1. Concept used: Electron gain enthalpy

Electron gain enthalpy is the enthalpy change when an electron is added to an isolated gaseous atom:

$X(g) + e^- \to X^-(g)$

More negative electron gain enthalpy means greater tendency to accept an electron.

For Group 16 elements, the general trend is that electron gain enthalpy becomes less negative down the group, but there is an important exception involving oxygen.

2. Elements given

The elements are:

$O,\ S,\ Se,\ Te$

These all belong to Group 16.

3. Expected group trend

Normally, on moving down a group, atomic size increases, so the added electron enters a larger orbital and experiences less attraction from the nucleus. Thus, electron gain enthalpy generally becomes less negative down the group.

So among

$S,\ Se,\ Te$

we expect:

$S > Se > Te$

(where “greater” means more negative / more favorable electron gain enthalpy).

4. Why oxygen is anomalous

Oxygen is very small, so the incoming electron must enter a compact $2p$ orbital. This leads to strong electron-electron repulsion in the small valence shell.

Because of this, oxygen has less negative electron gain enthalpy than sulfur.

Hence:

$S > O$

Also, oxygen is less favorable than selenium as well, so the actual order among these is:

$S > Se > Te > O$

5. Check each option

• A: $Te > Se > S > O$
  Incorrect, because down the group electron gain enthalpy does not increase like this.

• B: $S > Se > Te > O$
  Correct.

• C: $S > O > Se > Te$
  Incorrect, because oxygen does not have more negative electron gain enthalpy than selenium.

• D: $O > S > Se > Te$
  Incorrect, because oxygen is anomalously less negative than sulfur.

6. Final answer

The correct order is:

$\boxed{S > Se > Te > O}$

So the correct option is B.