1. From Rutherford to Schrodinger
Discovery of the Atom
From Rutherford to Schrodinger

2. Evolution of the atomic model

3. Rutherford Model
Derived from the gold foil experiment

4. Gold Foil Results

5. Milikan’s Oil Drop Experiment
Mass of oil drop determined by the rate of fall in absence of electric field
Charge adjusted to exactly hold oil drop in place.
Charge always multiple of -1.60e-19 Coulombs
Use e/m ratio, mass of electron x g.

6. Waves Low frequency High frequency long wavelength l
Amplitude
Low
frequency
short wavelength l
Amplitude
High
frequency 6

7. Wave characteristics of light
Frequency and wavelength are inversely related
c = λν

8. Electromagnetic Spectrum

9. Electromagnetic Spectrum

10. Visible Spectrum of Light
Waves 1/33,000” long
Waves 1/70,000” long
Red
Orange
Yellow
Green
Blue
Indigo
Violet
PRISM
Slit
Ray of
White Light
All light is bent (refracted) passing through a prism; violet is bent most and red least. A beam of sunlight produces a continuous band of rainbow colors showing that light is a mixture of colors. 10

11. Double Slit Experiment

12. Emission Spectra for Barium
1870- Kirchoff & Einstein
Each element emits a different line spectrum
Can be used to identify an element
Elements emit only certain wavelengths of light which correspond to particular color
“Fingerprint” of element
Emission Spectra for Barium

13. End of 1900 century: Believed that matter and energy were distinct
Matter = particles
Energy = light in the form of a wave
Max Planck (1900)
Found that matter could not absorb or emit any quantity of energy
Energy is emitted in small, specific amounts (quanta)
Quantum = fixed amount
Max Planck
Why are different ranges of wavelengths emitted at different temperatures?
Why do different elements exhibit different colors when heated, or characteristic colors for each element?
In 1900, Max Planck explained the “ultraviolet catastrophe” by assuming that the energy of electromagnetic waves is quantized rather than continuous—energy could be gained or lost only in integral multiples of some smallest unit of energy, a quantum.
• Classical physics had assumed that energy increased or decreased in a smooth, continuous manner.
• Planck postulated that the energy of a particular quantum of radiant energy could be described by the equation E = h, where h is the Planck’s constant and is equal to x joule•second (J•s).
• As the frequency of electromagnetic radiation increases, the magnitude of the associated quantum of radiant energy increases.
Courtesy Christy Johannesson 13

14. Continuous vs. Quantized
A B
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 330 14

15. Quantum Theory Albert Einstein (1905)
Used Planck’s theory to explain the photoelectric effect
Electrons are ejected from the surface of a metal when light shines on the metal
Albert Einstein
Planck’s quantization hypothesis was used to explain a second phenomenon that conflicted with classical physics.
• When certain metals are exposed to light, electrons are ejected from their surface.
– Classical physics predicted that the number of electrons emitted and their kinetic energy should depend only on the intensity of light, not on its frequency.
– However, each metal was found to have a characteristic threshold frequency of light — below that frequency, no electrons are emitted regardless of the light’s intensity, above the threshold frequency, the number of electrons emitted was found to be proportional to the intensity of light and their kinetic energy proportional to its frequency, a phenomenon called the photoelectric effect.
“The free, unhampered exchange of ideas and scientific conclusions is necessary for the sound development of science, as it is in all spheres of cultural life. ...
We must not conceal from ourselves that no improvement in the present depressing situation is possible without a severe struggle; for the handful of those who are really
determined to do something is minute in comparison with the mass of the lukewarm and the misguided. ... Humanity is going to need a substantially new way of thinking if it is to survive!" (Albert Einstein)
Courtesy Christy Johannesson 15

16. Photoelectric Effect

17. Photoelectric Effect No electrons are emitted Electrons are emitted
Bright
red light
infrared rays
Dim
blue light
ultraviolet rays or or
When a photon strikes the surface of a metal, it transfers its energy to an electron in a metal atom. The electron either “Swallows” the entire photon or does not swallow any of it. If the energy is too small, it will not be able to escape. The energy (frequency) matters, not the number of photons (intensity).
X-rays have high frequencies and high energy (can damage living organisms). Radio waves have low frequencies and low energy and pose no hazard.
Metal plate
Metal plate

18. Photoelectric effect
Light with a frequency below the threshold frequency will not cause the photoelectric effect
Light with a frequency higher than the threshold frequency is linearly related with the energy of the emitted electron
All metals had the same results, but had different threshold frequencies
E= hv

