The covalent radius is half the distance between two atoms that share a covalent bond. Usually, you see covalent radius in units of picometers (pm) or angstroms (Å), where 1 Å = 100 pm. For example the average covalent radius for hydrogen is 31 pm and the average neon covalent radius is 58 pm.
Why Are There Different Numbers?
When you look at a table of covalent radius values, its numbers may differ from those found on another table. This is because there are different ways of reporting covalent radius.
In reality, the covalent radius depends on an atom’s hybridization, the nature of the two atoms sharing a covalent bond, and on the chemical environment surrounding the atoms. For example, the covalent radius of carbon is 76 pm for the sp3, 73 pm for the sp2 hybridization, and 69 pm for the sp hybridization.
Also, covalent radius depends on whether the atom forms a single bond, double bond, or triple bond. In general, a single bond is longer than a double bond, which is longer than a triple bond.
A given table might generalize data or else offer values based on very specific conditions. Tables that cite an average value usually combine data for covalent bonds an atom forms in many different compounds. Some tables list the covalent radius for a homonuclear covalent bond. For example, this is the covalent radius for H2 or O2. Either use the idealized (calculated) or empirical average covalent radius for an atom for maximum transferability.
How Covalent Radius Is Measured
The most common methods of measuring covalent radius are x-ray diffraction and rotational spectroscopy. Neutron diffraction of molecular crystals is another method.
Covalent Radius Trend on the Periodic Table
Covalent radius displays a periodic table trend.
- Moving left to right across a period, covalent radius decreases.
- Moving top to bottom down a group, covalent radius increases.
Covalent radius decreases moving from left to right across a row or period because atoms gain more protons in their nucleus and electrons in their outer shells. Adding more protons increases the attractive pull on these electrons, drawing them in more tightly.
Covalent radius increases moving down a column or periodic table group. This is because increasing filled inner electron energy levels shield the outer electrons from the positive nuclear charge. So, the electrons are less attracted to the nucleus and increase their distance to it.
Covalent Radius vs Atomic Radius and Ionic Radius
Covalent radius, atomic radius, and ionic radius are three ways of measuring the sizes of atoms and their sphere of influence. The atomic radius is half the distance between the nuclei of atoms that are just touching each other, where “touching” means their outer electrons shells are in contact. The ionic radius is half the distance between two atoms touching each other that share an ionic bond in a crystal lattice.
All three measures of atomic size follow a periodic table trend, where radius generally increases in size moving down an element group and decreases in size moving from left to right across a period. However, the covalent radius and ionic radius often are different sizes from the atomic radius.
The Largest and Small Covalent Radius
The element with the smallest covalent radius is hydrogen (32 pm). The atom with the largest covalent radius is francium (223 pm when it forms a single bond). Basically, this is another way of saying hydrogen is the smallest atom and francium is the largest atom.
Covalent Radius Periodic Table
This covalent radius periodic table illustrates the periodic table trend. The table is available as a PDF for downloading and printing.
References
- Allen, F. H.; Kennard, O.; Watson, D. G.; Brammer, L.; Orpen, A. G.; Taylor, R. (1987). “Table of Bond Lengths Determined by X-Ray and Neutron Diffraction”. J. Chem. Soc., Perkin Trans. 2 (12): S1–S19. doi:10.1039/P298700000S1
- Cordero, B.; Gómez, V.; et al. (2008). “Covalent radii revisited”. Dalton Transactions. 21: 2832-2838. doi:10.1039/B801115J
- Pyykkö, P.; Atsumi, M. (2009). “Molecular Single-Bond Covalent Radii for Elements 1-118”. Chemistry: A European Journal. 15 (1): 186–197. doi:10.1002/chem.200800987
- Sanderson, R. T. (1983). “Electronegativity and Bond Energy”. Journal of the American Chemical Society. 105 (8): 2259–2261. doi:10.1021/ja00346a026