Given below are two statements: One is labelled as Assertion and the other is labelled as Reason . Assertion : HTe is more acidic than HS. Reason : Bond dissociation enthalpy of HTe is lower than HS. In light of the above statements, choose the most appropriate from the options given below:
Both and are true but is NOT the correct explanation of .
Both and are true and is the correct explanation of .
is false but is true.
is true but is false.
Correct answer:D
Standard Method
Given: Assertion states that HTe is more acidic than HS. Reason states that the bond dissociation enthalpy of HTe is lower than that of HS.
Find: Whether Assertion and Reason are true, and whether correctly explains .
The solution explains that acidity of binary hydrides increases down the group because the H–X bond becomes weaker as atomic size increases. Thus HTe has weaker H–Te bonds than the H–S bonds in HS, so it ionizes more easily and behaves as a stronger acid.
It also states that the bond dissociation enthalpy of HTe is lower than that of HS, which supports easier bond cleavage and therefore greater acidity.
Therefore, both Assertion and Reason are true, and is the correct explanation of .
However, the provided the solution also displays "The Correct Option is D," which contradicts the written explanation and the provided correct answer. Based on the actual solution reasoning, the correct option should be B.
Detailed Explanation
Given: Compare the acidity of HTe and HS, and examine whether lower bond dissociation enthalpy explains the trend.
Find: The correct assertion–reason relationship.
The extracted explanation states the trend of Group hydrides as:
So HTe is more acidic than HS. This makes Assertion true.
The explanation further states that down the group, atomic size increases, orbital overlap with hydrogen becomes poorer, and the H–X bond becomes longer and weaker. Hence the bond dissociation enthalpy decreases.
So for these two hydrides:
Therefore, the bond dissociation enthalpy of HTe is lower than that of HS. This makes Reason true.
Since lower bond dissociation enthalpy means the proton can be released more easily, correctly explains why HTe is more acidic than HS.
Therefore, both and are true, and is the correct explanation of . The correct option is B.
Discrepancy note: the heading in the solution says option D, but the written chemical reasoning clearly supports option B.
Students often compare acidity only using electronegativity. That approach is incomplete here because for hydrides down a group, bond strength becomes the dominant factor. Use the decreasing H–X bond strength trend to judge acidity.
A common mistake is to think lower bond dissociation enthalpy means lower acidity. In fact, a weaker H–X bond makes proton release easier, so acidity increases. Relate bond breaking to ease of deprotonation.
Some students confuse the acidic trend across a period with the trend down a group. For Group hydrides, acidity increases down the group from HO to HTe. Use the correct periodic trend before evaluating the assertion.
Get unlimited AI-adaptive practice, mastery tracking, and an AI tutor that explains every step — free to start.
