1. Electromagnetic Radiation and Light

2. I. Models of the Atom Many different models:
Dalton-billiard ball model (1803)
Thompson – plum-pudding model (1897)
Rutherford – Nuclear model of the atom (1911)
Bohr – uses quantized energy of the atom (1913)
Quantum Mechanical Model of the Atom (1926)

3. Each new model contributed to the model we use today.
Even our current model, does not give us an exact model of how electrons behave.

4. A. The Bohr Model
Bohr used the simplest element, hydrogen, for his model
Proposed electron is found in specific circular paths, or orbits around the nucleus
Each electron orbit was
thought to have a fixed
energy level.
Lowest level-ground state
Any Higher Level-
excited state

5. The Bohr Model cont.
One electron is capable of many different excited states (e- jumping to higher level)
Quantum: specific amount of energy an e- can gain or lose when moving energy levels
You can excite an e- with energy like electricity, the sun, or magnets
Electron dropping from higher level to lower-releases energy
energy

6. B. Problems with the Bohr Model
OOPS!- Model only works with hydrogen
Did not account for the chemical behavior of atoms
WRONG: Electrons do not move around the nucleus in circular orbits
Still very helpful!!

8. II. How do Neon Signs work?
They have “excited” gases in them.

9. Explanation
Step 2
Step 1: an electron absorbs energy and moves to a higher energy level
Step 2: e- drops back down to a lower energy level
During drop it gives off energy called a “photon”
Sometimes this energy is visible light (ROYGBIV)
Step 1
When a photon is emitted, energy is released. We can calculate the energy released using the equation: E = h  ν

10. Application: Atomic Emission Spectrum
Used to determine which elements are present in a sample
Used to determine which elements are present in a star (because stars are gases)
Each element has a unique spectrum
Only certain colors are emitted because the energy released relates to specific frequency

11. Spectroscope
A spectroscope is needed to see the atomic emission spectra, which acts similar to a prism, separating different frequencies of light

12. Electromagnetic Spectrum
Electromagnetic spectrum is the range of all energies emitted from photons acting like waves.

13. Electromagnetic Spectrum with Visible Light Spectrum

14. Lab: Atomic Emission Spectra of Several Gases

15. Light
Behaves like a particle
Behaves like a wave

16. Characteristics of a Wave
Wavelength  (lambda) – shortest distance between equivalent points on a continuous wave [Unit = meters]
Frequency  (nu) – the number of waves that pass a given point per second [Unit = 1/second = s-1 = Hertz (Hz)]
Crest – Highest point of a wave
Trough – Lowest point of a wave
Amplitude (a)– height from its origin to its crest (highest point) or trough (lowest point) [Unit = meters]
Amplitude
(Wavelength)

17. Wavelength and Frequency
Wavelength () and frequency () are related
As wavelength goes up, frequency goes down
As wavelength goes down, frequency goes up
This relationship is inversely proportional

18. Wavelength and Frequency cont.
Speed of light (c) = 3 x 108 m/s
c = 
= c / 
 = c / 
Speed of light  
wavelength
frequency

19. Question Time
Calculate the wavelength () of yellow light if its frequency () is 5.10 x 1014 Hz.   c

20. Question Time
What is the frequency () of radiation with a wavelength () of 5.00 x 10-8 m? What region of the electromagnetic spectrum is this radiation?   c

21. How Much Energy Does a Wave Have?
Planck’s constant
frequency
Energy
Energy of a wave can be calculated
Use the formula E= h
E= Energy
 = frequency
h = Planck’s constant = x Joule . Sec
Joule is a unit for energy (J)
Energy and frequency are directly proportional, as frequency increases, energy increases

22. Question Time
Remember that energy of a photon given off by an electron is E =h
How much energy does a wave have with a frequency of 2.0 x 108 Hz? ( h = x J.s)
E = 1.3 x Joule

23. Visible Light, Frequency, and Energy
Red: longest wavelength (),
smallest frequency ()
Red: frequency smallest (),
least amount of energy (E)
Violet: smallest wavelength (),
largest frequency ()
Violet: frequency largest (),
greatest amount of energy (E)

24. Flame Test
The flame test is a way to determine the element present in a sample
When placed in a flame, each element gives off a different color
Operates same as neon signs; electrons excited by heat and fall back down and give off different colors.