1. Development of the Atom

2. Essential Question
How were atoms and their subatomic particles discovered?

3. How small is the atom? https://www.youtube.com/watch?v=yQP4 UJhNn0I
What makes up the nucleus?
What surrounds the nucleus?
What contributes most of the mass of the atom?
What is the atom mostly made up of (volume)?

4. Democritus—400B.C. Greek philosopher World is made of two things
Empty space
Tiny particles, “atoma”
Believed shape determined properties
An idea only
Was NOT backed by experimental evidence

5. Aristotle—350B.C. Greek Philosopher Proposed that matter is continuous
NOT made of smaller particles
Accepted until the 17th century

6. John Dalton—Early 1800s Dalton’s Atomic Theory
All matter is made of atoms, which are indivisible
All atoms of an element are identical
Compounds are formed from two or more elements

7. Dalton’s Atom
Believed to be an indivisible sphere

8. J.J. Thomson—1897 Cathode Ray Tube Experiment
Discovered the existence of the electron

9. Thomson’s Atom Sphere with electrons embedded in it
Known as “Plum Pudding Model”

10. Robert Millikan—1909 Oil Drop Experiment
Charge of an electron = -1.6x10-19Coulombs
Standard Unit = -1

11. Ernest Rutherford—1911 Gold Foil Experiment
discovered the presence of a positive nucleus

12. Rutherford Atom Small dense nucleus, with electrons orbiting
Mostly empty space

13. What are the parts of the atom?
Work with your partner to fill in what you know on the chart below:
Particle
Symbol
Charge
Mass
Location
Proton
Neutron
Electron

14. What are the subatomic particles?
Symbol
Charge
Mass
Location
Proton p+ +1
1 amu
Nucleus
Neutron n0
Electron e- -1
1/1837amu, 0
Surrounds nucleus

15. A really corny joke
A neutron walked into a restaurant, ordered a drink, and asked “How much?”
The waiter replied “For you, no charge.”
Take out your periodic tables!

16. Atomic Number (Z) Number of protons Defines an element
Same # of protons = same element
Different # of protons = different elements
For neutral atoms = # e-
This number is always found on the periodic table!

17. Essential Question
Why do the masses of individual atoms differ?

18. Why do the masses of the same element differ?
Isotopes: atoms of the same element that differ in mass due to different numbers of neutrons
Same # protons (same element), different # neutrons (different isotope)

19. Check Point
Element A has 6 protons, 6 neutrons, and 6 electrons
Element B has 5 protons, 6 neutrons, and 5 electrons
Element C has 6 protons, 8 neutrons, and 6 electrons
Make a statement about the relationship between Elements A, B, and C using the terms “element” and “isotope”

20. Mass Number (A) Defines a particular isotope
Total number of protons & neutrons
Not directly found on the periodic table
Standard notation: element-A
Uranium-235
Uranium-238
To find # neutrons
Mass # = p+ + no
no = Mass # - p+
no = Mass # - atomic #

21. QUICK CHECK!
Determine the number of protons, neutrons and electrons in Nickel-60
Protons=
Neutrons=
Electrons=
What is the term used to describe the 60?

22. Isotopes—Shorthand Notation
Z → identity of element
A → particular isotope
#no = A - Z
Mass # (p+ + no)
Atomic # (p+=e-)

23. U-235 and U-238
Uranium-235
Uranium-238

24. The atomic number of silicon is 14. Therefore:
You Try…
Silicon (Si) has three naturally occurring isotopes: Silicon-28, Silicon-29, and Silicon-30. Write the shorthand notation of each and determine the number of p+, n0, and e- in each.
The atomic number of silicon is 14. Therefore:
28Si has 14p+, 14e- and 14n0 (28-14)
29Si has 14p+, 14e- and 15n0 (29-14)
30Si has 14p+, 14e- and 16n0 (30-14)

25. What if the number of protons and electrons aren’t equal?
Ion-an atom with a charge
Forms by changing number of electrons
Still the same element, only charged
Lose electron
+ ion
Cation
Gain electron
- ion
Anion
Finding # of electrons

26. *Common mistake…
When you lose electrons, the ion becomes positive because there are more protons than electrons
Do not associate a positive charge with adding electrons!
Example: Oxygen (Atomic number 8)

27. Ion Notation
Charge is represented as a superscript after the atomic symbol
Examples:
Na or O2-

28. Another really corny joke
An atom walked into a police station and said he wanted to report a lost electron. The officer asked “Are you sure you lost it?”
He replied “I'm positive.”

29. Practice 2. How many electrons are in the following ions?
A. Ca2+ B. Cl-

30. Essential Question
How is the average atomic mass of an isotope calculated?

31. Units for Masses on the P.T.
amu’s (atomic mass units, referring to mass of 1 atom)
Grams (referring to mass of 1 mole of atoms)
Example (Carbon):
12.01 amu = mass of 1 “average” carbon atom
12.01 g = mass of 1 mole of “average” carbon atoms

32. Decimal #s on the PT
What about the decimal numbers on the periodic table? What do they mean?

33. Masses on the P.T. Average Atomic Mass
the weighted average mass of an element’s isotopes
Weighted Average – depends on % abundance of each isotope
Percent Abundance—% of an element that is a particular isotope
More abundant isotopes = greater contribution

34. Calculating a Weighted Grade
Calculating average atomic mass is like finding a weighted grade.
Joey’s grades are as follows:
85% on tests, 92% on homework, 87% on quizzes, 90% on labs
His class is weighted as follows:
40% on tests, 30 % on labs, 20% on quizzes, 10% on labs
What affects his overall grade the most?
What is Joey’s overall grade in this class?

35. Joey’s Grade Grade=(85)*(0.40)+92*(0.30)+87*(0.20)+ 90*(0.10)

36. Calculating Avg. Atomic Mass
Convert all/any percents to decimal form
Multiply each mass by its percent abundance
Add together for avg. atomic mass

37. How does this number compare to the atomic mass on the periodic table?
The natural abundance for boron isotopes is: 19.9% 10B ( amu) and 80.1% 11B (11.009amu). Calculate the atomic weight of boron. 
Average Atomic Mass =
0.199 (10.013amu) +
0.801 (11.009amu)
How does this number compare to the atomic mass on the periodic table?

38. You Try…
Naturally occurring copper consists of % copper-63 ( amu) and % copper-65 ( amu). Find the average atomic mass of copper.

39. Corny Joke  What did the ion say to the other ion?
“I’ve got my ion you!”

40. Finding Relative Abundances
Work backwards to determine the relative abundances of isotopes.
You must use a system of equations to do this.

41. Relative Abundances
Gallium consists of two naturally occuring isotopes with masses of and amu. The average atomic mass of Ga is amu. Calculate the abundance of each isotope.
69.72amu=68.926x y
1.0 = x + y

42. Relative Abundances 69.72amu=68.926(1-y) + 70.925y
Isotope = 39.72%
Isotope = 60.28%

43. You Try….
Phosphorus has two isotopes, 32P (mass = 32.00amu) and 30P(mass = 30.00amu). The experimentally determined mass of phosphorus is amu. What are the percentage abundances of the two isotopes?
30.973= 32.00x (1-x)
= 32.00x – 30.00x
0.973 = 2.00x x=
32P= 48.65% 30P=51.35%