1. The Electromagnetic Spectrum and the Model of the Atom Part I
Chemistry – Mrs. Cameron

2. The Purpose of Science
The purpose of science to make models that explain natural phenomena.
A model is the best possible explanation which accounts for all observed phenomenon and has predictability.
1) An unanswered question means change the model.
2) Predictability is the test of a good or “true” model.

3. Radioactivity and Light have been tools to discover the structure of atoms.
Do you Remember?
Thompson’s Cathode Ray Experiment
Rutherford’s Gold foil Experiment

5. Thomson’s Plum-Pudding Model

6. Ernest Rutherford’s Gold Foil Experiment

7. Results of the Rutherford experiment
(a) The results that the metal foil
experiment would have yielded if the
plum pudding model had been correct
(b) Actual results

8. Rutherford’s Nuclear Model of the atom
Small, dense, positively charged nucleus
Contains protons (+1 charge)
Contains neutrons (no charge)
Remainder of the atom is mostly empty space
Contains electrons (-1 charge) in the “empty space”

9. New Evidence: Continuous and Line Spectra
Let’s Talk about LIGHT
White light
is actually made up of many colors
When white light is passed through a prism, the light refracts, or bends to display all of its component colors

10. The Wave Nature of Light
Crest 
Crest – the top of a wave
Trough – the bottom of the wave
Origin – the center line through
which the wave oscillates
Origin 
←trough
Amplitude – distance from the origin to the crest or the origin to the trough of a wave. (Indicates intensity)
Frequency () – the number of waves that pass a given point in a given amount of time
Wavelength (λ) - distance from crest to crest or trough to trough
Speed (Velocity) – the amount of distance covered in a specified amount of time.

11. The Wave Nature of Light

12. Wave Behaviors Reflection – a wave strikes an object and bounces off
Refraction – the bending of waves
Diffraction – the bending of waves around an opening or around the edge of an object
Interference – the ability of two or more waves to add together forming regions of large or small amplitude.

13. Wave Interference patterns
Constructive interference: waves “in-phase” create waves of greater amplitude ( they add)
Destructive interference: waves “out-of-phase” create waves of lower amplitude (they cancel out)

14. The Wave Nature of Light – The Electromagnetic Spectrum
 Ionizing Radiation 
 Nonionizing Radiation 

15. Units for Wave Characteristics
Wavelength (λ “lambda”) = meters with a Greek prefix (nanometers , nm)
Frequency ( “nu”) –cycles per second or Hertz (1/s, sec-1, Hz)
Speed – meters per second (m/s)
Wave Calculations
All electromagnetic radiation travels at the same speed
through the vacuum of space: 3.00 x 108 m/s (c)
Wavelength and frequency are inversely proportional: λ = c/
Energy and frequency are directly proportional: E = h
E = energy in joules,
h = Plank’s constant  J·s

16. The Work of Max Planck (1858-1947)
Able to predict the wavelengths of light changes with temperature
Energy must be emitted in a “Quantized” way, or restricted to certain quantities.
“Quantum” (singular) or “Quanta” (Plural)
A quantum is a packet of energy
VERY SMALL
Related through Planck’s Constant
h=  J·s, E = h

17. The Photoelectric Effect
When light is shined on a piece of metal, electrons are ejected from the metal
Only light containing enough energy (of a certain wavelength and frequency) works

18. Einstein and the Photoelectric Effect
Albert Einstein ( ) proposes that light has a particle nature too.
He relates Planck’s idea of “quantized energy” to light.
Light energy quanta = photons
Photons transfer energy to electrons when light strikes the metal.
Photons must be of sufficient energy for this to occur.

19. What do you see?
Depending on how you look at this it can be an old lady or a young lady turning her head.
The picture has a dual nature

20. Line Spectra of Elements
A line spectrum is a spectrum that contains only certain colors, or wavelengths of light.
The rainbow is a continuous spectrum
Elements emit line spectra when they are vaporized in an intense flame or with electricity.

21. Continuous spectrum of white light
Line spectra of elements

22. Line Spectra and the Quantization of Energy

23. Neils Bohr (1885-1962) Planetary Model of the Atom
Combines Rutherford’s nuclear model and Planck’s quanta of energy.
Electrons must have certain energy levels in which they travel. (Energy Level Postulate)
Energy of Electrons must be quantized to explain line spectra.
(Transitions Between Energy Levels)

24. Bohr Model Ground state = energy level closest to the nucleus
Excited state = energy levels farther away from the nucleus
Energy levels represented by quantum numbers ex. n = 1, n= 2, n = 3 etc.

26. Emission Spectrum of Hydrogen
In general, the line spectrum of an element is
rather complicated
The line spectrum of hydrogen, with a single electron,
is the simplest

27. The energy of the electron is said to be quantized
The Evidence of Line Sprectra:
Atomic line spectra tell us that when an excited atom loses energy, not just any arbitrary amount can be lost
This is possible if the electron is restricted to certain energy levels
The energy of the electron is said to be quantized

28. The energy emitted or absorbed =  energy levels
Radiation absorbed – electron moves from the ground state to an excited state.
Can’t maintain this higher energy level
Electron falls back down to the ground state
Radiation is emitted as it returns to the ground state
The energy emitted or absorbed =  energy levels

30. Bohr’s Calculations
Bohr proposed that the electrons moved around the nucleus is fixed paths or orbits much like the planets move around the sun
The orbits, labeled with the integer n, have energy
This equation allows the calculation of the energy of any orbit
Bohr was able to use this model to calculate the energies of the light given off by the hydrogen.
Unfortunately, the model became too mathematically complicated for any element with more electrons than hydrogen.
Also, electrons not completely explained as particles

31. Limitations of the Bohr Model
•It cannot explain the spectra of atoms other than hydrogen.
•Electrons do not move about the nucleus in circular orbits.
•However - the model introduces two important ideas:
•The energy of an electron is quantized: electrons exist only in certain energy levels described by quantum numbers.
•Energy gain or loss is involved in moving an electron from one energy level to another.

32. Matter Waves Louis DeBroglie (1892-1897)
Electrons (matter) have a dual nature just like light.
Proved it mathematically
Evidence: diffraction patterns produced by beams of electrons