The energy of an electron in the ground state of the hydrogen atom is $–2.18\times {10}^{-18}J$.  The ionization enthalpy of atomic hydrogen in terms of J ${\mathrm{mol}}^{-1}$ is -

1. 2.81 × 10⁶ J mol${}^{-1}$

2. 1.31 × 10⁶ J mol${}^{-1}$

3. 2.31 × 10⁶ J mol${}^{-1}$

4. 1.81 × 10⁶ J mol${}^{-1}$

The explanation for the fact that the first ionization enthalpy of sodium is lower than that of magnesium but its second ionization enthalpy is higher than that of magnesium would be -

1) pressure and volume

2) lustre and brightness

3) atomic size and effective nuclear charge

4) availability in nature

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.

An element with higher negative electron gain enthalpy in the given pair is-

(i) O or F  (ii) F or Cl

The relationship between 1^st and 2 ^nd electron gain enthalpy of oxygen atom (O) would be :

The statement that "the electronegativity of N on the Pauling scale is 3.0 in all the nitrogen compounds" is incorrect because -

1. Electronegativity of an element is a variable property.

2. Pauling scale is not used to measure electronegativity.

3. The electronegativity of N on the Pauling scale is 12.0

4. None of the above.

The correct statement about radius is :

The values of first ionization enthalpies for two isotopes would be -

The reactivity of alkali metals increases, whereas halogen decreases down the group, because -

The general outer electronic configuration of s, p, d , and f-block elements respectively would be :