The first ionization enthalpy values (in kJ mol^–1) of group 13 elements are :

The explanation for the deviation from the general trend can be -

1. Ga has lower ionization enthalpy than Al.

2. Ga has higher ionization enthalpy than Al.

3. Al has higher ionization enthalpy than Ga. 

4. Ga has a lesser valence electron than Al.

Boron has lesser ionization enthalpy than Beryllium, because:

The graph between ionization energy and atomic number for
the first group elements is shown below:


    

The element represented by Y in the graph above is -

1. Cs 

2. Rb

3. Ca

4.  K 

A configuration with the lowest ionization enthalpy among the following is:
1. \(1 s^2 2 s^2 2 p^5\)$\mathrm{}$
2. \(1 s^2 2 s^2 2 p^3\)
3. \(1 s^2 2 s^2 2 p^6 3 s^1\)
4. \(1 s^2 2 s^2 2 p^6\)

Abnormally high ionization enthalpy of B as depicted in the graph below can be due to : 

1. Completely filled 2p subshell.

2. Completely filled 2s subshell.

3. Completely filled 3d subshell.

4. Completely filled 4f subshell.

${\mathrm{IP}}_1$ and ${\mathrm{IP}}_2$ for Mg are 178 Kcal mol^–1 and 348 Kcal ${\mathrm{mol}}^{-1}$, respectively. The energy required for the reaction, $\mathrm{Mg}\to {\mathrm{Mg}}^{2+}$ $+$ $2{\mathrm{e}}^{-}$ will be:

1. +170 Kcal ${\mathrm{mol}}^{-1}$

2. +526 Kcal ${\mathrm{mol}}^{-1}$

3. –170 Kcal ${\mathrm{mol}}^{-1}$

4. –526 Kcal ${\mathrm{mol}}^{-1}$

The incorrect statement about ionization enthalpy is:

For the second-period elements, the correct increasing order of first ionisation enthalpy is:

Amongst the following electronic configurations, the highest ionization energy
is represented by:

1. [Ne]3s²3p³

2. [Ne]3s²3p²

3. [Ar]3d¹⁰4s²4p³

4. [Ne]3s²3p¹

The values of the first ionization enthalpies for two isotopes would be: