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Chemistry Question 57 – JEE-MAIN 2025

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Given below are two statements: Statement I : H2Se is more acidic than H2Te Statement II : H2Se has higher bond enthalpy for dissociation than H2Te In the light of the above statements, choose the correct answer from the options given below

Consider how atomic size changes down a group and its effect on bond length and strength.

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Ninja Strategy•Periodic Trends Application

Recall the periodic trends for acidity and bond enthalpy down a group to quickly evaluate each statement.

Step 1: Analyze Acidity Trend (Statement I)✦ Active

Acidity of hydrides of Group 16 elements (H2O,H2S,H2Se,H2Te) increases down the group. This is because as the atomic size of the central atom increases, the E-H bond length increases, leading to a weaker bond. A weaker bond is easier to break, facilitating the release of H+ ions, thus increasing acidity. Therefore, H2Te is more acidic than H2Se. Statement I, which claims H2Se is more acidic than H2Te, is false.

Step 2: Analyze Bond Enthalpy Trend (Statement II)○ Expand

Bond enthalpy (or bond dissociation energy) generally decreases as bond length increases. Down Group 16, the atomic size of the central atom increases, leading to longer E-H bonds. Longer bonds are weaker and require less energy to break. Therefore, the H-Se bond is stronger and has a higher bond enthalpy than the H-Te bond. Statement II, which claims H2Se has higher bond enthalpy for dissociation than H2Te, is true.

Step 3: Conclusion○ Expand

Based on the analysis, Statement I is false and Statement II is true. This corresponds to option 4.

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