Covalent radius, definition, examples, determination and periodic trend

Covalent Radius definition

Half of the single bond length between two similar atoms covalently bonded in a molecule is called the covalent radius.

What is a covalent Radius?

The covalent radius is a measure of the size of an atom that forms part of a covalent bond. In a covalent bond, atoms share electrons to achieve a more stable electron configuration. The covalent radius is defined as half the distance between the nuclei of two identical atoms that are bonded together.

It’s important to note that the covalent radius is an empirical measure, and there are different ways to define and measure atomic sizes, depending on the context and the type of bond involved. Different methods and data sources may yield slightly different values for covalent radii.

The concept of covalent radius is particularly relevant in covalent compounds, where atoms are linked by shared electrons. It helps to estimate the size of the atoms involved in the bond. Keep in mind that covalent radii are distinct from atomic radii, which are typically defined in the context of individual, isolated atoms.

Examples of Covalent Radius

Element	Covalent radius (A°)
Hydrogen	0.37 A°
Carbon	0.77 A°
Nitrogen	0.75 A°
Oxygen	0.73 A°
Fluorine	0.72 A°
Sodium	1.54 A°
Magnesium	1.36 A°
Aluminum	1.18 A°
Silicon	1.17 A°
phosphorus	1.10 A°

Determination of covalent radii:

Actually, the covalent radii are not constant but vary with the character of the covalent bond between the atoms. If the two bonded atoms vary in electronegativity value, the covalent radius will be smaller than the expected value.

The covalent radius of an atom can be used to determine the covalent radius of another atom.

Example 1: Calculating the Covalent Radius of Hydrogen (H)
Given:

A hydrogen atom has 1 electron and 1 proton.
The distance between the nucleus and the electron in a hydrogen atom is approximately 0.53 Angstrom (Å).
To calculate the covalent radius of hydrogen, we can use the formula:
Covalent Radius = (Distance between Nuclei) / 2

Covalent Radius = (0.53 Å) / 2 = 0.265 Å

So, the covalent radius of hydrogen is approximately 0.265 Å.

Example 2: Calculating the Covalent Radius of Oxygen (O)
Given:

An oxygen atom has 8 electrons and 8 protons.
The distance between the nuclei in an oxygen molecule (O2) is approximately 1.20 Ångstroms (Å).
To calculate the covalent radius of oxygen, we can use the formula:
Covalent Radius = (Distance between Nuclei) / 2

Covalent Radius = (1.20 Å) / 2 = 0.60 Å

So, the covalent radius of oxygen is approximately 0.60 Å.

Example 3: Calculating the Covalent Radius of Carbon (C)
Given:

A carbon atom has 6 electrons and 6 protons.
The distance between the nuclei in a carbon-carbon (C-C) single bond is approximately 1.54 Ångstroms (Å).
To calculate the covalent radius of carbon, we can use the formula:
Covalent Radius = (Distance between Nuclei) / 2

Covalent Radius = (1.54 Å) / 2 = 0.77 Å

So, the covalent radius of carbon is approximately 0.77 Å.

Variation of Covalent Radii in Periodic Table

In group

Covalent radii increase from top to bottom along with a group.

The trend of increasing covalent radii from top to bottom within a group in the periodic table can be explained by the following factors:

Atomic Size

Going down a group, the number of electron shells increases. Each new shell adds to the overall size of the atom.

As you move down a group, electrons are added to the outermost shell, increasing the average distance of the outer electrons from the nucleus.

The increased number of electron shells and the greater distance of the outer electrons contribute to a larger atomic size.

Shielding Effect

The inner electron shells shield the outer electrons from the attractive force of the positively charged nucleus.

With the addition of more electron shells down a group, there is an increase in the shielding effect, which reduces the effective nuclear charge felt by the outer electrons.

A reduced effective nuclear charge makes it easier for outer electrons to move farther from the nucleus, contributing to an increase in covalent radii.

Electron Repulsion

As more electron shells are added, there is an increase in electron-electron repulsion. Electrons in the outermost shell repel each other, leading to a slight expansion of the atomic size.

This electron repulsion counteracts the attractive force of the nucleus on the outer electrons, contributing to the observed increase in covalent radii.

In period

Covalent radii decrease from left to right in a period.

The trend of decreasing covalent radii from left to right within a period in the periodic table can be explained by the following factors:

Increasing Effective Nuclear Charge

As you move from left to right across a period, the number of protons in the nucleus increases, leading to a higher positive charge in the nucleus.

The increasing positive charge attracts electrons more strongly, creating a greater effective nuclear charge. Electrons experience a stronger pull toward the nucleus.

Constant Number of Electron Shells

In a period, electrons are added to the same principal energy level (shell).

The increasing effective nuclear charge without a corresponding increase in the number of electron shells results in a stronger attraction between the electrons and the nucleus.

Shielding Effect Remains Constant

In a period, electrons are added to the same energy level, so the shielding effect remains relatively constant.

The shielding effect, caused by inner electrons, does not increase as much as the effective nuclear charge does, leading to a net increase in the attractive force felt by the outer electrons.

Decreasing Atomic Size

The combination of a stronger attractive force from the nucleus and a constant shielding effect results in a decrease in the atomic size from left to right across a period.

The outer electrons are pulled closer to the nucleus, leading to smaller covalent radii.