Periodic Trends
1. Atomic Radius

Horizontal: The atomic radius of an element decreases from left to right within each period on the periodic table.

Explanation: In understanding this trend, one must consider the charge and energy of the valance electrons. As you move from left to right on the periodic table, the number of protons increases in increments of one. In addition to this, the number of electrons also increases by increments of one. These electrons are added to the same energy level (or ring). Therefore, the increased attraction of the nucleus pulls the electron closer to the nucleus, so the atom becomes smaller.

Vertical: The atomic radius increases within a group as you descend the periodic table.

Explanation: The atomic radius increases descending down a group of elements on the periodic table. Therefore, each row on the table contains one more energy level (or ring) than the previous row. The farther an energy level is from the nucleus, the less attraction force is exerted on the electrons in the energy level. Also, the more rings an atom has, cause a greater amount of shielding, since each ring contributes to a greater shielding effect.

2. First Ionization Energy

Horizontal: As the atomic number of elements in the periodic table increases from left to right, the first ionization energy also increases from left to right. This means that the elements of Group 1 (alkali metals) have the lowest first ionization energy, and Group 18 elements (noble gases) have the highest first ionization energy.

Explanation: When moving from left to right, the first ionization energy generally increases. This is due to the decreased atomic radius coupled with the increased effective nucleus charge. As a result, the force of attraction between the nucleus and outer most electrons increases, which leads to an increase in the amount of energy required to move an electron from the atom ( i.e., an increase in the first ionization energy).

Vertical: First Ionization energy decreases as you descend down the periodic table.

Explanation: Ionization energy decreases as you as you move down a family or group in the periodic table because the valence electrons of the elements are in orbits that are farther from the nucleus. As the distance from the nucleus increases, the attraction of the electrons to the nucleus decreases, even though the size of the nuclear charge increases. Thus less energy is required to remove an electron from an atom further down a group. (i.e., the ionization energy decreases down a group).

3. The Group that has the 2nd highest ionization energy is the Halogen group. This is due to the fact that the atomic radius of the elements decrease as you move from right to left on the periodic table, And smaller the atomic radius is, the greater the amount of ionization energy is required. The Halogen group is second to the right of the periodic table, and thus has the 2nd highest ionization energy.

4. Electron Affinity

Horizontal: The electron affinity for an element increases as you move from left to right across the periodic table.

Explanation: As you move from left to right across the periodic table, the number of valence electrons increases and the atomic radius decreases. The force of attraction between the nucleus and valence electron increases, so more energy is released when a new electron is acquired.

Vertical: The electron affinity decreases as you move down a group within the periodic table.

Explanation: As you descend the periodic table in a given group the atomic radius increases therefor the force of attraction between the valence electron and the nucleus of the atom decrease. This leads to a decreasing electron affinity within a group as you move down the periodic table.

Follow-up Practice Questions:

1. Smallest radius to largest radius: O (oxygen), Sb (antimony), Sn (tin), Ba (barium), Cs (cesium).

2.