Question

2A. Complete the following table.

Element	Zeff	Atomic orbital designation for highest energy valence electron (i.e., 2s)
Se
Kr

2B. Which atom (Se or Kr) has the smaller first ionization energy? Enter the chemical symbol in the space provided.

2C. Which one of the following statements best explains the trend observed in part 2?

Ionization energy decreases down a group. Elements lower in a group have larger atomic radii. The valence electrons are further from the nucleus so that the attraction between the valence electrons and nucleus is decreased and the energy required to remove the valence electrons is less.
Ionization energy decreases to the right across a period. Within a period the increasing number of valence electrons result in stronger electron-electron repulsion thereby decreasing the amount of energy required to remove the valence electrons.
Ionization energy increases down a group. The valence electrons occupy a higher energy atomic orbital, thereby requiring more energy to remove the valence electron.
Ionization energy increases to the right across a period. Within a period valence electrons of elements to the right have a larger Z_eff and therefore feel a stronger attraction to the nucleus, requiring more energy to remove the valence electrons.

Answer 1

first write electronic configuration and rearrange in increasing order of energy. Now calculate shielding constant and then effective nuclear charge.

(2B) More the value of effective nuclear charge more will be attractions between nucleus and outermost electron, so ionisation energy of 'Se' is lower.

(2C) correct explanation is fourth statement because from left to right in a period effective nuclear charge increases and also Nobel gases are completely filled. So they have very less tendency to eject electro hence their ionisation energy is larger.