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When two similar atoms are joined together by the bond then the atomic radius is one-half of the distance between the nuclei of the two atoms. As we move in the periodic table from top to bottom in groups and left to right in periods, a change in the atomic radius is observed. From left to right, the period size of the atoms generally decreases. But there are some exceptions observed such as the atomic radius of the oxygen atom is slightly more than the radius of the nitrogen atom. From top to bottom in a group, the size of atoms generally increases. Because of two reasons, one of them is that the atomic number increases down the group so, the positive nuclear charge increases, And other is, due to an increase in principal quantum number (n).
Atomic radius is the distance between the nucleus and the outermost shell of the electron.
As we move in the periodic table from one element to the next across a period, nuclear charge (Z) increases by one, but screening due to inner electrons increases only slightly.
Down the group, atomic radii increase because of the addition of an extra shell. Thus electrons are added in that outer orbital which increases the distance from the nucleus.
In this article, we will learn more about Atomic radii, Types of atomic radii, Trends of atomic radii, Effective nuclear charge and Shielding constant.
Types of Atomic Radii
(i) Covalent radius: It is half of the distance between two covalently joined atoms. For covalently bonded atoms we use a covalent radius.
(ii) Van der Waals radius: It is half of the distance between the two atoms belonging to the neighboring molecules. For example, the internuclear distance of two chlorine atoms is 360 pm. So, the Van der Waals radius of the chlorine is 180 pm.
(iii) Metallic Radius: It is the minimum distance between two nuclei of metal ions when they are adjacent to each other in a metal lattice. For example, the metallic radius of sodium is 186 pm whereas the covalent radius of sodium is just 154 pm.
Effective Nuclear Charge (Zeff):-It is the charge experienced by an electron in a polyelectronic atom after the screening effect of the inner shell electron.
Zeff = Z – shielding constant
Nuclear charge (Z):- It is the total number of protons present in the nucleus.
Shielding Constant / Screening constant: – It is the repulsive force experienced by the valence shell from the inner electrons.
Order of screening constant of orbital is:
s > p > d > f
This means that the electrons in the inner orbital s screen the electrons in the valence shell more effectively than p orbital and so on.
a) Trend of the periodic table in period:- In general, the atomic radii decrease with increase (from left to right) in a period. For example, in the second period, the atomic radii decrease from Li to Ne through Be, C, N, O, and F as given in the table;
This is explained based on increasing nuclear charge along a period. With the increase in atomic number from lithium to fluorine, the magnitude of the nuclear charge increases progressively by one unit while the corresponding addition of electrons takes place in the same principal shell. Since the electrons in the same shell do not screen each other from the nucleus, the increase in nuclear charge is not neutralized by the extra valence electron. As a result, electrons are pulled closer to the nucleus by the increased effective nuclear charge and thereby, cause a decrease in the size of the atom.
Thus, effective nuclear charge (Zeff) increases across the period from left to right. Due to this reason elements on the right side of the periodic table are small in size because they experienced the higher effective nuclear charge.
b) Trend of the periodic table in the group:- The atomic radii of elements increases from top to bottom in a group. In moving down a group, the nuclear charge is increasing as the atomic number increases and we expect that the size of the atom should decrease. However, while going from one atom to another, there is an increase in the principal quantum number. The effect of an increase in the size of the electron cloud is more pronounced than the effect of an increased nuclear charge. In other words, the size of the atoms goes on increasing as we move down the group.
Conclusion
The atomic radii of elements along with the period decrease due to the nuclear charge increase as a result electrons are pulled closer to the nucleus by the increased effective nuclear charge and thereby, causing a decrease in the size of the atom. Whereas down the group atomic radii increase due to an increase in principal quantum number (n).
The radius of the cation is always smaller than the atom because the effective nuclear charge per electron increases and the electrons are more strongly attracted and are pulled toward the nucleus. Due to this size of the ion decreases.
