Atomic Radii, Colour, Catalytic Behaviour

Introduction 

The total distance from an atom’s nucleus to the outermost orbital of electrons is known as the atomic radius. In simplest terms, it can be compared to the radius of a circle, where the nucleus is at the centre and the outermost orbital of the electron is at the periphery. As you walk up and down the periodic table, you’ll notice patterns that help explain how atomic radii fluctuate.

The net positive charge felt by the valence electron is the effective nuclear charge (Zeff) of an atom. Because the core electrons hide some positive charge, the valence electron does not feel the complete positive charge. Here is a full explanation of shielding and effective nuclear charge. The atomic size of an atom is substantially influenced by Zeff As the Zeff lowers, the atomic radius increases because the electrons are screened away from the nucleus more, lowering the attraction between the nucleus and the electron. Because Zeff reduces as you move down a group and right to left across the periodic table, the atomic radius increases as you move down a group and right to left.

Types of Radius in Relation to Bond Types

The atomic radii are difficult to calculate because the position of the outermost electron is unclear — we don’t know exactly where the electron is. The Heisenberg Uncertainty Principle can explain this behaviour. We determine the radius based on the distance between the nuclei of two bound atoms to acquire a precise, but still imperfect, measurement of the radius. Atoms’ radii are thus determined by the bonds that they make. An atom’s radius varies depending on the bond it makes, hence there is no such thing as an atom’s fixed radius.

Radius of Covalent Bonding

The covalent radius can be calculated when two atoms have formed a covalent connection. When two atoms of the same element are covalently connected, the radius of each atom is half the distance between the nuclei because the electrons are drawn in opposite directions. The diameter of an atom is determined by the distance between two nuclei, but you want the radius, which is half the diameter.

Covalent radii will follow the same pattern as atomic radii in increasing. This pattern is due to the fact that the larger the radii, the greater the distance between the two nuclei. For more information, see the Zeff explanation.

Ionic Radius

The radius of an atom creating an ionic bond or an ion is called the ionic radius. In an ionic link, each atom has a different radius than in a covalent bond. This is a crucial notion. The fact that the atoms in an ionic bond are of vastly different sizes accounts for the variation in radius. One of the atoms is a cation, which is smaller than the other, while the other is an anion, which is much larger. To account for this disparity, one must first calculate the entire distance between the two nuclei and divide it by the atomic size. The larger the atomic size, the greater the radius. 

The ionic radius of cations is lower than that of neutral atoms. Anions, on the other hand, have larger ionic radii than their neutral counterparts.

The following is a full explanation:

1.The cation, which is a positive-charged ion, contains fewer electrons than protons by definition. In comparison to the neutral atom of interest, the loss of an electron will result in a shift in atomic radii (no charge).

2.The loss of an electron means that the atom now has more protons than electrons, as indicated previously. Because there are fewer electrons for the protons to drag towards the nucleus, the atomic size will shrink, resulting in a stronger attraction of the electrons towards the nucleus. It will likewise drop when the number of electrons in the outer shell decreases, resulting in a smaller radius.

3.A magnet and a metallic object might be used as an analogy. If ten magnets and ten metallic objects represent a neutral atom, with the magnets representing protons and the metallic objects representing electrons, then removing one metallic object, which is equivalent to removing an electron, will cause the magnet to pull the metallic objects closer together because the number of metallic objects will decrease. This can also be said of protons dragging electrons closer to the nucleus, resulting in a reduction in atomic size.

As positive ions are generated, the ionic radius falls.

An anion, on the other hand, will be larger than the atom from which it was formed due to the gain of an electron.The addition of an electron to the outermost shell increases the radius because there are now more electrons further away from the nucleus and more electrons to pull towards the nucleus, causing the pull to become slightly weaker than that of a neutral atom, resulting in an increase in atomic radius.

Metallic Radius

The radius of an atom linked by a metallic connection is known as the metallic radius. In a metallic cluster, the metallic radius is half of the total distance between the nuclei of two adjacent atoms. The distance between each atom in a metal will be the same because it is made up of atoms of the same element.

Atomic Radius trend in  Periodic Trends

As the number of electronic shells on an atom grows, the radius of the atom increases as you move down the periodic table of elements.

In general, as you move from the left to the right of a period, the size of an atom shrinks. Effective nuclear charge rises with time as electron shielding remains constant. A higher effective nuclear charge attracts electrons more strongly, drawing the electron cloud closer to the nucleus and reducing the atomic radius.

Vertical Movement 

As we  progress through a group, the radius of atoms grows.

Horizontal Movement

As you move from the left to the right side of a period, the size of an atom shrinks.

EXCEPTIONS: The nucleus draws electrons inward because the electrons added in transition elements are added to the inner electron shell while the outer shell remains constant. This phenomenon is explained by the electron configuration of transition metals. 

The electronic configuration (E.C.) of Ga is:

[Ar]3d104s24p1

The screening effect (or shielding effect) for the outermost electrons reduces due to the low shielding effect of d electrons. As a result, the effective nuclear charge on the outer electrons rises, causing the size of the atom to shrink. As a result, the sizes of Ga and Al are nearly identical.

The order of penetration power of orbitals are: s>p>d>f

As a result, if the inner electrons are in the s -orbital, they will have a larger screening effect than those in the p -orbital, weakening the screening effect.

This explains why Ga is the same size as the atom before it and Sb is somewhat larger than Sn.

Conclusion 

When attempting to describe the behaviour of atoms or compounds, the size of atoms is critical. The atomic radius is one of the ways we can express the size of atoms. This information explains why some molecules fit together and why others have portions that become too crowded under specific circumstances.