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.
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.
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
=> 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+).
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.
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.