1. Happy MONDAY, Chemistry! 
What do you need for class?
Pencil
Calculator
Unit 2 Packet
Composition Notebook
Agenda:
The Atom and Light
Electromagnetic Spectrum
Light Equations
Notes on Light
Learning Targets
I CAN describe how light is produced.
I CAN describe the electromagnetic spectrum.
I CAN explain the relationships between energy, frequency, and wavelength of a light wave.
I CAN use the light equation to calculate wavelength, frequency, and energy of a light wave.

2. ELECTROMAGNETIC RADIATION

3. Electromagnetic Radiation
Wavelength (l) = the distance between wave peaks

4. What is the difference between these two waves?
So is radiation harmful?
Which wave would be more harmful to you?

5. Look at the Electromagnetic Spectrum handout
Find Visible light.
Which direction is increasing wavelength?
Which direction is increasing frequency?
Which direction is increasing Energy?

6. Electromagnetic Spectrum
Long wavelength  Low frequency
Short wavelength  High frequency
increasing frequency
increasing wavelength

7. E = energy c = speed of light
Light Equations
We will use two light equations, and one equation that is the combination of the two. You must know how to manipulate these algebraically to solve for each variable.
λ = wavelength
ν = frequency
E = energy c = speed of light
h = Planck’s constant

8. Light Equations

9. How do calculations relate to what we saw in the light lab?
What’s this thing called again?
How do calculations relate to what we saw in the light lab?

10. What are you seeing when you look through the spectroscope?
What is the prism doing?

11. Spectrum of Hydrogen Gas

12. What are you seeing when you look through the spectroscope?
What is the prism doing?

14. Electricity and Gases

15. What you should have seen in your lab.
Emission spectra!

17. What causes different colors of light?

18. Electromagnetic Radiation
wavelength
Visible light
Ultaviolet radiation
Amplitude
Node

19. Atomic Line Emission Spectra and Niels Bohr
Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the LINE EMISSION SPECTRA of excited atoms.
Niels Bohr
( )

20. Line Emission Spectra of Excited Atoms
Excited atoms emit light of only certain wavelengths
The wavelengths of emitted light depend on the element.

21. What happens when electrons get excited?
One view of atomic structure in early 20th century was that an electron (e-) traveled about the nucleus in an orbit.

22. Remember Niels Bohr?
Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the LINE EMISSION SPECTRA of excited atoms.
THIS MODEL ONLY WORKS FOR HYDROGEN.
Niels Bohr
( )

23. Quantum or Wave Mechanics
Schrodinger applied idea of e- behaving as a wave to the problem of electrons in atoms.
He developed the WAVE EQUATION
Solution gives set of math expressions called WAVE FUNCTIONS, 
Each describes an allowed energy state of an electron.
E. Schrodinger

24. Heisenberg Uncertainty Principle
Problem of defining nature of electrons in atoms solved by W. Heisenberg.
Cannot simultaneously define the position and momentum of an electron.
We define electron energy exactly but accept limitation that we do not know exact position.
W. Heisenberg

25. Interactives: Remember This?

26. Arrangement of Electrons in Atoms
Electrons in atoms are according to their
ENERGY LEVELS (Size)
SUBLEVELS (Shape)
ORBITALS (Orientation)

27. Energy Levels (Size)
n = 1
n = 2
n = 3
n = 4

28. Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

29. Sublevels (Shapes)
The most probable area to find these electrons takes on a shape.
So far, we have 4 shapes. They are named s, p, d, and f.

30. Sublevels (Shapes)
s orbital
p orbital
d orbital

31. s Orbitals (Orientation)
No more than 2 e- assigned to an orbital.
One spins clockwise, one spins counterclockwise

32. p Orbitals (Orientation)
The three p orbitals lie 90o apart in space
If no more than 2 e- can fit into each orbital, how many electrons are in the p sublevel?

33. d Orbitals (Orientation)
If no more than 2 e- can fit into each orbital, how many electrons are in the d sublevel?

34. f Orbitals (Orientation)
If no more than 2 e- can fit into each orbital, how many electrons are in the f sublevel?

