1. Unit 3 – Atomic Structure
Bravo – 15,000 kilotons

2. Democritus
400 BC
Greek philosopher
1st to come up with idea of atoms

3. John Dalton – 1800’s Major contributor of Atomic Theory
All matter made of atoms
All atoms of an element are alike
Atoms cannot be created or destroyed
Atoms combine in whole-number ratios to form compounds

4. JJ Thomson – late 1800’s
Cathode Ray Experiment – discovery of electrons
“Plum Pudding” model of atom
Measured charge to mass ratio of e-

5. Thomson’s Atomic Model
Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

7. Conclusions from the Study of the Electron
Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.
Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons
Electrons have so little mass that atoms must contain other particles that account for most of the mass

8. Millikan - 1910 Oil Drop Experiment
Determined actual charge and mass of an e-

9. Rutherford - 1910 Discovered nucleus and that it was positive
Gold Foil Experiment
1) Most of atom is empty space (majority of particles went straight through)
2) nucleus is small, dense and positively charged (some positive charges were greatly deflected)

10. Rutherford’s Gold Foil Experiment
Alpha particles are helium nuclei
Particles were fired at a thin sheet of gold foil
Particle hits on the detecting screen (film) are recorded

11. Rutherford’s Findings
Most of the particles passed right through
A few particles were deflected
VERY FEW were greatly deflected
“Like howitzer shells bouncing off of tissue paper!”
Conclusions:
The nucleus is small
The nucleus is dense
The nucleus is positively charged

12. Niels Bohr - 1915 Proposed early model of atom
“Planetary Model” electrons orbit nucleus like planets orbit sun
Lacks math of modern version
Has some errors/violates current theory
Radiation is emitted when electrons move from one orbit to another

13. Chadwick
Discovered neutrons

14. Modern Atomic Theory (changes from Dalton)
Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions!
Atoms of an element have a characteristic average mass which is unique to that element (isotopes)

15. Foundations of Atomic Theory
Law of Conservation of Mass: mass is neither created or destroyed in ordinary chemical reactions
Law of Definite Proportions: compounds contain same elements in same ratio by mass
Example: NaCl is always 39.9% Na and 60.66% Cl by mass
Law of Multiple Proportions: 2 or more different compounds composed of same two elements have ratios of small whole numbers
Example: CO vs CO2 ratio of oxygen to oxygen is 2 to 1

17. What is AMU?
Stands for atomic mass unit – used when describing “relative” atomic masses
This system is used because the actual masses of atoms are so small
Carbon-12 is the standard to which all other elements are compared (i.e. hydrogen-1 has a mass that is 1/12 that of carbon-12 so it’s mass would be 1 amu)

18. Atomic Number
Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element.
Element
# of protons
Atomic # (Z)
Carbon 6
Phosphorus 15
Gold 79

19. Mass Number Nuclide p+ n0 e- Mass #
Mass number is the number of protons and neutrons in the nucleus of an isotope.
Mass # = p+ + n0
Nuclide p+ n0 e-
Mass #
Oxygen - 10 - 33 42
- 31 15 18 8 8 18
Arsenic 75 33 75
Phosphorus 16 15 31

20. Composition of the nucleus
Atomic Masses
Atomic mass is the average of all the naturally isotopes of that element.
Carbon =
Isotope
Symbol
Composition of the nucleus
% in nature
Carbon-12
12C
6 protons
6 neutrons
98.89%
Carbon-13
13C
7 neutrons
1.11%
Carbon-14
14C
8 neutrons
<0.01%

21. Isotopes
Isotopes are atoms of the same element having different masses due to varying numbers of neutrons.
Isotope
Protons
Electrons
Neutrons
Nucleus
Hydrogen–1
(protium) 1
Hydrogen-2
(deuterium)
Hydrogen-3
(tritium) 2

22. ISOTOPES Example: Carbon, C, exists in 3 isotopes: Isotope Protons
Neutrons
Mass
Carbon-12 6 12
Carbon-13 7 13
Carbon-14 8 14

23. Isotope symbols Hyphen notation Nuclear notation
Carbon – C 6p+ and 6no
Carbon – C 6p+ and 7no
Carbon – C 6p+ and 8no 6 6 6

24. How to Calculate the Average Mass
What is the average atomic mass of sample of Cesium with 3 isotopes:
75% 133Cs, 20% 132Cs, and 5% 134Cs.
0.75 x 133 = 99.75
0.20 x 132 =
0.05 x 134 =
Total =
avg. atomic mass

25. The Mole 1 dozen = 12 1 gross = 144 1 ream = 500 1 mole = 6.02 x 1023
There are exactly 12 grams of carbon-12 in one mole of carbon-12.

26. I didn’t discover it. Its just named after me!
Avogadro’s Number
6.02 x 1023 is called “Avogadro’s Number” in honor of the Italian chemist Amadeo Avogadro ( ).
I didn’t discover it. Its just named after me!
Amadeo Avogadro

27. Calculations with Moles: Converting moles to grams
How many grams of lithium are in 3.50 moles of
lithium?
3.50 mol Li
6.94 g Li
= g Li
45.1
1 mol Li

28. Calculations with Moles: Converting grams to moles
How many moles of lithium are in 18.2 grams of
lithium?
18.2 g Li
1 mol Li
= mol Li
2.62
6.94 g Li

29. Calculations with Moles: Using Avogadro’s Number
How many atoms of lithium are in 3.50 moles of
lithium?
3.50 mol Li
6.022 x 1023 atoms Li
= atoms Li
2.11 x 1024
1 mol Li

30. Calculations with Moles: Using Avogadro’s Number
How many atoms of lithium are in 18.2 g of
lithium?
18.2 g Li
1 mol Li
6.022 x 1023 atoms Li
6.94 g Li
1 mol Li
(18.2)(6.022 x 1023)/6.94
= atoms Li
1.58 x 1024