1. History of the atom
Dalton
J.J. Thompson
Rutherford
Bohr

2. Dalton’s Atomic Theory
All elements are composed of atoms which are indivisible
Atoms of the same element are identical
Atoms of different elements can mix together in simple whole number ratios to form compounds
Chemical reactions occur when atoms are separated, joined or rearranged. Atoms of one element can’t be changed into atoms of another element
Law of Definite Proportions and Law of Multiple Proportions
Law of Conservation
of Matter

3. Early Atomic Models J.J. Thompson’s plum-pudding model
Rutherford’s Model of the atom

4. Problems with Rutherford Model
Couldn’t account for the chemical properties of the elements (why do elements react in the way that they do?)
Rutherford

5. The Wave Nature of Light
Electromagnetic Radiation
Forms of energy that exhibit wave like characteristics
Wavelength
How far apart the waves are
Frequency
The # of waves to pass a point in a certain timeframe
Wavelength and Frequency are Inversely Proportional

6. Wavelength and Frequency
Not Working very hard
Long Wavelength
Low Frequency

7. Wavelength and Frequency
Working VERY hard
Short Wavelength
High Frequency
What does this tell you about the relationship between wavelength, frequency and energy?

8. The Photo Electric Effect
In the early 1900s an experiment was done that COULD NOT be explained by light being a wave.
Different colors were shined on to a metal plate
Electrons would come off the metal plate for ONLY CERTAIN FREQUENCIES (colors) OF LIGHT
If light only behaved like a “wave” ANY frequency of light would cause the electrons to be released.

9. Brick Wall Analogy Different balls thrown at wall
Each one with different mass
All at same speed of 90 mile/hour
Each ball is like a color or light - each has its own energy
No effect
Dislodges Brick and send it flying
No effect
Super Dense Steel Ball
Ping Pong Ball
Softball JB

10. Einstein and the Photo Electric Effect
This observation led Einstein to believe that light acted like a particle and a wave
This is called the “dual nature” of light
Light carried packets of energy called quanta, or photons

11. Neils Bohr and the New Model of the Atom
Bohr hypothesized that electrons could only be at certain energy “levels” around the nucleus
Electrons could “jump” from lower to higher energy states by absorbing a quantum of energy
When an electron releases the energy it gained, specific colors or wavelengths of light are emitted
Mercury Line Spectra

12. Electrons and the Atom Add Energy e- electron at “Excited State”
Release Energy in the form of colored light
Add Energy e-
electron at “Excited State”
Electrons disappear from one orbit and reappear at another without visiting the space in between!
Electron at “Ground State” e-
Electron at “Ground State” e-
Quantum Leap
Nucleus
Just like a jumper has potential energy at the top of the jump, the electron has stored potential energy in the higher orbit. JB

13. Energy and the Electron
Ground State – lowest energy state for an electron
Excited state – high energy state for e-
Quantum – exact amount of energy to move an electron from one energy level to another

14. Heisenberg Uncertainty Principle
It is impossible to pinpoint the exact location and velocity of an electron at any point in time
You can estimate where an electron will be 90% of the time
An electron cloud shows where an electron spends most of its time

15. Quantum Mechanical Atomic Model
Similar to Bohr model except that e- cannot be found in distinct orbits around the nucleus
Determines how likely it is for an electron to be found in various regions around the nucleus.

16. Atomic Orbitals
Region around the nucleus where an electron of a given energy is likely to be found
Each orbital has a characteristic size, shape, and energy
There are four different orbitals: s, p, d, f
Different types of atomic orbitals

17. Principle Energy Levels
# of Sublevels
Type of sublevel and orbitals
n=1 1
1s (1orbital)
n=2 2
2s (1 orbital)
2p (3 orbitals)
n=3 3
3s (1 orbital)
3p (3 orbitals)
3d (5 orbitals)
n=4 4
4s (1 orbital)
4p (3 orbitals)
4d (5 orbitals)
4f (7 orbitals)
See table 5.1 on page 131
Symbolized by
n = 1,2,3,4 etc.
Each energy level contains sublevels denoted by a number and a letter (ex. 1s)
Each sublevel contains a certain # of orbitals
Every orbital can hold 2 e-

18. Electron Configurations
Aufbau principle: e- occupy the lowest energy sublevels 1st (see pg. 133 figure 5.7)
Pauli Exclusion Principle: every orbital can hold a maximum of 2 e- (paired e- spin in the opposite direction)
Hund’s Rule: e- fill all empty orbitals in a sublevel BEFORE they pair up
Aufbau Diagram

19. Valence Electrons
Valence electrons – electrons found in the highest energy level of an atom
Valence e- determine the chemical properties of the element
A filled outer energy level (stable octet) of 8 e- makes an atom stable (or for He 2e- fills its outer energy level)
Group
Valence e- (and the orbitals they occupy) 1A
1 (s1) 2A
2 (s2) 3A
3 (s2p1) 4A
4 (s2p2) 5A
5 (s2p3) 6A
6 (s2p4) 7A
7 (s2p5) 8A
8 (s2p6)

20. Electron Dot Diagrams Shows only number of valence e-
Electrons shown as dots
Carbon has 4 valence e-