Question

Give reasons for the following:
(A) Oxygen has lower ionisation energy than that of nitrogen.
(B) Electron gain enthalpy of chlorine is more negative than that of fluorine.
(C) Arrange \[{O^{2 - }},A{l^{3 + }},N{a^ + },{F^ - }and{\text{ }}M{g^{2 + }}\] in the increasing order of their sizes.

Answer 1

Hint: Ionisation energy is maximum for half-filled and completely filled electronic configurations. Electron gain enthalpy becomes more negative as the size increases. Size of an element increases in increasing the number of electrons in its valence shell.

Complete step-by-step solution:
(A) Ionisation energy is the energy required to remove an electron from the valence shell of an isolated gaseous atom. It increases with greater interaction between nucleus and electrons.
In case of nitrogen, the electronic configuration is half-filled providing extra stability to the atom. This tells us that there is a greater attraction between nucleus and valence electrons. Therefore, the tendency to lose an electron from the valence shell becomes difficult and requires a lot of energy.
While in an oxygen atom, there are four electrons present in the valence shell and the atoms tend to gain extra stability by losing one electron from the shell and attain half-filled electronic configuration. This explains why oxygen has lower ionisation energy than that of nitrogen.
EC of nitrogen is: \[1{s^2}2{s^2}2{p^3}\]
EC of oxygen is: \[1{s^2}2{s^2}2{p^4}\]

(B) Electron gain enthalpy is the change in energy when an electron is added in the valence shell of an isolated gaseous atom. When an electron is added to the system, there occurs an extra repulsion among the electrons. It becomes more negative for those elements which have a high tendency to accept an electron.
For example, halogens have very highly negative enthalpy because they can attain noble gas configuration on accepting a single electron.
In case of chlorine and fluorine, when an electron is added to the 2p orbital of fluorine, the electron experiences too much repulsion from the existing electrons. Adding an electron to chlorine is easy because it goes to a 3p orbital which is far from the nucleus and has large space to minimize electron-electron repulsion. Therefore, Electron gain enthalpy of chlorine is more negative than that of fluorine.
EC of chlorine is: \[1{s^2}2{s^2}2{p^6}3{s^2}3{p^5}\]
EC of fluorine is: \[1{s^2}2{s^2}2{p^5}\]

(C) Atomic size is the mean distance from the centre of the nucleus to the outermost shells of electrons. Atomic size of a neutral atom increases with increase in number of electrons as in anions and decreases with loss of electrons as in cations. Anions are generally bigger in size as compared to cations. These are called ionic sizes.
The increasing order of their sizes be \[A{l^{3 + }} < M{g^{2 + }} < N{a^ + } < {F^ - } < {O^{2 - }}\].
The trivalent aluminium cation has the smallest size as it has lost three electrons from its outermost shell. Due to which nuclear charge has increased and thereby size is reduced.
And in divalent oxygen anion, ionic size increased to this extent due to gain of two electrons and decrease in the nuclear charge.

Note: As per the trends followed in periodic table, ionisation energy increases along the period and electron gain enthalpy becomes more negative as we go down the group. The exception to the first one is discussed in (a).