Which of the following is an example of a conjugate acid-base pair?
A. $H _2 S$ and $HS ^{-}$
B. $NO _2$ and $NO _2^{-}$
C. HCl and HBr
D. HF and $H ^{+}$
Answer (2)
The correct answer is Option A: H₂S and HS⁻, as they form a conjugate acid-base pair by differing only by one proton (H+). The hydrosulfide ion (HS⁻) is formed when hydrogen sulfide (H₂S) donates a proton. The other options do not meet the criteria for conjugate acid-base pairs.
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A conjugate acid-base pair differs by only a proton ( H + ).
Examine each option to see if the two species differ by only a proton.
H 2 S can lose a proton to become H S − , fitting the definition.
Therefore, the conjugate acid-base pair is H 2 S and H S − .
Explanation
Understanding Conjugate Acid-Base Pairs A conjugate acid-base pair consists of two species that differ by only a proton ( H + ). We need to identify which of the given options fits this definition.
Analyzing Each Option Let's examine each option:
Option A: H 2 S and H S − . H 2 S can lose a proton ( H + ) to become H S − . This fits the definition of a conjugate acid-base pair.
Option B: N O 2 and N O 2 − . N O 2 would need to gain a proton to become H N O 2 to be related to N O 2 − . These are not a conjugate acid-base pair.
Option C: HCl and HBr . These are both strong acids, but they don't form a conjugate acid-base pair with each other. They are different acids altogether.
Option D: HF and H + . HF would need to lose F − to become H + . These are not a conjugate acid-base pair.
Identifying the Correct Pair Based on the analysis, the only option where the two species differ by only a proton is option A. Therefore, H 2 S and H S − form a conjugate acid-base pair.
Examples
Conjugate acid-base pairs are essential in understanding buffer solutions, which are used to maintain a stable pH in various chemical and biological systems. For example, in our blood, the carbonic acid ( H 2 C O 3 ) and bicarbonate ( H C O 3 − ) system acts as a buffer to keep the blood's pH within a narrow range, crucial for the proper functioning of enzymes and other biological processes. The equilibrium between H 2 C O 3 and H C O 3 − helps neutralize excess acids or bases, preventing drastic changes in pH that could be harmful.
