Periodic Trends

Zeff

 Effective Nuclear Charge - The amount of charge from the nucleus that is effectively able to reach the outside (valence) electrons. 

 As you go across a period (from left to right), the Z_eff increases because you are adding protons to the nucleus, but you are not adding any shielding electrons. 

As you go down a group, Z_eff stays approximately the same because you are adding the same number of protons and shielding electrons. 

Atomic Radius

The distance from the nucleus to the outermost (valence) electrons of an atom. Analogous to the size of an atom. 

 As you go across a period (from left to right), the atomic radius decreases. This is because elements on the right have a higher Z_eff (more charge coming from the nucleus). The higher Z_eff leads to a stronger force of attraction, which pulls the electrons in closer to the nucleus. 

As you go down a group, the atomic radius increases because more energy levels (electron shells) are added to the atom. 

Ionization Energy

The amount of energy required to remove the outermost (valence) electron from an atom.

As you go across a period (from left to right), ionization energy generally increases. Higher Z_eff and smaller atomic radius create a stronger attraction between the nucleus and valence electrons, making them harder to remove.

As you go down a group, ionization energy generally decreases. Z_eff stays approximately the same, but additional energy levels place the valence electrons farther from the nucleus. The weaker attraction makes them easier to remove. 

B < Be & Al < Mg: The electron removed from B or Al is in a higher-energy p sublevel and is partially shielded by the s electrons, making it easier to remove than expected.

• O < N & S < P: O and S have paired p-sublevel electrons. The extra electron-electron repulsion makes one of the paired electrons easier to remove.

Electron Affinity

 The amount of energy released when an electron is added to an atom. Electron affinities are usually listed as negative values; the more negative the value, the more energy is released.

As you go across a period (from left to right), electron affinity generally becomes more negative. Higher Z_eff and smaller atomic radius cause a stronger attraction for an added electron.

As you go down a group, electron affinity generally becomes less negative. Z_eff  stays approximately the same, but additional energy levels place the added electron farther from the nucleus. The weaker attraction means less energy is released when the electron is added. 

Group 2: The s sublevel is filled. An added electron must enter the higher-energy p sublevel, making electron addition less favorable.

• Group 15: The p sublevel is half-filled. Adding an electron requires pairing electrons, increasing electron-electron repulsion.

• Group 18: The valence energy level is filled. An added electron must enter a new, higher energy level, making electron addition very unfavorable and sometimes endothermic.

Electronegativity

The ability of an atom in a chemical bond to attract the shared electrons toward itself.

As you go across a period from left to right, electronegativity generally increases. Higher Z_eff  and smaller atomic radius create a stronger attraction for the shared electrons in a bond. 

As you go down a group, electronegativity generally decreases. Z_eff  stays approximately the same, but additional energy levels place the bonding electrons farther from the nucleus. As a result, the atom attracts shared electrons less strongly. 

Noble gases generally are not assigned electronegativity values because they rarely form bonds. Electronegativity trends are most useful for comparing main-group elements.