Periodic+Relationships


 * Atomic Radius**
 * Estimated size of an atom from inside of the nucleus to outside the electron cloud
 * About 1/2 the distance between the nuclei of 2 atoms of the same element when they are joined
 * Distance from the atomic nucleus to the valence electrons
 * The atomic radius trend going down a group comes from the increase in the principal quantum number of electrons in the atom. The farther down a column, the more likely an element's electrons will be farther from the nucleus, causing the atom to be larger.
 * The period trend comes from the increase in nuclear charge, which attracts outer electrons, pulling them closer to the nucleus of the atom. This makes the atoms smaller as you move across the rows.


 * Ionization Energy **
 * Energy required to remove an electron from a neutral atom (in other words, how difficult it is to take an electron)
 * The second ionization energy is greater than the first because it is more difficult to remove a second electron from an atom
 * Due to the fact that you are pulling from a positive charge
 * as atomic radius decreases, ionization energy increases because it becomes more difficult to take electrons away because of the stronger pull that the nucleus has on the outermost electrons
 * Measures the ability of an atom to attract electrons when the atom is in a compound
 * Measured in kJ/mol
 * The period trend for ionization energy results from an increase in atomic number, but there are some exceptions to this trend. The discrepancies are due to the electron configurations of some elements. For example, the ionization energy from beryllium to boron decreases instead of increases because the third valence electron of boron must occupy the 2//p// level but that is empty for Be.
 * The column trend is opposite of this, and decreases while the atomic number increases.

Periodic Table courtesy of: [|www.elementsdatabase.com] **Electron Affinities**
 * Energy change of adding an electron to an atom or ion


 * Oxidation State [or Oxidation Number]**
 * Number assigned to an element [is either positive or negative]
 * *Remember, the oxidation number is written with the sign __before__ the number, unlike ions
 * Based on the following rules:
 * 1) Electrons shared by 2 unlike atoms are counted with the more electronegative atom
 * 2) Electrons shared by 2 like atoms are divided equally between atoms
 * The Oxidation number of an atom is the charge it would have if its bonds were completely ionic. That is, in determining the oxidation number, all the shared electrons are counted with the more electronegative atom.

Example: In K2CrO4, what is the oxidation number for Cr? We know that the oxidation for one K is +1, and there are 2 atoms of K, so: 2(+1)= +2 We also know that the oxidation for one O is -2, and that there are 4 atoms, so: 4(-2)= -8 Therefore, since the sum of the positive atoms and that of the negative must equal 0, X, or Cr, must equal +6.

To learn more, visit: Oxidation-Reduction Reactions

Sources: BROWN, THEODORE L., LEMAY, JR., H. EUGENE, BURSTEN, BRUCE E., and MURPHY, CATHERINE J. __Chemistry, The Central Science__. 10th ed. New Jersey: Pearson Education, Inc. 2006.