As you move from left to right across a period in the periodic table, the atomic radius generally decreases. This trend occurs due to the increasing effective nuclear charge, which is caused by the increasing number of protons in the nucleus. The effective nuclear charge pulls the electrons closer to the nucleus, resulting in a smaller atomic radius.

The decreasing atomic radius can also be explained by the filling of electron shells. As you move from left to right, electrons are added to the same principal energy level, but occupy different sublevels. The increasing number of electrons leads to stronger electron-electron repulsion, causing the electron cloud to spread out and resulting in a larger atomic radius.

However, the decreasing trend in atomic radius is not consistent across all elements in a period. Beginning at group 13, there is a noticeable jump in atomic radius between the elements with electron configurations ending in p-block and those with d-block configurations. This jump is referred to as a "shielding effect," which occurs due to the insertion of electrons into d orbitals. The addition of d-block electrons can partially shield the outermost s electrons from the increasing nuclear charge, resulting in a larger atomic radius than would generally be expected.

Overall, while there is a general trend of decreasing atomic radius from left to right across a period, specific exceptions occur due to the shielding effect and other factors.

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