1. The Bohr model for the electrons
Electronic structure – how the electrons are arranged inside the atom
Applying the quantum principle of energy
Two parameters:
Energy
Position

2. Learning objectives Describe the basic principles of the Bohr model
Distinguish between the “classical” view and the “quantum” view of matter
Define atomic orbitals
Distinguish between the Bohr orbit and atomic orbital
Apply quantum numbers and atomic orbitals to building atoms and the periodic table
Describe periodic trends in terms of electronic structure

3. Bohr’s theory of the atom: applying photons to electronic structure
Electrons occupy specific levels (orbits) and no others
Orbits have energy and size
Electron excited to higher level by absorbing photon
Electron relaxes to lower level by emitting photon
Photon energy exactly equals gap between levels
Larger orbits are at higher energy

4. Size of energy gap determines photon energy
Small energy gap, low frequency, long wavelength (red shift)
High energy gap, high frequency, short wavelength (blue shift)

5. The full spectrum of lines for H
Each set of lines in the H spectrum comes from transitions from all the higher levels to a particular level.
The lines in the visible are transitions to the second level

6. The Bohr orbits Bohr orbits have quantum numbers n n = 1 (capacity 2)

7. Bohr orbits and the periodic table
Elements in the same group have the same number of electrons in outer Bohr orbit

8. Successes and shortcomings of Bohr
Couldn’t explain why orbits were allowed
Only successful agreement with experiment was with the H atom
Introduced connection between spectra and electron structure
Concept of allowed orbits is developed further with new knowledge
Nonetheless, an important contribution, worthy of the Nobel prize

9. Electrons are waves too!
Life at the electron level is very different
Key to unlocking the low door to the secret garden of the atom lay in accepting the wave properties of electrons
De Broglie wave-particle duality
All particles have a wavelength – wavelike nature.
Significant only for very small particles – like electrons or photons
As mass increases, wavelength decreases
Electrons have wavelengths about the size of an atom
Electrons are used for studying matter – electron microscopy

10. Electron microscopes can peer within – waves interacting with matter

11. Heisenberg Uncertainty Principle: the illusive electron
We can predict the motion of a ball;
But not an electron: problems locating small objects

12. The Quantum Mechanics: waves of uncertainty
System developed that incorporated these concepts and produced an orbital picture of the electrons
No longer think of electrons as particles with precise location, but as waves which have probability of being in some region of the atom – the orbital
Impossible with the classical mechanics of Newton

13. Orbitals are described by quantum numbers
Each orbital has unique set
1s, 2p, 3d etc.
Number describes energy
Letter describes shape
S zero dimensions
P one dimension
D two dimensions
F three dimensions

14. Getting from the orbitals to the elements
All elements have the same set
Atomic number dictates how many are filled – how many electrons are added
Filling orbitals follows a fixed pattern: lowest energy ones first

15. Orbital energy levels in H and other elements

16. How many per orbital? Electrons share orbitals (only two allowed)
A consequence of “spin”

17. How many electrons can be added to the orbitals
1s, 2s, 3s etc. 2 electrons
2p, 3p, 4p etc. 6 electrons
3d, 4d etc electrons
4f, 5f etc electrons

18. Add electrons to the orbitals – lowest first2p 3d 3p 4p 4s 3s 2s 1s
H(z = 1)

19. Fill lowest orbital 2p 3d 3p 4p 4s 3s 2s 1s
He(z = 2)

20. Begin next orbital 2p 3d 3p 4p 4s 3s 2s 1s
Li(z = 3)

21. Fill 2s 2p 3d 3p 4p 4s 3s 2s 1s
Be(z = 4)

22. Begin filling 2p 2p 3d 3p 4p 4s 3s 2s 1s
B(z = 5)

23. Electrons don’t like to pair
C(z = 6)

24. 2p 3d 3p 4p 4s 3s 2s 1s
O(z = 8)

25. 2p 3d 3p 4p 4s 3s 2s 1s
F(z = 9)

26. Filled 2p – neon unreactive4s 3s 2s 1s
Ne(z = 10)

27. Shape of the periodic table explained by orbital picture
6 groups
2 groups
10 groups
14 groups

28. Shells: echoes of the Bohr orbits
The orbitals with the same Principal Quantum number (1,2,3 etc) are grouped into shells
Filled shells have special significance

29. The periodic law

30. Ionization energy and the periodic law
Ionization energy is energy required to remove electron from the neutral atom