1. Trends in the Periodic Table
Periodicity
Trends in the Periodic Table

2. Atomic Radius
One half the distance between nuclei of identical atoms that are bonded together
Decrease across a period due to increasing positive nuclear charge
Increase down a group due to increasing number of energy levels (outer electrons are farther from the nucleus)

3. Ionization Energy
Ion: An atom or group of atoms that has a positive (cation) or negative (anion) charge
Ionization: The formation of an ion
Ionization Energy: The energy required to remove one electron from a neutral atom of an element, measured in kilojoules/mole (kj/mol)

4. Ionization Energy IE increases for each successive electron
Each electron removed experiences a stronger effective nuclear charge
Greatest increase in IE comes when trying to remove an electron from a stable, noble gas configuration.

5. Ionization Energy IE tends to increase across each period.
Atoms are getting smaller, electrons are closer to the nucleus
IE tends to decrease down a group
Atoms are getting larger, electrons are farther from the nucleus
Outer electrons becomes increasingly more shielded from the nucleus by inner electrons
Metals have characteristically low IE
Nonmentals have high IE
Noble gases have a very high IE

7. Ionic Radius
Cations have a smaller ionic radius than corresponding atom
Protons outnumber electrons
Less shielding of electrons
Anions have a larger ionic radius than corresponding atom
Electrons outnumber protons
Greater electron-electron repulsion
Ion size decreases down a group

8. Electron Affinity
Energy change that occurs when an atom accepts an electron
Halogens have the highest electron affinities
Metals have low electron affinities
Affinity tends to increase across a period and decrease down a group

9. Electronegativity The ability of an atom to attract electrons
Nonmetals have high electronegativity
Metals have low electronegativity
Electronegativity increases across a period and decreases down a group