Question

iodine has an electron configuration of $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{10}4p^{6}5s^{2}4d^{10}5p^{5}$. how would iodine and other elements ending in $p^{5}$ differ? their valence electrons are in different subshells they have different numbers of valence electrons their valence electrons are in different shells they would be in a different group

Explanation:

• Valence electrons: For elements ending in \(p^{5}\), the number of valence electrons is \(7\) (e.g., in iodine \(5s^{2}5p^{5}\), valence electrons \(= 2 + 5=7\)). So the option "They have different numbers of valence electrons" is wrong.
• Subshells: All these elements have valence electrons in \(p\) sub - shell (as the last filling is \(p^{5}\)), so "Their valence electrons are in different subshells" is wrong.
• Group: Elements with \(ns^{2}np^{5}\) configuration (where \(n\) is the principal quantum number of the valence shell) are in group \(17\) (halogens). So "They would be in a different group" is wrong.
• Shells: The principal quantum number (\(n\)) of the valence shell is different. For example, \(F:2s^{2}2p^{5}\), \(Cl:3s^{2}3p^{5}\), \(Br:4s^{2}4p^{5}\), \(I:5s^{2}5p^{5}\). The valence electrons of these elements are in different shells (\(n = 2,3,4,5\) respectively for \(F,Cl,Br,I\)).

Answer:

Their valence electrons are in different shells