1. Chapter 5

2. Chapter 5

3. Quantum Theory and Electron Configurations

4. It’s all about color… In terms of atomic models, so far:
Dalton (1803) = Tiny, solid particle
Thomson (1897) = “Plum Pudding” model – Electrons stuck on the outside of a big positive charge
Rutherford (1911) = Positively-charged nucleus with electrons moving around it
Rutherford’s model of the atom not quite right
Could not explain chemical properties of elements
Could not explain color changes when metal is heated

5. Bohr Model of the Atom Niels Bohr’s model of the atom
Electron found only on specific, circular paths around nucleus
Each orbit has fixed energy level
Hypothesis: When electrons are excited (added energy), jump into higher energy levels. When they moved back into lower energy levels - gave off light.
Electrons do not exist between levels (think of rungs on a ladder)
Electrons absorb and emit only certain quanta (amounts) of energy
Quantum of energy = fixed amount of energy required to move from one energy level to another energy level

6. Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
Chapter 5

7. Bohr’s Planetary Model of the Atom
Electrons must have enough energy to keep moving around the nucleus
Electrons orbit nucleus in defined energy levels, just like planets orbit the sun
Each energy level assigned a principal quantum number n.
Lowest energy level called ground state (n=1)
Higher energy levels (n=2, 3, 4...) excited states
Model worked OK for hydrogen but not so good for other elements
Nucleus
n = 1
n = 2

8. } Bohr’s Model Fifth Fourth Increasing energy Third Second First
Further away from the nucleus means more energy.
There is no “in between” energy
Energy Levels
Fifth
Fourth
Third
Increasing energy
Second
First
Nucleus
Chapter 5

9. Electron starts on lowest energy level (ground state)
Higher energy levels = excited states
Add energy to electron – moves to excited state
Energy levels are not evenly spaced
Energy Level 2
Energy Level 3
Energy Level 1
Nucleus

10. Electron starts on lowest energy level (ground state)
Higher energy levels = excited states
Add energy to electron – moves to excited state
Energy
Energy levels are not evenly spaced
Electron returns to lower state – emits/gives off quantum of energy
Energy Level 2
Energy Level 3
Energy Level 1
Nucleus

11. Bohr used this theory to explain the lines in the atomic emission spectra for hydrogen
Chapter 5

12. Each of these lines corresponds to different energy changes
434 nm
656 nm
486 nm
410 nm
Chapter 5

13. Chapter 5

14. Chem I - Mon, 9/22/15
Do Now
Get to work on the PEN worksheet from last class
Homework
MEAL paragraph if not finished
Agenda
History
Intro to quantum
Electron Config

15. Quantum-Mechanical Model of the Atom
Since the Bohr model had a very limited use, a new and very different model of the atom exists
The Quantum Mechanical Model (1926) contains:
Quantum energy levels
Dual wave/particle nature of electrons
Electron clouds
In the new model, don’t know exactly where electrons are - only know probabilities of where they could be

16. Heisenberg Uncertainty Principle
Heisenberg Uncertainty Principle = impossible to know both the velocity (or momentum) and position of an electron at the same time

17. Quantum-Mechanical Model of the Atom
Orbital = region around nucleus where an electron with a given energy level will probably (90%) be found
Four kinds of orbitals
s - spherical in shape, lowest orbital for every energy level
p - dumbbell shaped, second orbital
d - complex “flower” shape, third orbital
f - very complex shape, highest orbital

18. s-orbitals All s-orbitals are spherical.
As n increases, the s-orbitals get larger.

19. p- orbitals Three p-orbitals: px, py, and pz
Lie along the x-, y- and z- axes of a Cartesian system.
Dumbbell shaped, gets larger as n increases

20. d and f - orbitals
There are five d and seven f-orbitals.

23. Quantum Mechanical Model
Principle Energy Levels (n)
Labeled from 1-7
First energy level is n=1
Contains sublevels (s, p, d and f)
Each energy level contains the number of sublevels equal to its value for n
If n=3, there are three sublevels

24. Chapter 5

25. Quantum Mechanical Model
In each sublevel there are atomic orbitals
Atomic orbitals – describe a space where an electron is likely to be found
Type of subshell
Shape of orbitals
Number of orbitals
Orbital ‘names’ s
Spherical 1 p
Dumbbell 3
px, py, pz d
Cloverleaf (and one donut) 5 f
Multi-lobed 7

26. Quantum Mechanical Model
Each orbital can contain two electrons.
Since negative-negative repel, these electrons occupy the orbital with opposite spins.

