1. Electrons in Atoms Bohr Model of the Atom cjohannesson@satx.rr.com
Used by permission of Christy Johannesson
nucleus (+)
electron (-)
Courtesy Christy Johannesson

2. How color tells us about atoms
Atomic Spectrum

3. Prism
White light is made up of all the colors of the visible spectrum.
Passing it through a prism separates it.
Author: Thomas V. Green Jr.

4. If the light is not white
By heating a gas or with electricity we can get it to give off colors.
Passing this light through a prism does something different.
Author: Thomas V. Green Jr.

5. Atomic Spectrum Each element gives off its own characteristic colors.
Can be used to identify the atom.
How we know what stars are made of.
Author: Thomas V. Green Jr.

6. These are called line spectra unique to each element.
These are emission spectra
The light is emitted given off.
Author: Thomas V. Green Jr.

7. Line-Emission Spectrum
excited state
Wavelength (nm)
410 nm
486 nm
656 nm
434 nm
ENERGY IN
PHOTON OUT
Prism
Slits
ground state
Courtesy Christy Johannesson

8. Bohr Model
electrons exist only in orbits with specific amounts of energy called energy levels
Therefore…
electrons can only gain or lose certain amounts of energy
only certain photons are produced
Courtesy Christy Johannesson

9. 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

10. Other Elements
Each element has a unique bright-line emission spectrum.
i.e. “Atomic Fingerprint”
Helium
Bohr’s calculations only worked for hydrogen! 
Courtesy Christy Johannesson

11. Bohr’s Experiment
Kelter, Carr, Scott, Chemistry A Wolrd of Choices 1999, page 76
Animation by Raymond Chang – All rights reserved.

12. (a) Electronic absorption transition
(b) H2 emission spectrum (top), H2
absorption spectrum (bottom)
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

13. continuous spectrum gas absorption spectrum absorption spectrum
hot source
Source:
emission spectrum
emission spectrum

14. Hydrogen Spectral Lines
Lyman series
(ultraviolet)
Balmer series
(visible)
Paschen series
(infrared)
Frequency
(hertz)
1016
1015
1014
n =

15. HYDROGEN SPECTRAL LINES
(ultraviolet)
HYDROGEN SPECTRAL LINES
(visible)
(infrared)
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

16. Hydrogen Spectral Lines
Bohr’s model of the atom accounted mathematically for the energy of each of the transitions shown.
ionization IR
region E6 E5
656 nm
486 nm
434 nm
410 nm E4 E3
A B C D
Paschen (IR) E2
The energy of the light that is emitted from an atom is equal to the difference in energy between the excited state and the ground state. As Bohr predicted, the colors of light indicated by his theory matched those in the line spectrum of hydrogen.
• In 1885, Johann Balmer showed that the frequencies of the lines observed in the visible region of the spectrum of hydrogen fit a simple equation:
 = constant (1/22 – 1/n2) where n = 2, 4, 5, 6, and these lines are known as the Balmer series.
• Johannes Rydberg restated and expanded Balmer’s result in the Rydberg equation:
= (1/n21 – 1/n22),
where n1 and n2 are integers, n2 > n1, and , the Rydberg constant, has a value of x 107m-1.
A B C D E
Balmer (Visible)
Energy UV
region E1
A B C D E F
Lyman series (UV)
Davis, Metcalfe, Williams, Castka, Modern Chemistry, 1999, page 97

17. Electronic Transitions in the Excited Hydrogen Atom