1. 7. Describe the structure of a typical atom.
Identify where each subatomic particle is
located.
A typical atom consists of a central, small, dense nucleus containing protons and neutrons. The nucleus is surrounded by a cloud of negatively charged electrons.

2. Direct vs. Indirect Evidence
Direct Evidence: You can use your senses to know what is happening. You can see it!
Proves the existence of a particular fact without any assumptions.
Indirect Evidence: You can not see what is happening. You must use reasoning powers to make conclusions. [Circumstantial evidence]

3. 9. Evaluate the experiments that led to the
conclusion that electrons are negatively charged
particles found in all matter.
The deflection toward positively charged plates demonstrated the negatively charged nature of electrons; the fact that changing the type of electrode or the type of gas used in the cathode-ray tube did not affect the ray -led to the conclusion that electrons are present in all matter.

4. Thomson’s Experiment Voltage source + -
Passing an electric current makes a beam appear to move from the negative to the positive end
By adding an electric field he found that the moving pieces were negative

5. 8. Compare and contrast Thomson’s plum pudding
atomic model with Rutherford’s nuclear atomic
model.
Thomson’s plum pudding model describes atoms as spherical particles with uniformly distributed positive charge in which individual, negatively charged electrons are located in fixed positions. In contrast, Rutherford’s model states that an atom is mostly empty space, with a small, dense, central nucleus containing all of an atoms positive charge and most of its mass. The negatively charged electrons move through the empty space and are held in the atom by their attraction to the positively charge nucleus.

6. Thomson’s Model Found the electron
Couldn’t find positive (for a while)
Said the atom was like plum pudding
A bunch of positive stuff, with the electrons able to be removed

7. Rutherford’s Gold Foil experiment
Ernest Rutherford English physicist. (1910)
Believed in the plum pudding model of the atom.
Wanted to see how big atoms are
Used radioactivity
Alpha particles - positively charged pieces given off by uranium
Shot them at gold foil which can be made a few atoms thick

8. Florescent
Screen
Lead block
Uranium
Gold Foil

9. What he expected:
He expected the alpha particles to pass through without changing direction very much
Because he thought the positive charges were spread out evenly. Alone they were not enough to stop the alpha particles

10. What he got:

11. How he explained it + Atom is mostly empty space Small dense, positive
center (nucleus)
Alpha particles are
deflected by it if
they get close enough +

12. +

13. Modern View The atom is mostly empty space Two regions:
Nucleus- protons and neutrons
Electron cloud- region where you might find an electron

14. 10. Compare the relative charge and mass of each of the subatomic particles
Relative Mass
Electron -1
1/ (~0)
Proton +1 1
Neutron

15. - by the atomic number ( # of protons) The # of protons
20. Explain how the type of an atom is defined
- by the atomic number
( # of protons)
21. Recall Which subatomic particle identifies an atom as that of a particular element?
The # of protons

16. Average Mass Activity   
Find the average mass of all the spheres in the tray.
You may:
-Make no more than 3 mass measurements.
-Mass only 1 sphere at a time
Record all measurements and show your work.
Explain what you did and all assumptions.

17. 22. Explain how the existence of isotopes is related to the fact that atomic masses are not whole numbers
Atomic masses are not whole numbers because they represent weighted averages of the masses of the isotopes of an element.

18. Isotopic Notation: Symbolic Written Form p + n gain/loss of e-
Isotopes – atoms of the same element (same atomic #),
but with different # neutrons (so different mass #).
Isotopic Notation:
Symbolic
Written Form
p + n
gain/loss of e-
atomic mass charge
Symbol
Atomic number
Element – atomic mass p
___ n

19. #n = (mass # - atomic #) = 14 - 6 = 8
Example:
Find the number of protons, neutrons, and electrons 14 C 6
p + n
Carbon-14 p
# p = atomic # = 6
#n = (mass # - atomic #) = = 8
# e- = # p = atomic # = 6

20. Practice Problem
 Determine the number of protons, neutrons,
& electrons in each of the isotopes given:
O Br U Cu
8 p p p p+
8 n 45 n 143 n 35 n
10 e e e e-

21. 23. Calculate Copper has two isotopes: Cu-63 (abundance=69.2%, mass= amu) and Cu-65 (abundance= 30.8%, mass= amu) Calculate the atomic mass of copper.
.692 x = 43.5
.308 x = 20.0
63.5 amu

22. 24. Three magnesium isotopes have atomic masses and relative abundances of amu (78.99%), amu (10.00%), and (11.01%). Calculate the atomic mass of magnesium
x = 18.95
x =
x =
24.31 amu

23. Silver is found in two isotopes with atomic
Practice Problem
 Silver is found in two isotopes with atomic
masses amu and amu,
respectively. The first isotope represents
51.82% and the second 48.18%. Determine
the average atomic mass of silver.
(0.5182) (0.4818) = amu

24. Do Problems pgs
36, 39, 44, 46, 58-62, 64 & 65