1. Quantum Mechanical Model of the Atom

2. Purpose of Lesson
To learn the structure of atomic orbitals because they are essential to understanding chemical bonding.

3. Rutherford-Bohr model
Bohr changed Rutherford’s model of the atom stating that electrons move around the nucleus in orbits of fixed size and energy.
Electron energy in this
model is “quantized.” –
electrons can’t occupy values of
energy between fixed energy
levels.

4. Rutherford-Bohr model

5. Rutherford-Bohr diagrams
Putting all this together, we get R-B diagrams
To draw them you must know the # of protons, neutrons, and electrons (2,8,8,2 filling order)
Draw protons (p+), (n0) in circle (i.e. “nucleus”)
Draw electrons around in shells He Li
3 p+
4 n0
2e– 1e–
Li shorthand
3 p+
4 n0
2 p+
2 n0
Draw Be, B, Al and shorthand diagrams for O, Na

6. Be B Al O Na 8 p+ 11 p+ 8 n° 12 n° 4 p+ 5 n° 5 p+ 6 n° 13 p+ 14 n°
2e– 8e– 1e– Na
8 p+
8 n°
2e– 6e– O

7. Rutherford-Bohr Model
Bohr’s atomic model described how electrons could move from energy level, to energy level.
When electrons absorbed energy, they moved to a higher energy level.
As they gave off energy, they emitted a photon of light and fell back down to a lower energy level (becoming more stable).

8. The degree to which they move from level to level determines the frequency of light they give off.
Figure 5-14

9. Limitations of the Bohr Model
Problem…Moving electrons should give off energy and collapse into the nucleus; The Bohr Model didn’t work for heaver atoms!

10. Quantum Mechanical Model of the Atom
Bohr was correct in assigning energy levels to the electrons BUT he was not correct in assuming that electrons moved like planets around the nucleus.
Now we know that electrons move in less predictable ways.

11. Quantum Mechanical Model
Describes the possible locations of electrons around the nucleus.
The cloud is denser where the probability of finding an electron is higher.
Does NOT show the exact location of electrons

12. Quantum Mechanical Model - Shells
Electrons in an atom can be assessed according to the shell, subshell, and orbital to which they are assigned.
Shells are numbered as n=1,2,3,4, etc. and increase in size and energy as they get further away from the nucleus.

13. Shells

14. Shells Correspond to Periods on the Periodic Table

15. Quantum Mechanical Model - Subshells
Shells can be subdivided into subshells.
The maximum number of subshells is equal to the shell number. For example, when n=1 (first shell), only one subshell is possible;
When n=2 (second shell), two subshells are possible.

16. Types of Subshells There are 4 four types of subshells.
Denoted by the letters s, p, d, and f. Each subshell has a max. number of electrons which it can hold:
s - 2 electrons, p - 6 electrons, d - 10 electrons, and f - 14 electrons.
The s subshell is the lowest energy subshell and f subshell is the highest.

17. Remember - the shell number is equal to the possible number of subshells:
when n=1, the only subshell possible is the 1s subshell.
When n=2, two subshells are possible the 2s and 2p.
When n=3, three subshells are possible the 3s, 3p, and 3d.

18. Subshells
n = 1 has 1 subshell (1s) n = 2 has 2 subshells (2s and 2p) n = 3 has 3 subshells (3s, 3p, 3d)

19. Subshells
n = 1 has 1 subshell (1s) n = 2 has 2 subshells (2s and 2p) n = 3 has 3 subshells (3s, 3p, 3d)

20. Subshells
n = 1 has 1 subshell (1s) n = 2 has 2 subshells (2s and 2p) n = 3 has 3 subshells (3s, 3p, 3d)

21. Subshells
n = 1 has 1 subshell (1s) n = 2 has 2 subshells (2s and 2p) n = 3 has 3 subshells (3s, 3p, 3d)

22. Orbitals Each subshell is further divided into orbitals.
An orbital is a region of space in which an electron can be found.
Only two electrons are possible per orbital.
Thus, the s subshell may contain only one orbital and the p subshell may contain three orbitals.

23. Orbitals
2 electrons/orbital = 6 electrons total in the 2p subshell
n = 1 has 1 subshell (1s) n = 2 has 2 subshells (2s and 2p) n = 3 has 3 subshells (3s, 3p, 3d)

24. Quantum Mechanical Model
The wave function predicts the three-dimensional region around the nucleus called the atomic orbital.

25. Each orbital has its own distinct shape.
An s orbital found in a s subshell is spherical.
p orbitals found in p subshells are two-lobed
d orbitals found in d subshells are four-lobed.

26. Shapes of Orbitals

27. Orientation of Orbitals
Since there are three possible orbitals per p subshell, each orbital adopts its own orientation.
The px orbital lies along the x axis, the py orbital lies along the y axis, and the pz orbital lies along the z axis.

28. Orientation of Orbitals

29. Order of Subshells—Periodic Table

30. Day 2: Electron Configuration

31. Electron Configuration Rules
There are rules for filling electrons into orbitals. 1. Aufbau Principle: Lowest energy orbitals fill first. 2. Pauli Exclusion principle 3. Hund’s Rule See handout with more details

32. Electron Configuration
When writing electron configurations for atoms:
Record the principle quantum # “n” (shell)
Followed by the letter of the subshell
Followed by a superscripted number that represents the total number of electrons in that subshell.

33. Example #1
Carbon atom with 6 electrons would have the electron configuration:
1s22s22p2

34. Aufbau Diagram for Carbon
Carbon: 1s22s22p2

35. Orbital Diagram

36. Example #2
A Chorine atom with 17 electrons would have the electron configuration:
1s22s22p63s23p5

37. Aufbau Diagram for Carbon
Chlorine: 1s22s22p63s23p5

38. Orbital Diagram

39. Example #3
A Zinc atom with 30 electrons would have the electron configuration:
1s22s22p63s23p64s23d10
Condensed: [Ar] 4s23d10

40. Aufbau Diagram for Carbon
Chlorine: 1s22s22p63s23p5

41. Orbital Diagram