19. Quantum Theory
The energy of a photon is proportional to its frequency.
E = h
We are unaware of quantum effects in the world around us because quanta of energy are extremely small (hence the small Planck’s constant).
E: energy (J, joules)
h: Planck’s constant (  J·s)
: frequency (Hz)
Courtesy Christy Johannesson

20. Implications of Photoelectric Effect
Light exists as a quantized, or discrete, packet of energy called a photon
Amount of energy dependent on the frequency of light.
Therefore light has a dual nature. It behaves both as a wave and as a particle

21. Dual Nature of Light Light has properties of both waves and particles
Wave- Light has a wavelength and frequency
Einstein‘s photons of light were individual packets of energy that had many characteristics of particles.
• Einstein’s hypothesis that energy is concentrated in localized bundles was in sharp contrast to the classical notion that energy is spread out uniformly in a wave.
Particle- Light carries a packet of energy called a photon
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 325 21

22. Waves Low frequency High frequency long wavelength l
Amplitude
Low
frequency
60 photons
162 photons
low energy
short wavelength l
Amplitude
Einstein‘s photons of light were individual packets of energy that had many characteristics of particles.
• Einstein’s hypothesis that energy is concentrated in localized bundles was in sharp contrast to the classical notion that energy is spread out uniformly in a wave.
High
frequency
high energy 22

23. Electromagnetic Spectrum

24. The New Model of the Atom
Must explain the photoelectric effect and line spectra
1915 Niels Bohr – decided to start with hydrogen
1 electron in the atom
Line Spectrum
for Hydrogen

25. Bohr model of the atom

26. Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels

28. Bohr Model of Hydrogen 5 4 3 2 1 nucleus
Courtesy Christy Johannesson

29. Bohr model of the atom
How does the Bohr model explain the photoelectric effect and spectral lines?
Bohr Model

30. An Excited Lithium Electron
Excited Li electron
Energy
Photon of
red light
emitted
Lithium electron
lower energy state
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 326

31. Excitation of Hydrogen Atoms
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 328

32. Return to Ground State

33. How does the Bohr model fit in with the Periodic Table?

34. Problems with the Bohr Model
The energies of the electrons calculated by Bohr only worked for one electron systems

35. Problems with the Bohr Model
DeBroglie
λ = h mv
All objects with mass and in motion travel as a wave
Therefore, electrons cannot move around the nucleus in a fixed circular path

36. Problems with the Bohr Model
Heisenberg Uncertainty Principle
Impossible to know both the velocity and position of an electron at the same time
Werner Heisenberg
~1926 g
Microscope
Werner Heisenberg ( )
The uncertainty principle: a free electron moves into the focus of a hypothetical microscope and is struck by a photon of light;
the photon transfers momentum to the electron. The reflected photon is seen in the microscope, but the electron has moved
out of focus. The electron is not where it appears to be.
Electron 36

37. Problems with the Bohr Model
Heisenberg Uncertainty Principle
Therefore, the electron cannot be moving around the nucleus in a fixed circular path at a fixed distance from the nucleus.
A new model of the atom was proposed

38. 1926- Schrodinger Energy in electrons is quantized
Electrons move in a wave-like behavior
Uncertainty Principle
Believed that the atom contains energy levels
Needed to explain line spectra of elements

39. New Model: Quantum Mechanical Model Based on Probability

40. The Schrodinger model of the atom
Based on probability
Electron Clouds

41. A closer look at s and p orbitals

42. D orbitals

43. Results of Schroedinger Calculations

44. The Schrodinger Model
Energy Diagrams

45. Many Electron System

46. Maximum number of electrons
Orbital Type
Number of orbitals
Maximum number of electrons s 1 2 p 3 6 d 5 10 f 7 14

47. Many Electron System

49. Quantum Numbers n l (n-1 and 0) Orbital Designation ml (-l to l)
# of Orbitals 1 2 3 4 1s 2s 2p 3s 3p 3d
-1. 0, 1
-1, 0, 1
-2, -1, 0, 1, 2
1 **
3 ** 5

50. Atomic Radius
Atomic Radius Trends--Table

51. Ionization Energy Definition of ionization energy
Ionization energy: the energy required to remove an electron from an isolated atom or molecule in the gaseous phase

52. Graph of Ionization Energies

53. Electron Affinity Definition of Electron Affinity
The electron affinity is a measure of the energy change when an electron is added to a neutral atom to form a negative ion

54. Graph of Electron Affinities

55. Trends in Electronegativity
Definition of electronegativity: a property of an atom which increases with its tendency to attract the electrons of a bond.

56. Trends in Electronegativity