NaHSO4 are represented colourless in this piechart, as they do not affect the acid-base properties. Water is also not specifically mentioned. When an acid fully dissociates, 100 parts of acid produce 100 parts of the anion corresponding to the acid and 100 parts of hydrogen-ions (H3O+).

The appropriate equilibrium reaction and equations from which the equilibrium concentrations can be calculated are given under each pie-chart.

By way of illustration, this is demonstrated below for the acid dissociation of the hydrogen sulphate ion .
The appropriate equilibrium reaction is :

This can be used to obtain the value of x.

From the illustration, it can be seen that for a weak acid such as ethanoic acid, the degree of dissociation is very low. In a 1M solution, the degree of dissociation is 0.0042 or 0.42 %. The ethanoic acid concentration in household vinegar is 1.3mol/L, giving a degree of dissociation of 0.0037 and an equilibrium concentration of [H3O+]e of 0.0048 mol/L.

For a solution with a concentration of 1 mol/L the dissociation percentage is 10.4 %. Weak acids with such Ka-values (ca. 10-2) are not common.
Most acids are in fact weak acids with similar properties to ethanoic acid. Many of such acids are used in food and in our metabolism. Even with very weak acids, a degree of dissociation of 1 can be attained upon extreme dilution in water.