If the acid base reaction HA(aq)+B^-(aq)---> HB(aq)+A^-(aq)

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has a k = 10^-4,which of the following statements are true & why?
(a)HB is a stronger acid than HA
(b)HA is a stronger acid than HB
(c)HA & HB have same acidity
(d)B^- is a stronger base than A^-
(e)A^- is a stronger base than B^-
(f)B^-and HB are conjugate acid-base pair
(g)the acid & base strength can cannot be compared

To determine which of the statements are true, we need to analyze the given equilibrium reaction and the equilibrium constant (K). The equilibrium constant expression for this reaction is:

K = [HB][A^-] / [HA][B^-]

Given that the value of K is 10^-4, let's evaluate each statement:

(a) HB is a stronger acid than HA:
This statement is not necessarily true. The strength of an acid is determined by its ability to donate a proton (H+). The equilibrium constant does not directly indicate the strength of the acid. Therefore, we cannot conclude which acid (HA or HB) is stronger based solely on the given information.

(b) HA is a stronger acid than HB:
Similar to statement (a), we cannot conclude which acid is stronger based on the equilibrium constant. Thus, this statement is not necessarily true.

(c) HA & HB have the same acidity:
Without additional information about the reactants, we cannot determine whether the acidity of HA and HB is the same or not. Therefore, we cannot conclude that this statement is true.

(d) B^- is a stronger base than A^-:
Since the equilibrium constant expression only involves the concentration of the species in the reaction, it does not provide any information about the relative strengths of the bases (B^- and A^-). Comparing their strengths is beyond the scope of what can be deduced from the given information. Thus, this statement is not necessarily true.

(e) A^- is a stronger base than B^-:
Similar to statement (d), we cannot determine the relative strengths of the bases (A^- and B^-) solely based on the equilibrium constant. Therefore, this statement is not necessarily true.

(f) B^- and HB are conjugate acid-base pairs:
In this reaction, HB is formed when HA donates a proton (H+). Likewise, A^- is formed when B^- accepts a proton. Therefore, HB and B^- can be considered a conjugate acid-base pair. This statement is true.

(g) The acid and base strength cannot be compared:
Based on the given information and the equilibrium constant, we cannot directly compare the strengths of the acids and bases in this reaction. Hence, this statement is true.

In summary, the true statements are:
- (f) B^- and HB are conjugate acid-base pair
- (g) The acid and base strength cannot be compared.

To determine which of the statements are true, let's consider the given reaction and the value of K, the equilibrium constant.

The equation for the reaction is:
HA(aq) + B^-(aq) ↔ HB(aq) + A^-(aq)

The equilibrium constant, K, is equal to the ratio of the products over reactants, with each concentration raised to the power of its coefficient.

Now, let's analyze each statement one by one:

(a) HB is a stronger acid than HA:
False. In this reaction, HA and HB are not directly compared. The equilibrium constant (K) only provides information about the extent to which the reaction occurs, not the relative strength of the individual acids.

(b) HA is a stronger acid than HB:
False. Similar to statement (a), the equilibrium constant does not indicate the relative strength of the individual acids.

(c) HA and HB have the same acidity:
False. The equilibrium constant does not provide information about the relative acidity of HA and HB.

(d) B^- is a stronger base than A^-:
False. The equilibrium constant does not provide information about the relative strength of the bases.

(e) A^- is a stronger base than B^-:
False. Similar to statement (d), the equilibrium constant does not indicate the relative strength of the bases.

(f) B^- and HB are a conjugate acid-base pair:
True. In the reaction, B^- acts as a base by accepting a proton (H+) from HA and forms HB. Thus, B^- is the conjugate base of HB.

(g) The acid and base strength cannot be compared:
False. While the equilibrium constant itself does not directly compare the strength of the acids or bases, there are other methods (such as pKa or pKb values) that can be used to compare the relative strengths of acids and bases.

In conclusion, (f) is the only true statement. The others are false because the equilibrium constant does not directly provide information about the relative acid or base strengths.

In a Bronsted-Lowry acid-base reaction with a high Keq value,

Acid + Base = Weaker acid + Weaker Base
Acid1 + Base1 = Base2 + Acisd2
Acid1 is stronger than Acid2, and,
Base1 is stronger than Base2
In your question, the equilibrium constant is small. The forward (left to right) reaction does not go far. We can assume the the acid and the base on the right side are stronger than the acid and the base on the left.
Based on the above discussion, (a) is a correct statement. I will let you decide on the others.