1. Unit 3 – The Electron
Chapter 5
Test:

2. Review of Atomic Theory
Bohr determined that the e- travel around the nucleus according to energy
Electrons must have energy to keep them away for the nucleus (opposites attract!)
Closer to nucleus the lower the energy
Energies are also observed with the speed the electrons orbit the nucleus

3. Atomic Orbitals
Orbitals - describes the as clouds where the e- will be found 90% of the time (Quantum Mechanical Model)
3-D region of probability
Electron cloud gives volume of the atom
Each atomic orbital has its own general size – not defined as Bohr suggested

4. Principal Quantum Numbers
Model assigns principal quantum numbers (n) that indicate relative size & energies of the orbitals
On periodic table each row is an energy level
n specifies the energy levels (principal energy levels)
Lowest energy level assigned n =1
When electron is in that level it is at its’ ground state

5. Description of Spectrum
Energy Levels
Principal energy levels contain energy sublevels
There are 4 energy sublevels – s, p, d, and f and are given by the shapes the orbitals make.
Each sublevel contains a different number of orbitals
Each orbital can only contain 2 electrons
Sublevel
Description of Spectrum
# orbitals
Max # e- s
Sharp 1 2 p
Principal 3 6 d
Diffuse 5 10 f
Fundamental 7 14

6. Organization of Energy Levels

7. Orbital Shapes
Figures 5-15 & 5-16 on page 133 of Text Book

8. Orbital Shapes
Image from

9. Locating Electrons
We have two ways to show where the electrons are found in the atom
Orbital filling diagrams
Electron configurations

10. Orbital Filling Diagrams
Show how the electrons fill into the orbitals
Each box or circle represents an orbital which can hold a max of 2 electrons
Electrons must fill all of one energy sub-level before starting into another
Electrons are notated with an arrow (up or down)
Up arrows must fill the boxes first then double up with the down arrows
Arrows represent the spin of the electrons

11. Orbital Filling Diagrams
Figure 5-17 on page 135 of Text Book

12. Orbital Filling Diagrams
The three p orbitals fill in the order shown below:
The number of arrows must match the number of electrons contained in the atom
Example: Carbon has six electrons
Page 136 in Text Book

13. Electron Configuration
Shorthand method for describing the arrangement of electrons
Composed of the principal energy level followed by the energy sublevel and includes a superscript with the # of electrons in the sublevel
# electrons in sublevel
He 1s2
Energy Level
Sublevel

14. Putting it all together
Neon Atom
Electron Configuration: 1s22s22p6
Orbital Filling Diagram
Orbital image:

15. Electron Configuration Shorthand
Give the symbol of the noble gas in the previous energy level in brackets
Give the configuration for the remaining energy level
Example:
Sulfur 1s22s22p63s23p4
[Ne]3s23p4

16. Valence Electrons
Valence electrons: found in the outermost energy level
Electrons used for bonding
Represented visually in Lewis-Dot Structures
Example: Carbon 1s2 2s2 2p2
Add up the number of electrons (superscripts) in the highest energy level
So, carbon has 4 valence electrons

17. Lewis-Dot Structures
Element’s symbol represents the nucleus and inner-level electrons
Dots represent the valence electrons
Dots are placed one at a time on the four sides of the symbol then paired until all valence electrons are used.
Maximum of 8 electrons will be around the symbol
d sublevel electrons are not valence electrons – they are in a lower energy level!

18. Lewis-Dot Structures
Page 140 in Text Book

19. Ions Atoms that have gained or lost electrons
The word atom implies that it is neutral!
Denoted by a superscript charge (sign and number) to the right of the element symbol
Examples: Cl- and Mg2+

20. Ions Cation – positive ion (Cat Ions are Pawsitive)
Loses electrons to become more positive
Example: Be 1s22s2 → Be2+ 1s2
Anion – negative ion
Gains electrons to become more negative
Example: F 1s22s22p5 → F- 1s22s22p6
What do you notice about the ions’ electron configurations?

21. Ions- Practice
Determine the # of protons, neutrons, and electrons and name as a cation or anion
K1+
Cl1-
O2-
Mg2+

22. Lewis Dot Diagrams - Ions
Metals lose electrons to form cations
Nonmetals gain electrons to form anions
Set-up diagram using the ions electron configuration, then place brackets around the diagram with a superscript of the charge
Example: O2- 1s22s22p6 [ O ]2-

23. Light
Electromagnetic Radiation (light) is a form of energy with a wavelike nature
Visible light is only a small portion of the electromagnetic spectrum
Wave model does not explain all of light’s behavior
Use a dual wave-particle model
Explains why chemicals give off certain colors of light when heated in a flame

24. How Light is Emitted ) ) ) ) ) ) ) 1 2 3 4 5 6 7 energy levels
As energy is absorbed (heat gained) the electrons move from their ground state to an excited state (higher energy level)
final
position
initial
position
nucleus
energy levels
) ) ) ) ) ) )
Absorbs energy
ground state
excited state

25. How Light is Emitted ) ) ) ) ) ) ) 1 2 3 4 5 6 7 energy levels
Electrons’ unstable in excited state
Returns to ground state by releasing energy (light quantum)
Color of light determined by the # of energy levels moved & amount of energy electron had
initial
position
final
position
nucleus
energy levels
) ) ) ) ) ) )
excited state
ground state
Releases energy – gives of light

26. Visible Light B lue G reen Y ellow R ed ← (low energy)
(high energy)  V iolet
I ndigo
B lue
G reen
Y ellow
O range
R ed ← (low energy)

27. Electromagnetic Spectrum
Figure 5-5 on Page 120 in Text Book

28. Emission Of Light
Max Planck described the emission spectrum of objects that were heated, from this we get the following terms:
Quantum – minimum amount of energy that can be gained or lost by an atom (can be referred to as a packet of energy)
Photon – packet of light energy (light quantum), has wave & particle properties

29. Spectra
Emission Spectra - Series of colored lines used to identify an element (each element has different spectrum)
Shows all the wavelengths of light that are emitted
Spectroscope – instrument used to see the emission spectra
Absorption Spectra – Opposite of emission spectra
Shows all the wavelengths of light that are absorbed

30. Spectra
Emission Spectra
Absorption Spectrum

31. Photoelectric Effect
Phenomenon where electrons are emitted from a metal’s surface when light of a certain frequency shines on the surface
Solar panels use this to generate electricity (solar calculators too!)
Page 123 in Text Book

32. End of Unit 3 Notes
Study for Test on