Atomic Size & Atomic Radius - Practice Questions & MCQ

There are various physical properties that follow the trend according to the atomic numbers of elements. 

• Atomic Radius
  Atomic radius is the distance from the centre of the nucleus to the outermost shell of electrons. It can be measured by the x-ray diffraction, electron diffraction techniques and other spectroscopic methods. As it cannot be measured in absolute form thus it is measured in various definitions such as covalent radius, Van der Waal’s radius, metallic radius and ionic radius.

• Covalent Radius
  It is half of the total length between two successive nuclei covalently bonded to each other in a molecule. Suppose there are two same atoms’ A’ and ‘A’ in a molecule and their bond length is ‘a’, then covalent radius is half of the covalent bond length between A and A. Thus, covalent radius = (a/2).

• Metallic Radius

It is defined as half of the distance between nuclei of two adjacent metal atoms that are closely packed in the metallic crystal lattice. For example, there are two metals atoms ‘A’ and ‘A’ that are closely packed to each other and bond length is ‘a’ then the metallic radius is half of the distance between these two metallic atoms i.e, a/2.

• Van der Waal’s Radius

It is half of the distance between two nuclei of the adjacent non-bonded atoms of different molecules. For example, if ‘a’ is the distance between two adjacent atoms i.e, A and B, then the Van der Waals’s radius is the half of the distance between these two atoms A and B, i.e, a/2.

• Ionic Radius

It is the effective distance from the centre of the nucleus of an ion upto which it has the influence over the electrons.

Comparison of the ionic radii and atomic radii

Variation in a Period

In moving from left to right in a period, the nuclear charge increases and the last electron enters into the same shell, thus the effective nuclear charge increases. Thus in this way, the atomic size decreases in the period.

Variation in a Group

In moving from top to bottom in a group, the number of shells increases due to which the atomic size increases. 

• The size of the cation is always smaller than its parent atom. In the case of cations, the number of electrons in the ion decrease and the nuclear charge remains the same, thus the effective nuclear charge increases and the size decreases.

• The size of the cation decreases  as effective nuclear charge increases

• M⁺³ < M⁺² < M⁺ < M 

• The size of the anion is always greater than its parent atom. In the case of anions, the number of electrons in the ion increase and the nuclear charge remains the same, thus the effective nuclear charge decreases and the size increases.

• M^-3 > M^-2 > M^- > M

Atomic Properties

Atomic properties are the physical properties of elements that are related to the atomic number of the elements. These properties can be divided into two categories:

1. Properties of individual atoms: These are the properties of individual atoms that are directly dependent on their electronic configurations. Some examples include ionisation enthalpy, electron gain enthalpy, screening effect, effective nuclear charge, etc.

2. Properties of the group of atoms: These are properties of the group of atoms together that are indirectly related to their electronic configurations. Some examples include the melting point, boiling point, the heat of fusion, density, etc.

The Screening effect or Shielding effect

The decrease in the force of attraction between the outer electrons and the nucleus due to the presence of inner electrons is called the screening effect or shielding effect. Actually, these inner electrons generate the repulsion between these inner electrons and the outer electrons due to which the net force of attraction between the nucleus and the outer electrons decreases.

Calculation of the screening effect

• For ns or np orbital electrons

• All electrons in the (ns, np) group contribute to 0.35 each to the screening effect constant. Except for 1s electrons which contribute by 0.30.

• All electrons in (n-1) shell contribute by 0.85 each to the screening effect constant. 

• All electrons in (n-2) shell or lower contribute by 1.0 each to the screening effect constant. 

• For d- or f-electrons

• All electrons in the (ns, np) group contribute to 0.35 each to the screening effect constant.

• All the electrons in groups lower than (nd, nf) contribute by 1.0 each to the screening effect.

 Effective Nuclear Charge

Due to the screening effect of the inner or the same shell electrons, the net force of attraction between the nucleus and the outer electrons decreases. This decreased force of attraction is known as effective nuclear charge. It is represented by Z*. Mathematically, it can be formulated as:

Z* = (Z- ), where  is the screening effect constant.

• In a period, the effective nuclear charge increases in moving from left to right.

• In a group, the effective nuclear charge almost remains the same.

Isoelectronic species -

A series of atoms, ions and molecules in which each species contains the same number of electrons but a different nuclear charge.