1. Chapter 6
Section 2

2. Sec 6.5 Quantum Mechanics and Atomic Orbitals
Wave functions – describes the behavior of the electron, denoted with the Greek letter, ψ
The wave function has a known energy, but the electron location is unknown, so the probability of its position in space is given by probability density, ψ2
Electron density – distribution map of the probability of finding the electrons at the points of space (probability density); high probability density=high electron density

3. Orbitals and Quantum Numbers
Orbitals – specific distribution of electron density in space (given by probability density); quantum mechanical model

4. Orbitals and Quantum Numbers
principal quantum number, n,
relates to the size and energy of the orbital: integral values of 1,2,3, etc.
an increase in n would mean a larger orbital, farther from the nuclear, and more energy (less tightly bound to nucleus)
There is no n= 0 value.
The ground state is n = 1
Excited states are n = 2, 3, 4 etc

5. Quantum Numbers Azimuthal quantum number, l,
* defines shape of the orbital
* designated by letters s, p, d, and f
s = 0
p = 1
d = 2
f = 3

6. Quantum Numbers Magnetic quantum number, ml ,
* describes orientation of orbital in space
* ranges from l and –l
Ex: if l = 3
then ml could be -3, -2, -1, 0, 1, 2, 3
Take a look at an orbital diagram.
Compare the number of ml options with the number of boxes per sublevel.

7. Quantum Numbers Spin quantum number, ms ,
* describes the direction of electron spin
There are 2 options for the ms value
+ ½ or - ½
The first electron in the orbital is spin up and the second electron is spin down.

8. Subshell – set or orbitals that have the same “n” and “l” values
Electron shells – a set of orbitals with the same value of n, such as 3s, 3p, 3d
Subshell – set or orbitals that have the same “n” and “l” values
The shell with the principal quantum number n, will have exactly n subshells
Each subshell has a specific number of orbitals. For a given l, there are 2l + 1 allowed values of m1
The total number of orbitals in a shell is n2

9. Ground State – when the electron is in the lowest energy orbital
Excited State – when the electron is in any other orbital

10. Representations of Orbitals

11. S- orbital Appears to be spherical Size increases as n increases
All s-orbitals are spherically symmetrical
Nodes = the intermediate regions where ψ2 goes to zero; the number of nodes increases with increasing values of n

12. P - orbitals

13. P- orbitals
Concentrated on two sides of the nucleus, separated by a node at the nucleus, “two lobes”
the orbitals of a given subshell have the same size and shape but differ in spatial orientation

14. D - orbitals

15. d- and f- orbitals
The different “d” orbitals in a given shell have different shapes and orientations in space
When “n” is equal to or greater than 3, the d-orbitals are present. There are 5 d orbitals
When n is equal to or greater than 4, there are 7 equal f-orbitals present

16. Sec 6.7 Orbitals in Many-Electron Atoms
The presence of more than one electron greatly changes the energies of the orbitals
The electron-electron repulsions cause different subshells to be at different energies

17. Effective Nuclear Charge
Each electron is simultaneously attracted to the nucleus and repelled by the other electrons
Energy of the electron can be estimated by how it interacts with the average environment (created by the nucleus and other electrons)

18. Effective Nuclear Charge
Zeff = Z – S
Z is # of protons in the nucleus
S is avg # of electrons between the nucleus and the electron in question.

19. Effective Nuclear Charge
Screening effect – the effect of inner electrons decreasing the nuclear charge experienced by outer electrons

20. Energies of Orbitals
In a many-electron atom, for a given value of n, Zeff decreases with increasing value
of l
The energy of an electron depends on the effective nuclear charge, Zeff
In a many-electron atom, for a given value of n, the energy of an orbital increases with increasing value of l
Degenerate - orbitals with the same energy
(i.e. there are 3 p orbitals or 5 d orbitals)

21. Shielding Example
1 1H He
9 19F Cl

23. Electron Spin and the Pauli Exclusion Principle
Electron spin is a property of the electron that makes it behave as though it were a tiny magnet. The electron behaves as if it were spinning on its axis, electron spin is quantized
Electron spin quantum number, ms , is a quantum number associated with the electron spin; two possible values + ½ or - ½

24. Pauli’s Exclusion Principle states that no two electrons in an atom can have the same values for n, l, ml, and ms
Places a limit of two on number of electrons that can occupy any one atomic orbital

25. 6.8 Electron Configurations
Electron configuration – way in which the electrons are distributed among the various orbitals of an atom
Represent electron configuration through an orbital diagram – each orbital represented by a box and each electron by a half arrow

26. 6.8 Electron Configuration

27. Electron Configurations
Practice writing electron configurations.
Practice writing abbreviated electron configurations.

28. 6.8 Electron Configuration
Paired – electrons in the same orbital
Unpaired – electron alone in an orbital
Hund’s rule – for degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin in maximized
Valence electrons - outer-shell electrons
Core electrons- electrons in inner shells

30. Period 4 and Beyond Elements known as:
* Transition Elements (and Metals)-Elements in which the “d” orbitals are filled
* Lanthanide Elements- Elements in which the 4f sub shell is partly occupied
* Actinide Elements- Elements in which the 5f orbitals are partly occupied

31. 6.9 Electron Configurations and the Periodic Table

32. e- configuration exceptions
If you apply the Aufbau principle to all known elements you will find that their electron configurations do not always agree with those determined experimentally. These exceptions do not have any major chemical consequences, though they are interesting to chemists.

33. Exceptions Cu The normal configuration: 1s2 2s2 2p6 3s2 3p6 4s2 3d9
The exception:
1s2 2s2 2p6 3s2 3p6 4s1 3d10
Why do you think this happens?

34. Exceptions Cr The normal configuration: 1s2 2s2 2p6 3s2 3p6 4s2 3d4
The exception:
1s2 2s2 2p6 3s2 3p6 4s1 3d5
Why do you think this happens?

35. Iron What are the 2 most common oxidation states of Iron?
Write electron configurations to support the existence of the 2 most common oxidation states.

36. Brown & LeMay Problems
7-9, 13, 16, 21, 22, 27, 31, 42-46, 61, 67, 71, 74, 79, 86