Question

The first three successive ionization energies of an element Z are $900,1757,14850{\text{kJ}}.{\text{mo}}{{\text{l}}^{ - 1}}$ respectively. The element Z belongs to:
A. group-$1$
B. group-$2$
C. group-$3$
D. group-$4$

Answer 1

Hint: Ionization energy is the certain amount of energy that is necessary to knock off electrons from an atom to form a positive ion. They are expressed in electron volts per atom or kilocalories per mole or kilojoules per mole.

Complete step by step answer:
The energy needed to remove one or more electrons from a neutral gaseous atom and convert it into a positively charged gaseous ion. On moving from left to right in the periodic table, ionization energy increases. When atomic number increases, nuclear charge increases. Electrons are held tightly hence greater energy is required to remove electrons.
On moving down the group, ionization energy decreases. Cations are smaller than the neutral atom. Down the group, electrons are loosely held, hence lesser energy is required to remove electrons.
The energy gap between the ionization energies help us to guess the valency of an element. From the given successive ionization energy values, the energy gap between second and third ionization energy is comparatively higher than first and second ionization energy. This tells that the third electron is more close to the nucleus and less shielded by inner shells because these factors will increase the force of attraction and require more energy to overcome.
Thus the element will be in group-$2$. So it has two valence electrons in its outermost shell. It is easy to remove the first valence electron, thus low energy. It needs comparatively more energy to remove the last valence electron in the valence shell. When this is removed, it becomes stable. To remove electrons further, it is difficult. Thus needs higher energy.

Hence the correct option is B.

Note:
Ionization energy depends upon the following factors:
-Atomic size
-Number of protons
-Nature of orbital
-Shielding effect