The correct order of electron gain enthalpy is (a) S Se Te O (b) Te Se S O (c) O S S

Q: Why does electron gain enthalpy generally decrease down the group?

As atomic size increases down the group, the distance between the nucleus and the incoming electron increases. This reduces the electrostatic attraction, leading to less energy being released.

Q: Why is Oxygen an exception to this trend?

Oxygen has an exceptionally small atomic size. The high electron density in its compact $2p$ subshell causes strong inter-electronic repulsion for any incoming electron, reducing the net energy released.

Q: Does this anomaly happen in other groups as well?

Yes! This is a characteristic of second-period elements. For example, in Group 17, Fluorine has a lower electron gain enthalpy than Chlorine for the exact same reason.

Q: Is electron gain enthalpy always negative?

Not always. While adding an electron to most non-metals releases energy (negative enthalpy), adding an electron to noble gases or an already negative ion requires energy input (positive enthalpy) due to strong repulsions.

Q: What is the difference between electron affinity and electron gain enthalpy?

They represent the same physical process but with opposite sign conventions. Electron affinity is the energy released (positive value), while electron gain enthalpy is the change in enthalpy of the system (negative value when energy is released).

Question

The correct order of electron gain enthalpy is

Options:

(A) S > Se > Te > O

(B) Te > Se > S > O

(C) O > S > Se > Te

(D) S > O > Se > Te

Sigma Percentile · Accepted Answer

### The Expected Trend

When we talk about electron gain enthalpy, we are looking at the energy released when a neutral gaseous atom accepts an extra electron.

Mathematically, this process is represented as:
$$X(g) + e^- ightarrow X^-(g) + 	ext{Energy}$$

Generally, as we move down a group in the periodic table, the atomic size increases.

Because the added electron is further from the nucleus, the effective nuclear attraction decreases.

Therefore, the energy released (the magnitude of electron gain enthalpy) should also decrease down the group.

Following this logic for Group 16, we would expect Oxygen to release the most energy, giving the order: **$O > S > Se > Te$**.

### The "Small Room" Anomaly

However, Chemistry loves its exceptions, and Oxygen is a classic one!

Oxygen is a second-period element, which means its atomic size is exceptionally small.

Imagine a tiny room that is already crowded with people.

When a new electron tries to enter Oxygen's compact $2p$ orbital, it faces intense **inter-electronic repulsion** from the already present electrons.

### The Final Verdict

Because of this strong repulsion, the incoming electron isn't welcomed as easily.

The atom actually has to spend some energy to overcome this repulsive force.

As a result, the net energy released by Oxygen is significantly lower than expected.

In fact, it drops below that of Sulphur!

For the rest of the elements (Sulphur, Selenium, Tellurium), the size is large enough that this repulsion is negligible, so they follow the normal trend.

This gives us our final, correct order for the magnitude of electron gain enthalpy:
**$S > Se > Te > O$**

This is a high-yield concept for JEE, and you will see the exact same anomaly in Group 17, where Chlorine has a higher electron gain enthalpy than Fluorine!

Animated Solution for Chemistry - s and p-Block Elements: The correct order of electron gain enthalpy is

Select Answer:

Visualized Solution

Group 16 Elements

Electron Gain Enthalpy $Δ_{e g} H$

The Anomaly of Oxygen

Inter-electronic Repulsion

Oxygen vs Sulphur

Final Order

Key Takeaway

The Sigma Insight: Group 16 Elements

The Expected Trend

The "Small Room" Anomaly

The Final Verdict

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