1. 8.7-8.8 Ions, Electron Affinity and Metallic Character

2. Electron Configurations of Ions
When writing an electron configuration for an anion (an ion with a negative charge), electrons are added to the configuration.
Example: the electron configuration of F- is:
1s22s22p6 or [He] 2s22p6
When writing an electron configuration for a cation (an ion with a positive charge), electrons are removed from the electron configuration.
Example: the electron configuration of Mg2+ is:
1s22s22p63s2 or [He] 2s22p6
Transition metal ions are a different case. Experimentation has shown that although 4s is lower in energy than 3d, and electrons will fill the 4s sublevel first, when a transition metal ionizes, the electrons get removed from the 4s sublevel first.

3. Let’s Try a Practice Problem
Write the electron configurations of N3- and Ca2+. N3- 1s22s22p6 or [He] 2s22p6 Ca2+ 1s22s22p63s23p6 or [Ne] 3s23p6

4. Ionic Radii
The radius of a cation is smaller than the radius of the atom from with it was formed. This happens because when an metal atom ionized to form a cation, it loses its valence electrons while maintaining the same number of protons. Those electrons that were lost were being shielded from the nucleus by the core electrons which greatly contributed to its size.
The radius of an anion is larger than the atom from which it was formed. This happens because nonmetals gain electrons to become anions. The extra electrons increase the repulsion among the outer most electrons.

5. Let’s Try a Practice Problem
Choose the larger atom or ion from each pair: a.) K or K+ b.) F or F- c.) Ca2+ or Cl- a.) K b.) F- c.) Cl-

6. Ionization Energy
Ionization energy (IE) of an atom or ion: The amount of energy required to remove an electron in the gaseous state.
This value is always positive, because it always takes energy to remove an electron.
With each electron removed, the IE increase. IE1 < IE2 < IE3 … and after all valence electrons are removed, the energy required to remove the first core electron of the next lowest principle energy level is extremely large in comparison.

7. General Trend in Ionization Energy
Ionization energy decreases down a group (family) on the periodic table. Why?
The outermost principle energy level are increasingly farther away from the positively charged nucleus and therefore less tightly held.
Ionization energy increases across a period on the periodic table. Why?
The electrons in the outermost principle energy level generally experience a greater effective nuclear charge (Zeff).

8. Let’s Try a Practice Problem!
On the basis of periodic trends, determine the element in each pair with the higher first ionization energy (if possible).
a.) Sn or I
b.) Ca or Sr
c.) C or P
d.) F or S
a.) I
b.) Ca
c.) can’t tell
d.) F

9. Electron Affinities
The (EA) of an atom or ion, is the energy change associated with the gaining of an electron by an atom in gaseous state. This energy is usually negative, since gaining an electron is an exothermic reaction.
The coulombic attraction between the nucleus of an atom and an incoming electron usually results in the release of energy as an electron is gained.
General trends:
In general, electron affinity decreases (becomes more positive) down a group.
For the representative elements, as you move across a period, EA becomes more negative (more exothermic).

10. Metallic Character
As we move across a period, metallic character decreases.
As we move down a group (or family), metallic character increases.
On the basis of periodic trends, choose the more metallic element from each pair (if possible).
a.) Ge or Sn
b.) Ga or Sn
c.) P or Bi
d.) B or N
a.) Sn
b.) Can’t tell
c.) Bi
d.) B

11. pg. 376 #’s 65, 69, 73, 79, and 81.
Study for exam on Chapters 7-8