1. Bohr Model Deficiencies
Experimental:
The calculated line spectra results for atoms with more electrons didn’t match observations.
Only works for 1 electron species, H, Li+, etc
Doesn’t account for spectral lines that appear as
2’s, 3’s, etc.

2. Observed periodic structural trends in both valence level and electrons differed beyond Atomic Number 18.
Bohr’s model predicts 2n2 valence electrons, but the periodic table predicts a maximum of 8.
2, 8, 18, 32, 50 vs 2, 8, 8, 8

3. Theoretical:
Equates position with energy and does not allow Potential and Kinetic Energy inter-conversion.
The electron stays at the same distance from the nucleus, but doesn’t move?

4. Fails to account for how electrons make energy (orbit) transitions if they are not permitted to be at any distance but the allowed orbit distance.
The electron can’t “jump” creating line spectra, which is the basis of his theory.

5. Probability vs Probability Density
The probability graph shows the most likely place to find the electron.
But, it is not in the area closest to the nucleus.

6. The Scatter Diagram The probability graph doesn’t tell the whole story
The scatter diagram of the stone tosses (electrons) shows that the density is actually close to the nucleus.

7. Probability Density
The area around the nucleus is smaller than the surrounding rings.
When the density of throws is graphed, it shows that the electron is actually as close to the nucleus as expected.

8. PROBABILITY MODEL OR QUANTUM MECHANICS
Electrons can be anywhere in their orbital, traveling in any direction with a certain, fixed total energy.
As the electron travels, its PE and KE constantly exchange but the total electron energy remains fixed. Called the Principle Energy.
Some locations about the nucleus are more probable than others.

9. Position and direction of travel may not be simultaneously known (Heisenberg Uncertainty Principle).
An orbital is an equation solution to a more complicated equation called the Schroedinger Equation.
To solve the Schroedinger Equation and generate orbital equations, four separate constants called Quantum Numbers must be assigned values. These assigned values identify the specific orbital.
An orbital equation is its most probable physical location over a period of time as a 3D mapping or scatter diagram.

10. Scanning Tunnelling Microscopy (STM)
SCH 4U1

11. STM Image of Graphite

12. STM Image of Quantum Corral
Iron on Copper

13. STM of Two Electron Corral
Iron on Copper

14. Quantum Numbers
Each Bohr orbit is viewed as a Principle Energy Level defined by Principle Quantum Number, n where n = 1, 2, 3, 4....
Each Principle Energy Level contains “n” energy sub levels or orbital types defined by the Secondary or Orbital Quantum Number, l where l = 0 to n-1.
n l = 0 l = 1 l = 2 l = 3
1 sharp
2 s principal
3 s p diffuse
4 s p d fundamental
These 2 quantum numbers give the 1st 2 parts of an electron configuration. eg. 2s22p3

15. Orbitals for Electrons
s, p, d, and f

16. Each Principle Energy Level contains a total of “n2” orbitals defined by their spatial orientation and the Magnetic Quantum Number, ml where -l ≤ ml ≤+l
1 of type s ml = 0
3 of type p ml = -1, 0, +1
these give the different orbital boxes
5 of type d ml = -2, -1, 0, +1, +2
7 of type f ml = -3, -2, -1, 0, +1, +2, +3

17. The sublevels now explain the line spectra{ { { { f s p d

18. Each orbital describes the behavior of two electrons at most, defined by the Spin Quantum Number, mS = +1/2 or -1/2
This gives the up and down arrows, 
Now each electron has its own set of Quantum numbers.
These numbers can be entered into the Schroedinger equation and the orbital for that electron can be mapped out and drawn
We can also give the quantum numbers for any electron in an orbital box configuration.

19. Photon Electron Spectroscopy, PES
Is based upon a single photon in/electron out.
The energy of the photon is E = hv
For core e-, X-rays are used (XPS), for valence electrons, UV radiation is used, (UPS).
Overall process: A hv  A e-
NRG conservation: E(A) + hv = E(A+) + E(e-)
e- NRG is KE, KE = hv - ( E(A+) + E(e-) )
difference is BE KE = hv - BE

20. Where: h = Planck constant (6.62 x 10-34 J●s)
v = frequency (Hz) of the radiation
KE = kinetic energy of the emitted e-
BE = binding energy or IE
Then BE = hv – KE

21. Principles used to Interpret PES Spectra:
All e- in the same subshell have the same BE (IE)
There is 1 peak for every different type of e-
Peak height shows the relative # of e-s in each subshell of an atom
As effective nuclear charge , the BE will .

22. 1s2
2s2
2p6
3s2
The element is Mg

23. 1s2
2p4
2s2
The element is O