1. Atomic Structure
Atom = the smallest particle of an element that retains the properties of that element

2. Law of Conservation of Mass
Matter cannot be created or destroyed
What happened in Lab 2?

3. Law of Definite Proportions
A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size or source of the sample
H2O is water, all water is H2O, or it’s NOT water
H2O2 is hydrogen peroxide, NOT WATER

4. 2 regions of the atom Nucleus Electron cloud Protons (p+)
Neutrons (n0)
Electron cloud
Electrons (e-)

5. Discovery of Electrons
Cathode ray experiments
Thomson
e- have large charge:mass ratio
Mass of e- was determined by measuring charge and calculating mass based on ratio
Mass of e- = x 10-31

6. Atoms are electrically neutral
Negatively charged e- are balanced by positively charged particles (p+)
Mass of e- is negligible, other particles must account for atomic mass
Rutherford’s model = plum pudding model
Rutherford, Geiger, and Marsden = densely packed, positively-charged nucleus

7. Subatomic Particles
With 1 exception, all atomic nuclei contain p+ and n0
Charge of p+ = charge of e-
n0 has NO charge
Atom is neutral, because #p+ = #e-
Mass of p+ > mass of e-

8. Forces in Nucleus + attracts –
+ repels +, so how are the p+ in the nucleus held together?
When p+ are in close proximity to other p+ and n0, strong forces are created = nuclear forces
But a nucleus can only hold so many protons in proximity and remain stable

9. Atomic Number and Mass Number
Atomic # = # p+
#p+ = #e-
Mass number = #p+ + #n0
Remember: mass of e- is negligible
The atomic number or # of p+ determines the atom’s identity (which element)
The atomic number for every element can be found in the Periodic Table
Element builder

10. Atomic # appears above symbol
Atomic mass appears below symbol
Why is it a decimal?

11. Average Atomic Mass The mass number on a PT is the average atomic mass
Isotopes = atoms of the same element that have different numbers of neutrons and therefore different atomic masses
A specific type of isotope is also called a nuclide
The average atomic mass is the average mass number of all the atoms of a given element including isotopes

12. Nuclides of Hydrogen Protium (99.9885% of H) = 1p+
Deuterium (0.0115% of H) = 1p+ +1n0
Tritium (negligible in nature, can be synthesized) is radioactive = 1p+ +2n0
How many e- do these isotopes have?

13. Most isotopes/nuclides don’t have names…..
Isotopes are referred to by the element name and the mass number (#p+ + #n0)
Ex. Carbon
Carbon-14 (14 = 6p+ + 8n0) = radioactive
Carbon-12 (12 = 6p+ + 6n0) = most common
NOTICE: both types of carbon have same number of protons
What if the atom had 7 protons?
Practice Problems pg 80

14. Atomic mass and mass units
Metric unit for mass =
Too large to measure mass of something so small, so we use amu (atomic mass units) to measure atomic mass
Relative atomic mass = the mass of an atom expressed in relation to a defined standard (carbon-12 atom)
The carbon-12 atom has a relative mass of 12 amu, so 1 amu = 1/12 mass of C-12 atom

15. Calculating average atomic mass
Depends on mass and relative abundance of isotopes
Avg. atomic mass = %A1m1 + %A2m2 + …..

16. The mole
Mole = the amount of a substance that contains as many particles as there are atoms in exactly 12 g of C-12
1 mol of atoms = x 1023 atoms
Just like…
1 dozen eggs = 12 eggs
Avogadro’s number
The mole is a counting unit, just like a dozen
If all the people on Earth worked to count 1 mole of pennies at 1 penny/second, it would take us all 3 million years to count them!

17. Molar Mass
Mole = the amount (mass) of a substance that contains Avogadro’s number (6.022 x 1023) of particles
The mass of 1 mole of a substance = molar mass (found under symbol for each element in PT)
I do practice problem 1 on page 85, you do 2-4 alone. We do practice on p 87. handout practice worksheet solo. Assign homework. Be ready for quiz
Molar mass for neon = g/mol

18. SUMMARY Atomic # = number of p+ Number of p+ = number of e-
Mass number = #p+ + n0
Atomic mass in PT = average atomic mass
Units for avg. atomic mass = g/mol
1 mole = x 1023 particles
e.g.
6.022 x 1023 carbon atoms has a mass of g