27. Quantum Mechanical Model
The total number of orbitals of an energy level is n2.
For the third principle energy level, n=3, which means there are 9 orbitals
These orbitals are 3s, 3px, 3py, 3pz and the 5 d orbitals
Remember, we no longer think of orbitals as concentric circles, but we can say that n=4 extends farther from the nucleus than n=1.

28. Valence Electrons
Only those electrons in the highest principle energy level

29. Electron Configuration and Orbital Notation
Aufbau Principle – electrons fill lower energy orbitals first, “bottom-up”
n=1 fills before n=3
Will an electron fill the 1s or the 2s orbital first?
Energy
2px
2py
2pz 2s 1s

30. Electron Configuration &Orbital Notation
Hund’s Rule – electrons enter same energy orbitals so that each orbital has one electron before doubling up
Each of the first electrons to enter the equal energy orbitals must have the same spin
If we have 7 electrons, how will they fill in the below orbitals?
Energy
2px
2py
2pz 2s 1s

31. Electron Configuration and Orbital Notation
Pauli Exclusion Principle – an orbital can contain no more than 2 electrons. Electrons in the same orbital must have different spins.
If we have 8 electrons, how will they be arranged?
Energy
2px
2py
2pz 2s 1s

33. Apartment Analogy Atom is the building Floors are energy levels
Rooms are orbitals
Only two people per room

34. Orbital Diagrams Draw each orbital as a box.
Each electron is represented using an arrow.
Up arrows – clockwise spin
Down arrows – counter-clockwise spin
Determine the total number of electrons involved.
Start with the lowest energy level (1s) and start filling in the boxes according the rules we just learned.

35. Orbital Diagram 2s 1s 3s 4s 2p 3p 4p 3d
Energy

36. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
Chapter 5

37. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
The first to electrons go into the 1s orbital
Notice the opposite spins
only 13 more
Chapter 5

38. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
The next electrons go into the 2s orbital
only 11 more
Chapter 5

39. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
The next electrons go into the 2p orbital
only 5 more
Chapter 5

40. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
The next electrons go into the 3s orbital
only 3 more
Chapter 5

41. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s
The last three electrons go into the 3p orbitals.
They each go into separate shapes
3 unpaired electrons
1s22s22p63s23p3
Chapter 5

42. Orbital Diagrams
Orbital diagrams are used to show placement of electrons in orbitals.
Need to follow three rules (Aufbau, Pauli, Hund’s) to complete diagrams Li Be B C N Ne Na

43. Orbitals and Energy Levels
Principal
Sublevels
Orbitals
Energy Level
n = 1 1s 1s
(one)
n = 2 2s 2p , 2s
(one) + 2p
(three)
n = 3 3s , 3p , 3d 3s
(one) + 3p
(three) + 3d
(five)
n = 4 4s , 4p , 4d , 4f 4s
(one) + 4p
(three) + 4d
(five) + 4f
(seven)
Chapter 5

44. Summary
shapes
Max electrons
Starts at energy level s p d f
Chapter 5

45. Orbitals and Energy Levels4p 4d 4f
n = 4
and so on.... 3s 3p 3d
n = 3
Increasing energy 2s 2p
n = 2 1s
n = 1

46. Electron Configuration
Let’s determine the electron configuration for Phosphorus
Need to account for 15 electrons
Chapter 5

47. Writing Electron Configuration
Determine the total number of electrons.
Write the principle energy level number as a coefficient, the letter for the subshell, and an exponent to represent the number of electrons in the subshell.
He: 1s2

48. The Kernel (Noble Gas) Notation
Determine the total number of electrons
Find the previous noble gas and put its symbol in brackets
Write the configuration from that noble gas forward as usual

49. Writing electron configurations
Examples
O 1s2 2s2 2p4
Ti 1s2 2s2 2p6 3s2 3p6 3d2 4s2
Br 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5
Core format
O [He] 2s2 2p4
Ti [Ar] 3d2 4s2
Br [Ar] 3d10 4s2 4p5
Chapter 5