1. Chapter 5 Electrons In Atoms 5.3 Atomic Emission Spectra
5.1 Revising the Atomic Model
5.2 Electron Arrangement in Atoms
5.3 Atomic Emission Spectra
and the Quantum
Mechanical Model
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2. Light and Atomic Emission Spectra
The Nature of Light
Light consists of electromagnetic waves.
Amplitude - wave’s height from zero to the crest.
Wavelength () - distance between the crests.
Frequency () - number of wave cycles to pass a given point per unit of time.
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3. Light and Atomic Emission Spectra
The Nature of Light
Speed of Light = wavelength x frequency
c =  108 m/s.
c = ln
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4. Light and Atomic Emission Spectra
Frequency () and wavelength () are inversely proportional.
As wavelength increases, frequency decreases.
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5. Calculating the Wavelength of Light
Sample Problem 5.2
Calculating the Wavelength of Light
Calculate the wavelength of the yellow light emitted by a sodium lamp if the frequency of the radiation is 5.09 × 1014 Hz (5.09 × 1014/s).
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6. Sample Problem 5.2
c = ln
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7. What is the frequency of a red laser that has a wavelength of 676 nm?
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8. What is the frequency of a red laser that has a wavelength of 676 nm?
c = ln c 
 =
 = = = 4.43  1014 m
c  108 m/s
  10–7 /s
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9. Light and Atomic Emission Spectra
The Nature of Light
When sunlight passes through a prism, different wavelengths separate into a spectrum of colors.
Red has the longest wavelength and lowest frequency.
Violet has the shortest wavelength (highest frequency)
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10. Light and Atomic Emission Spectra
Low energy
( = 700 nm)
High energy
( = 380 nm)
Frequency  (s-1)
3 x 106
3 x 1012
3 x 1022
102
10-8
10-14
Wavelength  (m)
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11. Light and Atomic Emission Spectra
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12. Light and Atomic Emission Spectra
When atoms absorb energy, their electrons move to higher energy levels. These electrons lose energy by emitting light when they return to lower energy levels.
No two elements have the same emission spectrum.
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13. Light and Atomic Emission Spectra
Electrons orbit around the nucleus in specific energy levels (ground state)
Electrons can absorb energy, become excited and move to a higher energy orbital (excited state)
Excited electrons release energy as photons of light in order to get back to the ground state.
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14. Light and Atomic Emission Spectra
A sample of H atoms receives energy from an external source.
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15. Light and Atomic Emission Spectra
The excited atoms release energy by emitting photons.
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16. Light and Atomic Emission Spectra
Atomic Emission Spectra of Hydrogen
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17. The Quantum Concept and Photons
The Quantization of Energy
The amount of radiant energy (E) of a single quantum absorbed or emitted by a body is proportional to the frequency of radiation (n).
h = Plank’s Constant = x J·s
E = hn
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18. The Quantum Concept and Photons
Photoelectric effect: electrons are ejected when light of sufficient frequency shines on a metal.
Light quanta are called photons
Energy of a photon = E = h x ⋎
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19. Calculating the Energy of a Photon
Sample Problem 5.3
Calculating the Energy of a Photon
What is the energy of a photon of microwave radiation with a frequency of 3.20 × 1011/s?
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20. E = h  = (6.626  10–34 J·s)  (3.20  1011/s) = 2.12  10–22 J
Sample Problem 5.3
E = h  = (6.626  10–34 J·s)  (3.20  1011/s)
= 2.12  10–22 J
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21. What is the frequency of a photon whose energy is 1.166  10–17 J?
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22. What is the frequency of a photon whose energy is 1.166  10–17 J?
E = h n
n = h E
 = = =  1016 Hz
E  10–34 J
h  10–17 J·s
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23. END OF 5. 3 EM Waves https://www. youtube. com/watch
END OF 5.3 EM Waves Fireworks
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