1. ATOMIC STRUCTURE

3. Problem solving companion

4. Electromagnetic Radiation
Waves have a frequency and a wavelength
Use “nu”, , for frequency, and units are “cycles per sec” or Hertz (Hz)
Use “lamda”, , for wavelength,(m)
All radiation:  •  = c
c = speed of light = 3.00 x 108 m/sec

5. Electromagnetic Radiation
wavelength
Visible light
Wavelength, 
Ultaviolet radiation
Amplitude
Node
, for frequency

6. Electromagnetic Spectrum
 •  = c = 3.00 x 108 m/sec

7. Electromagnetic Spectrum
Short wavelength -->
high frequency
high energy
increasing frequency
increasing wavelength
Long wavelength -->
small frequency
low energy

8. Electromagnetic Radiation
Red light has  = 700 nm. Calculate the frequency.
 •  = c = 3.00 x 108 m/sec

9. Quantization of Energy
Max Planck ( )
An object can gain or lose energy by absorbing or emitting radiant energy in QUANTA.
Energy of radiation is proportional to frequency
E = h • 
h = Planck’s constant = x J•s

10. Quantization of Energy
E = h • 
Light with large  (small ) has a small E.
Light with a short  (large ) has a large E.

11. Photons and Energy  = 700. nm
Understand experimental observations if light consists of particles called PHOTONS of discrete energy.
PROBLEM: Calculate the energy of 1.00 mol of photons of red light.
 = nm
 = x 1014 sec-1

12. Energy of 1.00 mol of photons of red light.
E = h•
= (6.63 x J•s)(4.29 x 1014 sec-1)
= x J per photon
(2.85 x J/ph)(6.02 x 1023 ph/mol)
E per mol = kJ/mol
This is in the range of energies
that can break bonds.

14. Excited Gases & Atomic Structure

15. Atomic Line Emission Spectra and Niels Bohr
Bohr’s contribution;
a simple model of the atom
based on an understanding of the SHARP LINE EMISSION SPECTRA of excited atoms.
Niels Bohr
( )

16. Spectrum of Excited Hydrogen Gas
Excited atoms emit light of only certain wavelengths
The wavelengths of emitted light depend on the element.
Spectrum of Excited Hydrogen Gas

17. Line Spectra of Other Elements
Figure 7.9

18. Atomic Spectra and Bohr
One view of atomic structure in early 20th century was that an electron (e-) traveled about the nucleus in an orbit.
Any orbit should be possible and so is any energy.

19. Atomic Spectra and Bohr
If e-’s are in quantized energy states, then ∆E of states can have only certain values.
This explain sharp line spectra.

20. Atomic Spectra and Bohr.
Atomic Spectra and Bohr n = 1 2 E - C ( / ) N R G Y
Calculate ∆E for e- “falling” from high energy level (n = 2) to low energy level (n = 1).
∆E = Efinal - Einitial
∆E = negative therefore
Note that the process is EXOTHERMIC

23. Line Emission Spectra of Excited Atoms
High E
Short 
High 
Low E
Long 
Low 
Visible lines in H atom spectrum are called the BALMER series.

24. Problem solving companion

26. Atomic Line Spectra and Niels Bohr
Rec’d Nobel Prize, 1922
Problems with theory —
theory only successful for Hydrogen.
So, we go on to QUANTUM or WAVE MECHANICS
Niels Bohr
( )

27. n (major) ---> shell l (angular) ---> subshell
QUANTUM NUMBERS
The shape, size, and energy of each orbital is a function of 3 quantum numbers:
n (major) ---> shell
l (angular) ---> subshell
ml (magnetic) ---> orbital within a
subshell

28. Shells
n = 1
n = 2
n = 3
n = 4

29. “Your Best Friend”
Periodic table