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The ionization equation of an acid HA
is characterized by a constant Ka, called acidity constant (or dissociation constant, or ionization
constant) :
The more the equilibrium is driven to the right, the stronger the acid is. Thus, Ka (or pKa) is a
direct measure of the acid strength.
Note:
To avoid the use of acidity constants Ka expressed in terms of powers of 10, an acid-base couple
HA/A- is characterized by its pKa.
The stronger the acid is, the larger the Ka and the smaller the pKa (which can even be negative
for very strong acids) !
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Strong and weak acids and bases
An acid stronger than H3O+ is called strong acid; it is totally dissociated in water and its pKa is
negative (pKa 0).
Examples:
Hydrochloric acid HCl
Nitric acid HNO3
Sulfuric acid H2SO4 (as far as the first ionization is concerned)
Hydrobromic acid (HBr)
Solvent leveling effect and pKa scale
The strongest acid which can exist in water is H3O+. All the stronger acids react with H2O to
yield H3O+. See for instance the hydrochloric acid : HCl dissociates completely in water to give
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the H3O+ and the very weak conjugate base Cl-. This can be interpreted as a leveling of HCl
acidity by the solvent.
Similarly, the strongest base that can exist in water is OH-. Any stronger base reacts with H2O to
form OH-. For instance, a strong base such as NH2- is completely replaced by OH- in aqueous
medium. That is the solvent can also level off the alkalinity of the solution.
Note: This leveling effect can be understood by the simple statement:
In solution, the strongest acid and the strongest base will always react to form a weaker
conjugate base and a weaker conjugate acid.
Calculation of pH and pOH
1. Strong Acid
Reminder: a strong acid HA is completely dissociated (pKa
0, ∆G°diss << 0)
Note: the introduction of a strong acid HA into water drives equilibrium 2 towards the left!
If [HA]0 is the total concentration of acid initially introduced into the aqueous solution.
Case (1): [HA]0 > 10-7 M : the H3O+ ions present in the solution come primarily from the
dissociation of HA and :
Case (2) : [HA]0< 10-7 M : one cannot neglect the autoionization of water :
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2. Strong Base
Reminder: a strong base B is completely protonated (pKb
0, ∆G°ass << 0)
The following two equations have to be taken into account:
Note: the introduction of the strong base B into water drives equilibrium 2 towards the left
If [B]0 is the total concentration of base initially introduced into the aqueous solution.
Case (1): [B]0 > 10-7 M: one can assume that the OH- come exclusively from the dissociation of
H2O protons under the action of B.
Case (2): [B]0< 10-7 M : the autoionization of water can no more be neglected.
3. Weak Acid
Reminder: a weak acid HA has a pKa in the range 0 - 14 (∆G°
(∆ ass > 0).
A weak acid is partially dissociated in water. At equilibrium, the species present in solution are
HA, A-, H3O+ and HO- (arising from water autoionization). The following equilibria must be
taken into consideration :
Calculation of the H3O+ concentration is quite complex, but approximations are often
used, leading to welcome simplifications:
(a) If the acid is concentrated enough (> 10-6 M), water autoionization can be
neglected
(b If the acid is sufficiently weak, it is little dissociated and can be considered as undissociated
4. Weak Base
Reminder: the pKb of a weak base is in the range 0 to 14 (∆
∆G°diss > 0).
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A weak base is partially protoned in water. At equilibrium, the species present in solution are B,
BH+ as well as H3O+ and HO- arising from the autoionization of water. The following equilibria
must be taken into consideration :
Calculation of [OH-] is complex, but approximations are often used to simplify it :
(a) If the base is concentrated enough (> 10 -6) the autoionization of water can be
neglected
(b) If the base is sufficiently weak (pKb
considered as being almost un-protonated.
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2), it is not much protonated and can be