1. Atomic Structure II (Ch.5)
Introduction:
Wave-Particle Nature of
Light
Discoveries about the nature of light led to
greater understanding of the properties of
electrons and atomic structure.

2. 1800’s: Electrons were thought of as particles;
light was thought of as waves.
Experimental observations can support both these theories.
Early 1900’s: New experiments showed that electrons have
some wave-like properties and light has some
particle-like properties.
These discoveries led to the idea
that both electrons and light
have a dual wave-particle nature.
(This was the beginning of
quantum physics.)

3. I. Light is a form of electromagnetic radiation.
A. electromagnetic radiation (EMR)- energy that has
wave characteristics as it travels through space.

4. 1. Visible light is one form of electromagnetic radiation;
others include x-rays, ultraviolet light, infrared light,
TV and radio waves.
2. Electromagnetic spectrum – consists of all the forms of
EMR arranged in order by wavelength.
3. All types of electromagnetic radiation travel at a speed
of 3.00 x 108 m/s. This is called the speed of light (c).

5. Electromagnetic Spectrum

6. Waves
Waves transfer energy from one place to
another
B. Electromagnetic radiation has the
following wave characteristics:

7. 1. Wavelength,  (lambda) = distance between
successive points
10
m
2m

8. 2. Frequency,  (nu) = the number of waves
that pass a given point per unit of time;
Cycles per second
t=0 t=5 t=0 t=5

9. Units for Frequency
• 1/s
• s-1
• hertz, Hz
FM Radio stations are identified by their frequency in MHz. ;
AM Radio stations are identified by their frequency in kHz.

10. 3. Wave Velocity
• Velocity of a wave (m/s) =
wavelength (m) x frequency (1/s)
• c = 
• 3.00x108 m/s = 

11. Because all electromagnetic waves travel at
the speed of light, wavelength is determined
by frequency:
Low frequency = long wavelengths
High frequency = short wavelengths
Wavelength and frequency are inversely proportional.

12. 1. The yellow light given off by a sodium vapor light
has a wavelength of 589 nm. What is the lights frequency?
2. A laser used to fuse detached retinas has radiation with
a frequency of 4.7 x 1014 Hz. What is its wavelength in nm?

13. 3. Calculate the wavelength (in nm) of a microwave that has
the frequency of 3.25 x 1010 Hz.
4. What is the frequency of blue light that has
a wavelength of 475 nm?
9.23 x 106 nm
6.31 x 1014 Hz

14. 1. An x-ray has a frequency of 2.5 x 1018 Hz.
What is the wavelength of the x-ray?
2. A red light has a wavelength of 725 nm.
What is the frequency of the red light?

15. C. Prisms
A prism can be used to “spread out” the wavelengths of light.

16. Continuous spectrum – a spectrum in which all the
wavelengths within a given range are included.
A rainbow would be an example of a continuous spectrum
of visible light.

18. Light shows the wave property of Interference.
(Physics 2000 – Interference Experiments)
D.

19. II. Light as a Particle
two properties of light could not be explained in terms of waves:
A) emission of light by hot objects; B) the photoelectric effect.

20. A. (1900) Max Planck says energy can only be
released or absorbed by atoms in
"packets of energy" called quanta.
The energy of a quantum (E) is given by the equation:
E = h  = frequency;
h = Planck's constant: 6.63 x 10-34 Js
EX: Calculate the energy of a quantum of radiation whose
frequency is 3.0 x 1011 Hz.

21. • Higher-frequency electromagnetic waves
have higher energy than lower-frequency
electromagnetic waves
• All forms of electromagnetic energy
interact with matter, and the ability of these
different waves to penetrate matter is a
measure of the energy of the waves

22. Photoelectric Effect Animation
B. Photoelectric Effect - the phenomena whereby a metal
will emit electrons if light with a certain minimum frequency
is shined on it.

23. (1905) Einstein explains photoelectric effect by saying
that light consists of quanta.
Only if the light has a high enough frequency will the quanta
have enough energy to dislodge an electron.
Einstein called these quanta or "particles" of light photons.
Light (and all electromagnetic radiation) has
both wave-like and particle like properties.

24. EX Problem 1: A purple light has a frequency of
7.5 x 1014 Hz. What is the energy of a photon
of purple light?
EX2: A green light has a wavelength of 500. nm.
Calculate the energy of a photon of green light.
5.0 x 10-19 J
3.98 x 10-19 J

25. •A blue light has a frequency of 6.7 x 1014 Hz.
Calculate the energy of a photon of blue light.
•Radio signals from station WIP have a frequency of
6.1 x 105 s-1. Calculate the energy of a photon
of this radio signal.
•A yellow light has a wavelength of 575 nm.
Calculate the energy of a photon of this light.
4.4 x 10-19 J
4.0 x 10-28 J
3.46 x 10-19 J

26. III. Line Spectra (bright-line spectra, atomic emission spectra)
1. Line Spectra ( or bright-line spectra) – these lines of color are produced
when the light from an element is passed through a prism.
2. Each element produces a characteristic set of lines called the
line spectrum of the element.
Elements can be identified by their line spectrum.
3. Each line in a line spectrum represents a particular wavelength and
frequency of light given off.
You can calculate the energy of the photons of light using E=h
4. Bohr used the line spectrum of hydrogen to devise his model of the atom.

27. III. Line Spectra
1. Physics 2000 – Quantum Atom - Line Spectra 2. Line Spectra Animation

29. Quantum Model of Atom
A. (1924) Louis de Broglie says if light has a dual wave-particle nature,
then matter (electrons) could also have a dual wave-particle nature.
I. Quantum Mechanics

30. Electron Interference Animation(Web) Drive
Experiments showed that electrons display
the wave properties of diffraction and interference.

31. B. Heisenberg Uncertainty Principle
It is impossible to know exactly both
the location and velocity of an electron
at the same time.

32. C. Schrodinger Wave Equation
Describes the probability of where an electron can be found.
Schrodinger’s equation works for all atoms.
D. Quantum theory – describes mathematically the wave
properties of electrons. In the quantum model of the
atom, electrons are found in orbitals, not specific orbits.

34. Quantum Theory

35. Orbitals