1. Arrangement of Electrons in Atoms
The Development of a New Atomic Model

2. Light Before 1900, scientists thought that light behaved only as wave
discovered that also has particle-like characteristics

3. Light is electromagnetic radiation.
form of energy that acts as a wave as it travels
includes: gamma, X rays, ultraviolet (UV), visible and infrared (IR) light, microwaves, and radio waves
All forms are combined to form electromagnetic spectrum

4. Light as a Wave

5. c = λf Wave Properties: all form of EM radiation travel at a
speed (c) of 3.0 x 108 m/s in a vacuum
wavelength: (λ) distance between points on adjacent waves; in nm (109nm = 1m)
frequency: (f) number of waves that passes a point in a second, in waves/second
Inversely proportional!
c = λf

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

8. E = hf Quantum Theory Max Planck (1900) German physicist
Observed - emission of light from hot objects
Concluded –energy is emitted in small, specific amounts (quanta)
Quantum - minimum amount of energy change
Max Planck
Planck’s constant (h) =
6.626 x J*s
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.
When studying elements, scientists found that when certain elements were heated in a flame those elements emitted varying colors of light.
Light carries energy
Something in the atom must be carrying energy in order to emit light
E = hf
Courtesy Christy Johannesson

9. Photoelectric Effect Albert Einstein (1905)
Observed - photoelectric effect
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

10. Photoelectric Effect No electrons are emitted Electrons are emitted
Bright
red light
infrared rays
Dim
blue light
ultraviolet rays or or
Metal plate
Metal plate

11. Photoelectric Effect
when light is shone on a piece of metal, electrons can be emitted
no electrons were emitted if the light’s frequency was below a certain value
scientists could not explain this with their classical theories of light

12. Solar Calculator
Solar Panel

13. Photoelectric Effect Einstein added on to Planck’s theory in 1905
suggested that light can be viewed as stream of particles
photon- particle of EM radiation having no mass and carrying one quantum of energy
energy of photon depends on frequency

14. Photoelectric Effect
EM radiation can only be absorbed by matter in whole numbers of photons
when metal is hit by light, an electron must absorb a certain minimum amount of energy to knock the electron loose
this minimum energy is created by a minimum frequency
since electrons in different metal atoms are bound more or less tightly, then they require more or less energy

15. H Line-Emission Spectrum
Why had hydrogen atoms only given off specific frequencies of light?
current Quantum Theory attempts to explain this using a new theory of atom

17. H Line-Emission Spectrum
ground state- lowest energy state of an atom
excited state- when an atom has higher potential energy than it has at ground state
(excited by heat or electricity)
line-emission spectrum- pattern of wavelengths of light created when visible light from excited atoms is shined through a prism

18. Line-Emission Spectrum
pattern of wavelengths of light created when visible light from excited atoms is shined through a prism
excited state
Wavelength (nm)
410 nm
486 nm
656 nm
434 nm
ENERGY IN
PHOTON OUT
Prism
Slits
ground state
Courtesy Christy Johannesson

19. H Line-Emission Spectrum
scientists using classical theory expected atoms to be excited by whatever energy they absorbed
continuous spectrum- emission of continuous range of frequencies of EM
radiation

20. H Line-Emission Spectrum
when an excited atom falls back to ground state, it emits photon of radiation
the photon is equal to the difference in energy of the original and final states of atom
since only certain frequencies are emitted, the differences between the states must be constant

21. Bohr Model created by Niels Bohr (Danish physicist) in 1913
linked atom’s electron with emission spectrum
electron can circle nucleus in certain paths, in which it has a certain amount of energy

22. Bohr Model Can gain energy by moving to a higher rung on ladder
Can lose energy by moving to lower rung on ladder
Cannot gain or lose while on same rung of ladder

23. Bohr Model
a photon is released that has an energy equal to the difference between the initial and final energy orbits

24. Bohr Model Energy of photon depends on the difference in energy levels6
Energy of photon depends on the difference in energy levels
Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom 5 4 3 2 1
nucleus
Courtesy Christy Johannesson

26. Other Elements
Each element has a unique bright-line emission spectrum.
i.e. “Atomic Fingerprint”
Helium
Bohr’s calculations only worked for hydrogen! 
Didn’t explain chemical behavior of atoms.
Courtesy Christy Johannesson

27. Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

28. Electrons as Waves In 1924, Louis de Broglie (French scientist)
suggested the way quantized
electrons orbit the nucleus is similar to behavior of wave
electrons can be seen as waves confined to the space around a nucleus
waves could only be certain frequencies since electrons can only have certain amounts of energy

29. Uncertainty Principle
In 1927 by Werner Heisenberg (German theoretical physicist)
electrons can only be detected by their interaction with photons
any attempt to locate a specific electron with a photon knocks the electron off course
Heisenberg Uncertainty Principle- it is impossible to know both the position and velocity of an electron

30. Electrons as Waves
shows that anything with both mass and velocity has a corresponding wavelength

31. Schrödinger Wave Equation
In 1926, Erwin Schrödinger
(Austrian physicist)
his equation proved that
electron energies are quantized
only waves of specific energies provided solutions to his equation
solutions to his equation are called wave functions

33. Teacher, may I be excused? My brain is full!