1. The Bohr Atomic Theory
Moving toward a new atomic model….

2. Back to Planck and Einstein…
Once the ideas of Planck and Einstein (re: quanta) had been accepted, the study of atomic and molecular structure began.
With the invention of the spectroscope (Bunsen & Kirchhoff, Fig 2/175) the quantitative analysis of the atomic spectra of elements began.
The range of frequencies of light emitted or absorbed by an atom is called its spectrum (pl. spectra).

3. Spectroscope:
If the light emitted by an gaseous element is passed through a spectroscope, it produces a bright line atomic spectrum that is characteristic to that element (fingerprint).

5. When an element is excited, atoms absorb energy, then as e-s fall back to “ground state”, energy is released in the form of light
Visible line spectra called emission spectra of discrete(distinct) coloured lines - specific wavelengths = specific energies, not a continuous spectrum
Atoms absorbing energy - discrete black lines - absorption spectrum.

6. Emission Spectra:

7. Emission Spectrum (Na)
Absorption Spectrum (Na)

8. Recall: Problems with Rutherford’s planetary model re stability
Atomic spectra not explained - b/c if Rutherford’s model was correct, continual acceleration means continuous spectrum of light - spectral evidence shows this is not the case!

9. The Confusion About Bright Line Spectra
Niels Bohr – Danish physicist, working for Thomson (raisin bun guy) suggested that line spectra could be explained using ideas from the quantum theory!
Bohr began to work with Rutherford (1913) - modified Rutherford's atomic theory.

10. Bohr's Model of the Atom
54 years after line spectra were first described qualitatively by Bunsen and Kirchhoff.
28 years after line spectra were described quantitatively by Jacob Balmer.

11. Bohr's Postulates
The atom has only specific allowable energy levels (stationary states) & each stat.state corresponds to the e-s occupying fixed circular orbits around nucleus.
While in stat.state, atoms do not emit energy.
An e- changes stat. state by emitting or absorbing a specific quantity of energy that is exactly equal to the difference in energy between the two stat. states. (Fig 4/176)

12. Higher energy level  lower: loss of energy in the form of light
Lower energy level  higher: energy is absorbed
Only certain quanta of light can be emitted or absorbed by an atom, so the energy of the electrons inside the atom is also quantized. In other words, an electron can only have certain allowable energies. (staircase analogy).

13. http://www. upscale. utoronto

14. H-spectra example
If an e- that has been excited to the 3rd NRG level & then falls from 3rd to 2nd NRG level, emits a photon of red light with wavelength 656nm & red line is visible on emission spectrum.
Each line on e. spectrum corresponds to a specific energy transition (I.e. 4th to 2nd = green, 5th to 2nd = blue)

15. Hydrogen

16. Evaluating Bohr's Model:
Was reasonably successful at explaining Mendeleev's Periodic Table
Periods result in filling an additional energy level (i.e. Na: 3 levels)
A period ends when the maximum # of electrons in the valence level is reached (i.e. 2, 8, 8, 18)
BUT - could only explain one-electron systems (so still more complex than he thought).

17. Bohr Energy-Level Diagrams
Same rationale as orbit diagrams, but the focus is on the energy that the electrons have, rather than the position of the electrons, so orbits are not used.
For a Bohr energy level diagram, start from the bottom and work your way up!

18. Example:
5e- (3rd energy level) (from group 15)
8e- (2nd energy level) (from 8 elements in period 2)
2e- (1st energy level) (from 2 elements in period 1)
15p+ (protons) (from atomic number)
P (symbol) (from PT)
phosphorous (name) (lowercase)
Note: The diagram is drawn with the energy levels appearing to be evenly spaced, BUT we know this is not the case, b/c line spectra evidence tells us that energy levels are increasingly closer together at higher energy levels!!

19. Homework:
180 #1-5,8,9