CHEMISTRY COMMUNITY

Can someone explain the reasons behind the trends we see in the periodic table?
- Atomic radius
- Ionization energy
- Electronegativity
How do they relate to shielding effect and number of electrons or protons?

Does anyone know why the atomic radius of Al is greater than the atomic radius of silver?

1. Atomic radius generally increases down a group and decreases across a period. Down a group, valence electrons are in shells that are further away from the nucleus and hence atomic radius increases down the group.
Across a period there is a decrease because effective nuclear charge increases and valence electrons are more strongly attracted to the nucleus. So the atom becomes more compact and the atomic radius decreases. Shielding effect remains relatively the same across a period because electrons are added to the same shell while nuclear charge increases. So effective nuclear charge increases.

2. First ionization energy generally increases across a period and decreases down a group. Down a group, valence electrons are further away from the nucleus and less tightly bound hence less energy is required to remove a valence electron from the atom and first I.E. decreases.
Across a period, effective nuclear charge increases and valence electrons are more strongly attracted to the increasingly positive nucleus so it is harder to remove a valence electron from the atom. Hence, first I.E. increases across a period.

Kristen Kim 1I wrote:Can someone explain the reasons behind the trends we see in the periodic table?
- Atomic radius
- Ionization energy
- Electronegativity
How do they relate to shielding effect and number of electrons or protons?

Atomic radius increases going down the periodic table because a valence electron is added into a new energy level and the core electrons repel each other (shielding effect) which contribute to the atomic radius increasing. Atomic radius decreases going across the periodic table because the effective nuclear charge increases. The more protons there are means that they attract the electrons more, thus making the atomic radius smaller.

Ionization energy and electronegativity have the same trend. Both increase from left to right and decrease going from top to bottom. This is the general trend but there are many exceptions based on how the orbitals are filled. If the orbitals are fully filled or half filled then it will take more energy to remove an electron from them. Since the effective nuclear charge also increases from left to right, then the attraction between the protons and electrons are greater; therefore a higher ionization energy is required to remove the electron.