explain how the apparent nuclear strength and the shielding effect influences periodic trends.?

Answer 1

well, as you go down the table, the extra shells of electrons cause shielding of force applied on the electrons by the protons. this causes three effects that i know of. 1. the electrons are easier to remove. 2.the atom is less likely to attract another electron. 3. and the atoms bond radius increases.
as you move across the table the greater number of protons and electrons(without adding shells)causes the attraction of electrons from the protons to increase. and causes three effects that i know of. 1.the electrons are harder to remove. 2.the atom is more likely to attract electrons. 3.and it causes the atom's bond radius to become smaller

Answer 2

(1) As you go from left to right along a row, the atomic radius decreases. This is because the number of protons in the nucleus increases by 1 as you move from one element to the next, so the nucelus has an increased positive charge, which attracts the electrons more 'strongly' for lack of a better word, and pulls them closer to the nucleus, hence the radius decreases. (2) As you go down a column, the atomic radius increases, This is because although the nuclear charge is increasing, the electron shells/orbitals are another 'level' so to speak further away from the nucleus, and so the attraction is not as strong. Also, the inner 'levels' of electrons shield the outer 'levels' from the full attraction of the nucleus, hence again why the atomic radius increases. Hope that helps!