Q.ii. “Acidic character of hydrides of group 16 elements increases from H$_2$O to H$_2$Te” Explain.

Step 1. All four hydrides (H2O, H2S, H2Se, H2Te) are weakly acidic in water, but not equally so.

Step 2. Table 7.5 shows the H-E bond dissociation enthalpy falls steadily down the group: 463 (H2O) -> 347 (H2S) -> 276 (H2Se) -> 238 kJ/mol (H2Te).

Step 3. A weaker H-E bond breaks (ionises) more readily when the hydride dissolves in water, releasing H+ more easily -- so the weakest-bonded hydride, H2Te, is the most acidic, and the strongest-bonded, H2O, is the least acidic among the four (though H2O's acidity here refers to its own very weak self-ionisation/Bronsted behaviour, not to be confused with water simply being 'neutral' in the everyday sense).

Step 4. This directly matches the falling pKa values in Table 7.5 (14.0, 7.0, 3.8, 2.6) -- a lower pKa meaning a stronger acid, confirming H2Te as the strongest and H2O as the weakest of the four.

Connecting the stated acidity trend to its underlying physical cause (falling H-E bond dissociation enthalpy down the group) and cross-checking against the pKa data in Table 7.5.

• Explaining the trend purely in terms of 'electronegativity decreasing down the group', which is directionally true but is not itself the acidity-determining factor here; the decisive quantity is bond strength (dissociation enthalpy), not electronegativity.