1. The Atom
Chapters 4-5

2. Atomic Theories Democritus ~ 400 BC Dalton ~ 1800’s
believed that atoms were indivisible and indestructible
Dalton ~ 1800’s
Developed through experiments
First Atomic Model

3. Dalton’s Atomic Model
All elements are composed of tiny indivisible particles called atoms
Atoms of the same element are identical. Different elements have different atoms.

4. Discovery of Electron
1897 – JJ Thomson, using cathode ray tube, discovered negatively charged particles called electrons
1909 – Robert Millikan - Oil Drop Experiment
Determined charge on an electron.

5. “Plum Pudding” Model
Uniform positive sphere with negatively charged electrons embedded within.

6. Radiation Late 1800’s – discovery of radiation Three Types Alpha Beta
Gamma

7. Rutherford Gold Foil Experiment - 1909
Shot alpha particles at gold foil
Most went through foil with little or no deflection.
Some were deflected at large angle and some straight back.

8. Rutherford Gold Foil Experiment - 1909

9. Rutherford Model Conclusions from Gold Foil Experiment
Atom is Mostly Empty Space
Dense positive nucleus
Electrons moving randomly around nucleus

11. The Atom An Atom is made of 3 Subatomic Particles Electrons Protons
Neutrons

12. Subatomic Particles Electron Discovered in 1897 by JJ Thomson
Negative charge (-1)
Mass = *10-28g
Approx mass ~ 0
Found outside of nucleus

13. Subatomic Particles Proton Discovered in 1919 by Rutherford
Positive charge (+1)
Mass = *10-24g
Approx mass ~ 1 atomic mass unit (u)
Found inside nucleus

14. Subatomic Particles Neutron Discovered in 1932 by James Chadwick
No charge (0)
Mass = *10-24g
Approx mass ~ 1 atomic mass unit (u)
Just slightly larger than a proton
Found inside nucleus

15. Atomic Structure Atoms have no net charge
# of electrons = # of protons
# of electrons around nucleus = # of protons in nucleus

16. Atomic Structure Atomic Number Mass Number
Number of protons in an element
All atoms of the same element have the same number of protons
Mass Number
Number of protons and neutrons in an atom

17. Atomic Structure # of Neutrons = Mass Number – Atomic Number
Atoms of the same elements can have different numbers of neutrons
Isotope – atoms of the same element with different number of neutrons

18. Ion
Atom or group of atoms that have gained or lost one or more electrons
Have a charge
Example:
H+, Ca2+, Cl-, OH-

19. Chemical Symbols Cl-35 Chlorine-35 Mass Number Charge Atomic Number
Number of Atoms

21. Atomic Theories
Rutherford’s model could not explain the chemical properties of elements
Niels Bohr believed Rutherford’s model needed to be improved
Bohr proposed that electrons are found only in circular paths around the nucleus

22. Bohr Model Dense positive nucleus
Electrons in specified circular paths, called energy levels
These energy levels gave results in agreement with experiments for the hydrogen atom.

23. Bohr Model

24. Bohr Model
Each energy level can only hold up to a certain number of electrons
Level 1  2 electrons
Level 2  8 electrons
Level 3  18 electrons
Level 4  32 electrons

25. Electron Configuration
The way in which electrons are arranged in the atom
Example: Na
Valence Electrons
Electrons in the outermost energy level

26. Wave Mechanical Model More detailed view of the Bohr Model
Schrödinger Wave Equation and Heisenberg Uncertainty provides region of high probability where electron COULD be.
Orbital
Modern Model
AKA Quantum Mechanical Model, Electron Cloud Model

27. Wave Mechanical Model Orbital
Regions of space where there is a high probability of finding an electron

29. M&M’s Demo What colors are found in a regular M&M’s bag? Green Yellow
Orange
Blue
Red
Brown

30. M&M’s Demo Do you get an equal amount of each color in each bag?
If we opened up all the regular M&M bags in the world would we get an equal number of each color?
Are you supposed to?

31. M&M’s Demo Color 1 bag World Blue % 24% Green 16% Yellow 14% Orange
20%
Red
13%
Brown

32. M&M’s Demo M&M’s come in certain abundances (percentages)
So do isotopes of each element
Relative Abundance
Percent of each naturally occurring isotope found in nature

33. Average Atomic Mass Atomic Mass Example
Weighted average based on the relative abundance and mass number for all naturally occurring isotopes
Example
C % u
C %

34. Atomic Mass
C %
C %
Carbon = 0.989* *13 = u