1. Electron Configuration
Chapter 4

2. Light Today scientists believe there are 2 properties of light: waves
particles
Electromagnetic waves consist of both electric and magnetic fields.

3. Electromagnetic Waves

4. Wave Activity
Use the rope to devise an explanation of the following to be given to a person without sight:
Amplitude
Wavelength
Frequency
Crest
Trough
Relate wavelength and frequency

5. Waves All waves have 4 characteristics Amplitude –
the height of the wave
Wavelength –
the distance between crests

6. Waves - 4 characteristics
Frequency –
tells how fast a wave oscillates up and down.
Speed –
Light moves at a constant speed of 3.00x108m/sec.

7. Waves
Explain the relationship between frequency and wavelength.

8. Electromagnetic Spectrum
Visible Light – each color has a different wavelength and therefore a different color.
Roy G. Biv
Violet has the shortest wavelength (highest frequency)
Red has the longest wavelength (lowest frequency)

9. Wavelengths

10. Quantum Theory

11. Quantum Theory Questions about electromagnetic radiation:
Why does a heated piece of metal emit different wavelengths at different temperatures?
Why do some metals emit different colors when heated?

12. Planck

13. Planck’s Theory
Energy emitted or absorbed by an object is restricted to certain quantities.
Each “piece” of this energy is a quantum.
Energy absorbed or emitted by atoms are quantized. Their values are restricted to certain quantities.

14. Photoelectric Effect

15. Photoelectric Effect (Einstein)
When light shines on a metal – electrons are ejected from the surface.
Light consists of quanta of energy that act like tiny particles of light. These are called photons.
When a photon strikes the metal it transfers its energy to an electron in the metal.
Light with a greater frequency has more energy per photon.

16. Atomic Emission Spectra or Line Spectrum
When different elements are heated to high temperatures, they emit a specific color (wavelength) of light.
The atoms absorb energy and then release it as a specific wavelength.

17. Line Spectra

18. The Bohr Model
Bohr proposed that each electron has a certain quantum of energy called the ground state of energy.
The electrons at ground state are in the lowest energy level.
When electrons absorb radiant energy they are excited and move to a higher energy level.
When they move back to the ground state level they emit radiation (color)

20. Lab
Flame Tests and Emissions

21. Heisenberg’s Uncertainty Principle
The position and the momentum of a moving object cannot simultaneously be measured and known exactly.
You can’t know where a particle is and how it is moving at any one moment.
So, you can’t really predict where a particle will be in the future.

22. Lab
Quantum Leap Lab

23. Electron Density
Electron Density = the density of the electron cloud – where electrons are most likely to be found
Atomic Orbital (“electron cloud”)
Region around the nucleus of an atom where an electron of given energy is likely to be found 23

24. Quantum Numbers Four Quantum Numbers:
These numbers identify the location of the electron.
UPPER LEVEL 24

25. Quantum Numbers Principal Quantum Number ( n )
n is the symbol for the principle energy levels (n=1, n=2, n=3, n=4, etc.)
Energy of the electron increases as “n” increases
The size of the energy level increases as “n” increases 1s 2s 3s
Courtesy Christy Johannesson 25

26. Quantum Numbers
2. Each principle energy level is divided into Sublevels
The number of sublevels that a principal energy level has is the number of that principal energy level.
So, energy level 1 has one sublevel, level 2 has two sublevels, etc.
Sublevels are labeled: s, p, d, f
Principal energy level 1 = s sublevel
Principal energy level 2 = s and p sublevels
Principal energy level 3 = s, p and d sublevels
Principal energy level 4 = s, p, d and f sublevels 26

27. Quantum Numbers 3. Orbitals
The area where there is the probability of finding an electron.
The s sublevels have one orbital.
The orbitals increase in size and energy as the sublevel increases in number. 1s 2s 3s 27

28. Shapes of Orbitals in the s,p,d,f sublevels
s = 1 orbital
p = 3 orbitals
d = 5 orbitals
f = 7 orbitals

29. Energy Levels
As the value of n increases
the electron energy increases.
electrons also spend more time farther from the nucleus making the orbital larger

30. Quantum Numbers 4. Spin of the electron Electron spin  +½ or -½
An orbital can hold 2 electrons that spin in opposite directions.
The spin creates a magnetic field (N and S poles)
Courtesy Christy Johannesson 30

31. Each orbital can hold a maximum of 2 electrons:
Maximum Number of Electrons
In Each Sublevel
Maximum Number
Sublevel Number of Orbitals of Electrons s p d f

32. Electron Configuration
Electron configuration is the distribution of the electrons among the orbitals of an atom.
Describes:
Where the electrons are found
What energies the electrons possess
When electrons are in the lowest energy orbitals possible they are at ground state
Ground state is the most stable state of the atom.
The electron configuration is determined by a set of principles.

33. Principles for Determining Electron Configuration
Aufbau Principle: Electrons are added one at a time to the lowest
energy orbitals available until all the electrons of the atom
have been accounted for.
Pauli Exclusion Principle: An orbital can hold a maximum of two electrons.
To occupy the same orbital, two electrons must spin in opposite
directions.
Hund’s Rule: Electrons occupy equal-energy orbitals so that a maximum
number of unpaired electrons results. Each orbital in a sub-level gets
electron first and then they each get a second electron.
*Aufbau is German for “building up” 33

34. Determining Electron Configuration
Orbital Diagrams

35. 1s2 2s2 2p4 O Notation 1s 2s 2p 8e- O Orbital Diagram 8
Notation
Orbital Diagram 1s 2s 2p O
8e-
Electron Configuration
1s2 2s2 2p4
Courtesy Christy Johannesson

36. How the sublevels are filled:
Because there can only be 8 valence electrons at any given time – the 4s level must be filled before the 3d, the 5s before the 4d, etc.

37. Aubfau Principle
Electrons are added one at a time to the lowest energy orbitals available until all of the electrons are accounted for. 2s 2p

38. Pauli Exclusion Principle
Max of 2 electrons per orbital
Must have opposite spin

39. Hund’s Rule WRONG RIGHT
Within a sublevel, place one electron per orbital before pairing them.
“Empty Bus Seat Rule”
WRONG
RIGHT
Courtesy Christy Johannesson

40. Practice Problems Al P Cl