Atoms do not have sharp boundaries because the electrons around nucleus are diffused like cloud. Because of this, it is not possible to define the exact size of an atom.
One way to describe atomic size is by looking at the distance between two atoms that are not bonded but touch each other. For example, in solid krypton, atoms come into contact without forming bonds. The distance between the nuclei of two krypton atoms can be calculated from the density of the solid. Half of this distance is taken as the atomic radius. This type of atomic radius is called the non-bonding atomic radius or the van der Waals radius. It represents the size of an atom when it is not bonded to any other atom. For instance, the distance between two krypton nuclei in the solid state is about 400 pm, so its van der Waals radius is approximately 200 pm.
Another type of atomic radius is the bonding atomic radius, also called the covalent radius. Its definition is slightly different for nonmetals and metals. In nonmetals, it is half the distance between the nuclei of two identical atoms bonded by a covalent bond. In metals, it is half the distance between the nuclei of two adjacent atoms in a metallic crystal. For example, in a bromine molecule (Br₂), the distance between the nuclei of two Br atoms is 228 pm. Thus, the covalent radius of bromine atom is half of the bond distance between two bromine atoms which is 114 pm.
Similarly, in a chlorine molecule (Cl₂), the distance between the two nuclei is 198 pm, which gives a covalent radius of 99 pm. By this method, atomic radii can be assigned to most elements that form covalent bonds or metallic crystals.
A more general term, atomic radius, refers to the average bonding radius obtained from many experimental measurements of elements and compounds. The atomic radius shows the size of an atom when it is bonded to another atom. It is always smaller than the van der Waals radius, because bonded atoms are pulled closer together, whereas non-bonded atoms remain farther apart.
The concept of atomic radius is also useful for estimating bond length. The bond length between two covalently bonded atoms can be approximated as the sum of their atomic radii. For example, in iodine monochloride (ICl), the atomic radius of iodine is 133 pm and that of chlorine is 99 pm. Adding these gives a predicted bond length of 232 pm. The experimentally measured bond length is 232.07 pm, which closely matches the calculated value.
Comparison of Atomic Radii
| Type of Radius | Definition | Example | Relative Size |
| Covalent Radius | Half the distance between nuclei of two bonded identical atoms | Br = 114 pm; Cl₂ = 99 pm | Smaller |
| Metallic Radius | Half the distance between nuclei of two adjacent metal atoms in a crystal | Na = 186 pm | Moderate |
| van der Waals Radius | Half the distance between nuclei of two non-bonded atoms in contact | Kr = 200 pm | Largest |
Periodic Trends
Trend Along Periods
When we travel along periods from left to right, atomic radius decreases. The reason is that nuclear charge (number of protons) increases which brings electrons closer to the nucleus.
Trend Down the Groups
Down the groups, atomic radius of elements increases because number of electronic shells increases which , in turn, increases shielding effect. Nuclear charge also increases but the electronic shell factor dominates it.
Difference Between Nuclear Charge and Effective Nuclear Charge
The electrons around the nucleus of an atom experience an attractive force from the positively charged protons inside the nucleus. At the same time, however, electrons also repel one another (especially inner shell electrons repel outer shell electrons) because they all carry a negative charge. This repulsion shields or screens the outer electrons from the full attractive force of the nucleus-also called shielding or screening effect. As a result, the outermost shell electrons do not feel the actual force of attraction from the nucleus. This actual and net positive charge felt by these electrons is called the effective nuclear charge.
Periodic Trends in Transition Metals
Going across a period in transition metals, the trends in atomic radii is different from that explained above (group IA to VIIA). In transition metals, there are two factors working at the same time.
One is that nuclear charge increases in the same way as above which pulls electrons closer to the nucleus, decreasing atomic radii. The other factor is that electrons added to the atoms go into inner shell, (n-1)d, rather than outermost shell , ns. These electrons cause shielding the outermost shell electrons from the nucleus, which might be expected to increase the atomic radius.
In reality, the atomic radii of transition metals decrease only slightly because the increase in nuclear charge is almost counterbalanced by the shielding effect of the additional inner-shell electrons. There are, however, a few exceptions to this trend, which are beyond the scope of this course.
Frequently Asked Questions
Q1: Why can’t we measure atomic radius exactly?
Because orbitals do not have fixed boundaries; they only give a probability of where electrons are likely to be found.
Q2: Which radius is larger — covalent or van der Waals?
The van der Waals radius is always larger, since non-bonded atoms are farther apart than bonded atoms.
Q3: Can bond length always be predicted by adding atomic radii?
Yes, approximately. The bond length is usually close to the sum of the radii, though actual experimental values may differ slightly.
Q4: Which elements usually have their covalent radii measured?
Mainly nonmetals and metals that form covalent bonds or metallic crystals. Noble gases rarely form covalent bonds, so their van der Waals radii are used instead.
Q5: Why is atomic radius important?
It helps us understand bond lengths, bond strengths, periodic trends, molecular size, and physical properties such as melting and boiling points.