What trends do you see on the periodic table for Ionization energy? What causes these trends?

As we face the Periodic Table, we find that ionization energy DECREASES down a Group, a column, and INCREASES across a Period, a horizontal row of the Table.

This trend is influenced by two factors: shielding by other electrons and nuclear charge.

See here and here. Once a valence shell is completed, shielding becomes reasonably effective, and electrons add to a valence shell farther removed from the nuclear core. Incomplete electronic shells shield the nuclear charge VERY ineffectively, and thus across the Period, the atomic radii markedly contract - this is a macroscopic variable that you should investigate.

The trend for ionization energy generally increases across a period and decreases down a group on the periodic table. This is due to the increase in effective nuclear charge across a period and the increase in atomic size down a group.

The trends observed in ionization energy on the periodic table generally include:

1. Ionization energy tends to increase across a period (from left to right) and decrease down a group (from top to bottom).
2. There are exceptions to the general trend, notably with the group 2 elements and the group 15 elements, which have slightly lower ionization energies than would be expected based solely on their position in the periodic table.
3. Transition metals generally have lower ionization energies compared to main group elements in the same period.

These trends can be attributed to the following factors:

1. Effective nuclear charge: As you move across a period, the effective nuclear charge (the positive charge felt by the outermost electrons) increases due to the increase in the number of protons in the nucleus. This greater nuclear charge exerts a stronger pull on the outermost electrons, making them more difficult to remove, thus increasing ionization energy.

2. Shielding effect: Although the number of protons increases across a period, the number of inner shell electrons also increases, providing some shielding or repulsion to the outermost electrons from the increased nuclear charge. This shielding effect partially offsets the increase in ionization energy.

3. Distance from the nucleus: As you move down a group, the outermost electrons are located in higher energy levels, which are farther from the nucleus. The increased distance weakens the attraction between the outermost electrons and the nucleus, making it easier to remove the outermost electron, thus decreasing ionization energy.

4. Sublevel and orbital stability: Sublevel and orbital stability also play a role in determining ionization energy. Half-filled and fully filled orbitals are more stable, so elements with configurations that result in half-filled or fully filled orbitals may have lower ionization energies than neighboring elements.

Overall, these trends in ionization energy reflect the interplay between effective nuclear charge, shielding effects, and electron distance from the nucleus, along with considerations of sublevel and orbital stability.