Problem 61 - A Complete Resource Book in Chemistry for JEE Main

Question

The electron affinities of \(\mathrm{N}, \mathrm{O}, \mathrm{S}\) and \(\mathrm{Cl}\) are (a) \(\mathrm{O} \approx \mathrm{Cl}<\mathrm{N} \approx \mathrm{S}\) (b) \(\mathrm{O}<\mathrm{S}<\mathrm{Cl}<\mathrm{N}\) (c) \(\mathrm{N}<\mathrm{O}<\mathrm{S}<\mathrm{Cl}\) (d) \(\mathrm{O}<\mathrm{N}<\mathrm{Cl}<\mathrm{S}\)

Step-by-Step Solution

Answer

The correct order is (c) \( N < O < S < Cl \).

Step 1: Understanding Electron Affinity

Electron affinity is the amount of energy released when an electron is added to a neutral atom in the gaseous state. Higher electron affinity indicates a greater tendency of the atom to gain an electron.

Step 2: Analyzing Trends Across Periods

In the periodic table, moving from left to right across a period, electron affinity generally increases except for some anomalies.

Step 3: Analyzing Trends Down Groups

Moving down a group, the electron affinity tends to decrease because additional electron shells reduce the effective nuclear charge experienced by the valence electrons.

Step 4: Consider Specific Elements

For the given elements: - Chlorine (431Cl432) has the highest electron affinity due to its position in Group 17 (halogens), which is very electronegative. - Oxygen (431O432) has lower electron affinity than Sulfur (431S432) despite being above Sulfur in the group, due to its small size. - Nitrogen (431N432) generally has a low electron affinity due to a half-filled 2p orbital that provides extra stability.

Step 5: Ordering Electron Affinities

Using the trends and specific element considerations, the order of electron affinities should be: \[ N < O < S < Cl \]

Key Concepts

Periodic Table Trends

When we analyze electron affinity within the context of the periodic table, a pattern emerges. Generally, as you move from left to right across a period, electron affinity increases. This trend occurs because atoms become more effective at attracting electrons to fill their outermost shell. More protons in the nucleus enhance the attraction of electrons. However, there are exceptions to this trend.

Notably, elements like Nitrogen show a deviation. The reason for this lies in the stability provided by its half-filled orbitals, which we'll discuss later. While the left-to-right increase is prominent, there are blips due to such stability considerations and electron-electron repulsions.

Halogen Electron Affinity

Halogens, like Chlorine, are well-known for their high electron affinity. These elements are in Group 17 and need just one additional electron to achieve a noble gas configuration. This motivation is why halogens have such a strong desire to gain an electron, releasing more energy in the process.

Chlorine, in particular, has one of the highest electron affinities of all elements due to its electronegativity. This property makes it very reactive, often found in compounds rather than in its elemental form. The strong pull Chlorine exerts on electrons is a reason for its position as the top of electron affinity lists.

Stability of Half-Filled Orbitals

Nitrogen provides us with a good example of how stability from half-filled orbitals can affect electron affinity. Nitrogen's electronic configuration ends with a half-filled 2p orbital. This half-filled state is relatively stable due to symmetrical distribution of electrons, which reduces repulsion between electrons.

Adding an electron to this arrangement increases electron-electron repulsion, making it less favorable. Therefore, nitrogen displays a notable deviation in electron affinity compared to other elements in its period. Such stability not only explains nitrogen's lower electron affinity but also offers insights into other anomalies on the periodic table.

Group Trends in Electron Affinity

As you move down a group in the periodic table, electron affinity generally decreases. This decline happens because each subsequent element adds a new electron shell. This increase in electron shells causes the valence electrons to be further from the nucleus.

The increased distance reduces the effective nuclear charge felt by these outer electrons, meaning the electron's attraction to the nucleus diminishes. This concept explains why oxygen has a lower electron affinity than sulfur, even though it is higher up in Group 16. Each added shell contributes to this shielding effect, which is a fundamental reason for decreasing electron affinities down a group.