Periodic Table Trends
The atomic radius of an element is half of the distance between the centers of two atoms of that element that are touching each other. Generally, the atomic radius decreases across a period from left to right and increases down a given group. The atoms with the largest atomic radii are located in group 1 and at the bottom of groups.
Moving from left to right across a period, electrons are added one at a time to the outer energy shell. Electrons within a shell cannot shield each other from the attraction of protons. Since the number of protons is also increasing, the effective nuclear charge increases across the period. This causes the atomic radius to decrease. Moving down a group in the periodic table, the number of electrons and filed electron shells increases, but the number of valence electrons remains the same. The outermost electrons in a group are exposed to the same effective nuclear charge, but electrons are found farther from the nucleus as the number of filled energy shells increases. Therefore, the atomic radius increases.
The ionization energy or ionization potential is the energy required to completely remove an electron from a gaseous atom or ion. The closer and more tightly bound an electron is to the nucleus, the more difficult it will be to remove, and thus the higher its ionization energy will be. The first ionization energy is the energy required to remove one electron from its parent atom. The second ionization energy is the energy required to remove a second valence electron. The second ionization energy is always greater than the first ionization energy.
Ionization energies increase moving from left to right across a period (deceasing atomic radius). Ionization energy decreases moving down a group (increasing atomic radius). Group 1 elements have low ionization energies because the loss of an electron forms a stable octet.
Electron affinity reflects the ability of an atom to accept an electron. It is the energy change that occurs when an electrons is added to a gaseous atom. Atoms with stronger effective nuclear charge have greater electron affinity. Group IIA elements and alkaline earth metals have low electron affinity values. These elements are relatively stable because they have filled “s” subshells. Group VIIA elements, the halogens have high electron affinities because the addition of an electron to an atom results in a completely filled shell. Group VIII elements, noble gases, have electron affinities near zero, since each atom possesses a stable octet and will not accept an electron readily. Elements of other groups have low electron affinities.
In a period, the halogens will have the highest electron affinity; while the noble gas will have the lowest electron affinity. Electron affinity decreases moving down a group because a new electro would be further from the nucleus of a large atom.
Electro-negativity is a measure of the attraction of an atom for the electrons in a chemical bond. The higher the electro-negativity of an atom, the greater its attraction will be for bonding electrons. Electro-negativity is related to ionization energy. Electrons with low ionization energies have low electro-negativities because their nuclei exert a strong attractive force on electrons. Elements with high ionization energies have high electro-negatives due to the strong pull exerted on electrons by the nucleus. in a group, the electro-negativity decreases as atomic number increases, as a result of increased distance between the valence electron and nucleus (greater atomic radius).
Summary of trends
Moving left to right
- Atomic radius decreases
- Ionization energy increases
- Electron affinity generally increases
- Electro-negativity increases
Moving top to bottom
- atomic radius increases
- ionization energy decreases
- electron affinity generally decreases
- electro-negativity decreases