Concept explainers
Interpretation:
Why the setup for
Concept Introduction:
There are associated scales within the chemistry utilized to find out how basic and acidic a solution is, as well as the strength of bases and acids. Though most familiar is pH scale, Ka, pKa, Kb and pKb are the common calculation which offer understanding into acid-base reactions.

Answer to Problem 81P
Scale of Kb values for bases is unnecessary because the strength of two bases can be compared from the strength of their conjugate acids.
Explanation of Solution
The acid ionization constant, Ka, is the equilibrium constant for acid’s ionization. For the ionization of an acid, say HA, in an aqueous solution we get.
The square brackets denote the concentration. The concentration of water is not considered as water is exist in excess.
The Ka equation displays that the greater the dissociation of acid, the greater the value of Ka.
If there is a stringer acid, it will dissociate to a greater extent. Therefore, the numerator within the equation of Ka is larger, and thus, Ka will have the greater value.
Therefore, the weak acid’s strength is expressed in terms of Ka.
As per the Bronsted-Lowry concept of bases and acids, an acid is known as the proton donor, and a base is the proton acceptor.
When an acid donates a proton to another ion or molecule, it is converted to the conjugate base. For illustration,
When a base accepts a proton, it is converted to the conjugate acid. For illustration.
As per the Bronsted-Lowry definitions any molecules or ions pairs which might be interconverted by the transfer of proton is known as conjugate acid-base pair. Therefore, CH3 COOH and CH3 COO- is a conjugate acid-base pair and NH3 andNH4 + is a conjugate acid-base pair.
There is an inverse relationship amid the strength of acid and its conjugate base. The stronger the acid, the weaker is the conjugate base. The acid ionization constant, Ka, expresses the weak acid’s strength.
The below table shows values of ka for weak acids:
Ka values of weak acids.
Formula | Name | Ka | |
H3 PO4 | Phosphoric acid | ||
HCOOH | Formic acid | ||
CH3 CH(OH)COOH | Lactic acid | ||
CH3 COOH | Acetic acid | ||
H2 CO3 | Carbonic acid | ||
H2 PO4 - | Dihydrogen phosphate ion | ||
H3 BO3 | Boric acid | ||
NH4 + | Ammonium ion | ||
HCN | Hydrocyanic acid | ||
C6 H5 OH | Phenol | ||
HCO3 - | Bicarbonate ion | ||
HPO42 - | Hydrogen phosphate ion |
If we want to compare the strength of two bases, we can compare the strength of their conjugate acids.
For illustration, let us compare the strength of two bases, phenoxide ion C6 H5 O- and bicarbonate ion HCO3-.
The conjugate acid of phenoxide ion i.e. C6 H5 O- is phenol C6 H5 OH.
Ka of phenol =
The conjugate acid of bicarbonate ion HCO3 - is carbonic acid H2 CO3.
Ka of carbonic acid =
Therefore, Ka of the carbonic acid is greater than Ka of phenol. Hence, the carbonic acid is stronger than phenol. The stronger the acid, the weaker is the conjugate base. Therefore, a phenoxide ion is stronger base than a bicarbonate ion.
Therefore, we could compare the strength of weak bases from the Ka scale of acids. Hence, a scale of Kb values for bases is unncessary.
Since, the strength of two bases can be compared from the strength of their conjugate acids thus, scale of Kb values for bases is unnecessary.
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