1. Subatomic particles

2. Average Atomic Mass
All atoms of an element are not identical – some have different #’s of neutrons. This makes different forms of the element called isotopes
11B 5p+ 6no 10B 5p+ 5no
Both are boron b/c both have 5 proton. (Same atomic number)

3. Isotopes
Isotopes have slightly different masses b/c of the # of neutrons. We use average atomic mass.
Units for avg. atomic mass are amu. amu = atomic mass units
To find avg. atomic mass you must take relative amounts into consideration. (Use isotopes)

4. Average atomic mass  (mass isotope) (% isotope)  = sigma (summation)
% isotope must = 1
If given % abundance, you will need to divide the numbers by 100.

5. Strontium and Chlorine examples

6. Counting Subatomic Particles
Atomic number 
Average Atomic Mass 
X Charge X = element symbol 6 C
12.011 A Z

7. Counting Subatomic Particles
Z = Atomic number (from Periodic Table)
# of p+ = # of e- in a neutral atom
A = Atomic Mass Number (from the nucleus) it is not the Average Atomic Mass
  # of p+ + # of n0
Charge = the number of electrons lost or gained.

8. Isotopes: Same element but with a different number of neutrons
Isotopes: Same element but with a different number of neutrons. The protons must be the same!!!
Examples:
1) Lead – ) Lead – 207

9. Electrons are negative!!!
If you have a negative charge you gain electrons
more e- than p+
If you have a positive charge you lost electrons
more p+ than e-

10. Ions have charges
Let’s compare O 2- and Na 1+

11. Ions have charges O 2- Na 1+ Gained 2 e- Anion (Angry – negative)
Lost 1 e-
Cation
(Happy cat – positive)

12. 3) Aluminum – 27, 3+ charge 4) Sulfur – 33, 2- charge