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Atomic radius is the distance from the nucleus to the outermost orbital of an atom. We can grasp the concept by comparing it to a circle’s radius. The centre of the circle is the nucleus here. There are three types of atomic radii- Covalent radius, Metallic radius, and Ionic radius. Suppose a covalent bond is present between two similar atoms. Then, the covalent radius is half of the single bond length between the two similar atoms that are covalently bonded.
Covalent bonds
Langmuir (1919) redefined Lewis’s postulations by introducing covalent bonds. According to him, when two atoms share one pair of electrons, they are in a covalent bond. If they are pairing more than one pair of electrons, they can be in double or triple covalent bonds. When a covalent bond is formed, we can obtain a covalent radius too.
Overview of covalent radii
As per the theory, we assume that the distance between nuclei and the outer shell will be half of the distance between two nuclei in covalently bonded atoms. This is postulated on the idea that the nuclei will have an equal attraction to their electrons. So, when the bond is formed between two similar atoms with the same number of electrons, the atomic radius would be half the distance between nuclei. In that case, the distance between two nuclei will be the atom’s diameter. That’s why we say, ‘covalent radius is the measure of the size of the atom.’
Since electrons are practically invisible and small, the atomic size can be calculated using the distance between atoms when they are in a combined state. X-ray techniques measure this distance between two nuclei, known as bond length.
For example, the bond distance in the chlorine molecule (Cl2 ) is 198 pm, and half this distance (99 pm) is taken as the atomic radius of chlorine.
The bond length in a covalent molecule AB can be depicted as
R= rA +rB (where, R is the bond length, and rA and rB are the covalent radii of atoms A and B, respectively)
Atomic radius and trend in the periodic table
Atomic radius is the distance between the nucleus and the valence shell in an atom, and it can be covalent, metallic, or ionic.
- As we move from left to right in the periodic table, the atomic radii decrease and the force of attraction between the nucleus and electrons increases. This causes small atomic sizes.
- Whereas, moving down the group, the atomic radius increases due to two reasons.
- Increase in the number of shells
- Shielding effect, which causes a decrease in the force of attraction
Covalent radius and trends in the periodic table
Similar to the atomic radius,
- As we move from left to right, the covalent radius decreases due to greater nuclear charges, which will pull the electrons close to the nucleus.
- The covalent radius increases down the group due to the increased number of occupied shells with electrons.
Covalent radii notes
The covalent radius is measured using X-ray diffraction, and Rotational spectroscopy gives a much more accurate value of bond length. Rarely do we use neutron diffractions on molecular crystals.
It is usually measured in picometre (pm) or Angstroms (Å).
We assume that the sum of covalent radii between two atoms should equal the covalent bond length between those two atoms. But practically, it is not possible because atomic size might be different in a different environment. Therefore, the covalent bond length also differs. This explains the phenomenon of shortness of polar covalent bonds.
Tabulated covalent radii are prepared from thousands of experimental data. But they are only either idealised or averaged values.
Conclusion
Atomic radius is the distance between the nucleus and the valence shell in an atom, and it can be covalent, metallic, or ionic. Covalent radius is half the single bond length between two similar atoms that are covalently bonded. The covalent radius is measured using X-ray diffraction.
We assume the sum of covalent radii between two atoms should equal the covalent bond length between those two atoms. But practically, it is not possible because the atomic size might be different in a different environment. Therefore, the covalent bond lengths also differ, disrupting the atomic size measurement.