35. Remember: A maximum of two electrons can be placed in an orbital.
How many electrons can be in a sublevel?
Remember: A maximum of two electrons can be placed in an orbital.
Shape of
sublevel s p d f
Number of
orbitals
Number of electrons

36. Where are the electrons at any given moment?
To be honest, we don’t exactly know.
Remember Heisenburg said, the more we know about an electron’s location, the less we know about its speed. So, by the time we have isolated an electron, it moved anyway.
BUT, based on probabilities, we can assume that electrons are in certain locations.

37. Locations of Electrons:
The Rules!
The Aufbau principle states that electrons must fill the lowest energy orbitals first.
 Hund's rule states every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied. (Seats on a bus)
The Pauli Exclusion principle states that no more than two electrons can occupy an orbital, and they must do so with opposite spins.

38. Electron Configurations
A list of all the electrons in an atom
Must go in order (Aufbau principle)
2 electrons per orbital, maximum (Hund’s Rule)
We need electron configurations so that we can determine the number of electrons in the outermost energy level. These are called valence electrons.
The number of valence electrons determines the propertied of an atom, including the type of bonds it can make to in order to make a molecule

39. How will I ever remember all of the rules?!
Diagonal Cheat Sheet!
But Ms. Kovach,
How will I ever remember all of the rules?!
What do they mean!??!
Must be able to write it for the test!
The diagonal rule is a memory device that helps you remember the order of the filling of the orbitals from lowest energy to highest energy
_____________________ states that electrons fill from the lowest possible energy to the highest energy

40. Diagonal Cheat Sheet! 1 2 3 4 5 6 7 s s 2p s 3p 3d s 4p 4d 4f
Steps:
Write the energy levels top to bottom.
Write the orbitals in s, p, d, f order. Write the same number of orbitals as the energy level.
Draw diagonal lines from the top right to the bottom left.
To get the correct order, follow the arrows! 1 2 3 4 5 6 7 s
s 2p
s 3p 3d
s 4p 4d 4f
By this point, we are past the current periodic table so we can stop.
s 5p 5d 5f 5g?
s 6p 6d 6f 6g? 6h?
s 7p 7d 7f 7g? 7h? 7i?

41. Electron Configurations
2p4
Number of electrons in the sublevel
Energy Level
Sublevel
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.

42. Let’s Try It!
Write the electron configuration for the following elements: H Li N Ne K Zn Pb

43. Shorthand Notation A way of abbreviating long electron configurations
Since we are only concerned about the outermost electrons, we can skip to places we know are completely full (noble gases), and then finish the configuration

44. Shorthand Notation
Step 1: Find the PREVIOUS noble gas to the atom. Write the noble gas in brackets [ ].
Step 2: Find where to resume by finding the next energy level.
Step 3: Resume the configuration until it’s finished.

45. Let’s Try It!
Write the shorthand electron configuration for the following elements: H Li N Ne K Zn Pb

46. Orbitals and the Periodic Table
But Ms. Kovach,
How am I going to remember the diamond cheat sheet
…..and the shorthand shortcut
….and EVERYTHING ELSE?
Orbitals and the Periodic Table
Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental) s d p f

47. Let’s Try It!
Write the electron configuration for the following elements: H Li N Ne K Zn Pb

48. Electron Configuration: Chlorine
Longhand is 1s2 2s2 2p6 3s2 3p5
You can abbreviate the first 10 electrons with a noble gas, Neon. [Ne] replaces 1s2 2s2 2p6
The next energy level after Neon is 3
So you start at level 3 on the diagonal rule (all levels start with s) and finish the configuration by adding 7 more electrons to bring the total to 17
[Ne] 3s2 3p5

49. Practice Shorthand Notation
Write the shorthand notation for each of the following atoms: Cl K Ca I Bi

50. Orbital Diagrams Graphical representation of an electron configuration
One arrow represents one electron
Shows spin and which orbital within a sublevel
Same rules as before (Aufbau principle, Pauli Exclusion principle, and Hund’s Rule).

51. Lithium
Group 1A
Atomic number = 3
1s22s1  3 total electrons

52. Carbon Group 4A Atomic number = 6 1s2 2s2 2p2  6 total e-
Here we see for the first time HUND’S RULE. When placing electrons in a set of orbitals having the same energy, we place them singly as long as possible.

53. Draw these orbital diagrams!
Oxygen (O)
Chromium (Cr)
Bromine (Br)

54. Interactives: Remember This?