Problem 118 - A Complete Resource Book in Chemistry for JEE Main

Question

Decreasing order of electron affinity is (a) \(\mathrm{F}>\mathrm{Cl}>\mathrm{Br}>\mathrm{I}>\mathrm{S}>\mathrm{Si}\) (b) \(\mathrm{Si}>\mathrm{S}>\mathrm{F}>\mathrm{Cl}>\mathrm{Br}>\mathrm{I}\) (c) \(\mathrm{Cl}>\mathrm{F}>\mathrm{Br}>\mathrm{I}>\mathrm{S}>\mathrm{Si}\) (d) \(\mathrm{Cl}>\mathrm{F}>\mathrm{Br}>\mathrm{S}>\mathrm{Si}>\mathrm{I}\)

Step-by-Step Solution

Answer

The correct order is (c): \\ \( \mathrm{Cl} > \mathrm{F} > \mathrm{Br} > \mathrm{I} > \mathrm{S} > \mathrm{Si} \).

Step 1: Understanding Electron Affinity

Electron affinity is the energy released when an electron is added to a neutral atom. Generally, across a period, electron affinity increases from left to right due to an increasing effective nuclear charge. Down a group, electron affinity decreases because of increased atomic size, which means added electrons are farther from the nucleus.

Step 2: Comparative Electron Affinities of Halogens

Among the options given, we know that halogens generally have high electron affinities due to their high nuclear charge and small size. However, fluorine (F) has a slightly lower electron affinity than chlorine (Cl) due to electron-electron repulsions in its small 2p orbital.

Step 3: Position of Sulfur and Silicon

Sulfur (S) also has a relatively high electron affinity but less than that of halogens. Silicon (Si), being a metalloid and positioned further left on the periodic table, has a much lower electron affinity compared to the others listed.

Step 4: Determining the Order

Based on previous steps, the decreasing order of electron affinity is: Chlorine (Cl) > Fluorine (F) > Bromine (Br) > Iodine (I) > Sulfur (S) > Silicon (Si). This sequence represents option (c).

Key Concepts

Periodic Trends

Periodic trends in the context of electron affinity refer to the way certain properties of elements change as you move across periods or down groups in the periodic table. One common trend is that electron affinity increases as you move from left to right across a period. This is because the effective nuclear charge increases, pulling electrons in more strongly.

However, as you descend a group, electron affinity tends to decrease. The added electron sits in a higher orbital further from the nucleus, reducing the energy released. Understanding these trends can help predict the reactivity and characteristics of different elements.

Halogens

Halogens are part of Group 17 in the periodic table and are known for their high reactivity and high electron affinity. This includes the elements like fluorine, chlorine, bromine, and iodine. Halogens need one more electron to achieve a stable configuration of a noble gas, making them highly reactive.

Interestingly, although fluorine is more electronegative, chlorine has a higher electron affinity. This is due to fluorine's small size, which leads to increased electron-electron repulsion in its compact electron shell. Thus, while both elements have high electron affinities, chlorine tops the list in terms of electron affinity among halogens.

Effective Nuclear Charge

Effective nuclear charge (Ze) is the net positive charge experienced by an electron in an atom. This is calculated by considering the total positive charge of the nucleus minus the shielding or screening effect of inner electrons.

Across a period, the effective nuclear charge increases as protons are added to the nucleus without a corresponding increase in the core electrons, thus increasing the nuclear pull on the added electron and thus, electron affinity. This concept helps explain the periodic trend where electron affinity increases across periods in the periodic table.

Atomic Size

Atomic size or atomic radius refers to the distance from the center of the nucleus to the outermost electron shell. Changes in atomic size affect electron affinity.

As you move down a group in the periodic table, atomic size increases due to added electron shells. This increase decreases electron affinity since the outer electrons are further from the nucleus.

• Greater distance lessens the nuclear attraction experienced by added electrons.
• Increased distance results in lower energy release when an electron is added.

Consequently, elements higher in a group often exhibit higher electron affinities compared to those further down.