The order of first four elements of group-17 are as follows.
F<Cl<Br<I (Covalent radius)
Cl>F>Br>I (Electron affinity)
F<Cl<Br<I (Ionic radius)
F>Cl>Br>I (First ionization energy)
Electron affinity order is irregular.
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Approach Solution -2
This question asks to identify which of the listed properties show an irregular trend for the first four halogens (Fluorine, Chlorine, Bromine, and Iodine).
Concept Used:
We need to analyze the periodic trends for the following properties down Group 17:
Covalent Radius and Ionic Radius: The size of atoms and their corresponding ions. The general trend is that atomic and ionic radii increase down a group due to the addition of a new electron shell in each successive period.
First Ionization Energy: The energy required to remove the most loosely bound electron from a neutral gaseous atom. The general trend is that ionization energy decreases down a group because the valence electrons are farther from the nucleus and experience greater shielding from inner electrons, making them easier to remove.
Electron Affinity: The energy change that occurs when an electron is added to a neutral gaseous atom. A more negative value indicates a greater affinity for an electron. Generally, electron affinity becomes less negative (decreases in magnitude) down a group due to increasing atomic size. However, exceptions exist, particularly in the 2nd period.
Step-by-Step Solution:
Step 1: Analyze Covalent Radius and Ionic Radius.
For the first four halogens (F, Cl, Br, I), a new principal energy level is added for each element down the group. This leads to a steady and predictable increase in both the covalent and ionic radii.
As we move down Group 17, the atomic size increases, and the shielding effect of the inner electrons becomes more significant. Consequently, the outermost electron is held less tightly by the nucleus, and the energy required to remove it decreases.
\[ \text{Order of First Ionization Energy: } \text{F} > \text{Cl} > \text{Br} > \text{I} \]
This trend is also regular.
Step 3: Analyze Electron Affinity.
Based on the general trend, one would expect electron affinity to decrease down the group (F > Cl > Br > I). However, an important exception occurs between Fluorine and Chlorine.
Fluorine, being a very small atom (in the 2nd period), has a very high electron density in its compact 2p subshell. When an external electron is added, it experiences significant repulsion from the electrons already present. This interelectronic repulsion reduces the net energy released.
Chlorine (in the 3rd period) is larger and has a more diffuse 3p subshell. The incoming electron can be accommodated with less interelectronic repulsion. As a result, Chlorine has a more exothermic (more negative) electron affinity than Fluorine.
The actual experimental order for the magnitude of electron affinity is:
\[ \text{Cl} > \text{F} > \text{Br} > \text{I} \]
This trend is irregular because the first element (F) has a lower value than the second element (Cl), breaking the expected monotonic decrease.
Step 4: Conclusion.
Among the given options, only electron affinity shows a significant irregularity for the first four elements of Group 17. The other properties (covalent radius, ionic radius, and first ionization energy) follow a regular, predictable trend.
Therefore, the correct property is Electron affinity.
