1. Chapter 5 Introductory Assignment
1. Describe the properties of a wave of light. Draw a wave and label its properties.
2. Draw the electromagnetic spectrum. Correctly color the portion of visible light with map pencils. Explain the EM spectrum.
3. What formula and what constant will be used to calculate wavelength of an EM wave?
4. What formula and what constant will be used to calculate the energy of a photon?
5. Define and explain what a photon is.

2. UNIT 3: Electrons in Atoms
-Light and Quantized Energy
-Quantum Theory and the Atom
- Electron Configuration

3. The Development of a New Atomic Model

4. Waves Wavelength () - length of one complete wave
Frequency (f) - # of waves that pass a point during a certain time period
hertz (Hz) = 1/s
Amplitude (A) - distance from the origin to the trough or crest

5.  A A  Waves crest greater amplitude (intensity) origin trough
greater frequency
(color)

6. EM Spectrum
HIGH
ENERGY
LOW
ENERGY

7. EM Spectrum HIGH ENERGY LOW ENERGY R O Y G. B I V red orange yellow
green
blue
indigo
violet

8. Wavelength & Frequency
Frequency & wavelength are inversely proportional
c = 
c: speed of light (3.00  108 m/s)
: wavelength (m, nm, etc.)
: frequency (Hz, s-1)

9. Wavelength & Frequency Example
EX:What is the wavelength of a photon whose frequency is 3.90 X 104 Hz
GIVEN:
 = 3.90 X 104Hz =?
c = 3.00  108 m/s 9

10. Wavelength & Frequency Example
EX: Find the frequency of a photon with a wavelength of 434 nm.
GIVEN:
 = ?
 = 434 nm = 4.34  10-7 m
c = 3.00  108 m/s
WORK:
 = c 
 = 3.00  108 m/s
4.34  10-7 m
 = 6.91  1014 s-1

11. E = h Energy & Frequency E: energy (Joules, J)
Energy & wavelength are directly proportional
E = h
E: energy (Joules, J)
h: Planck’s constant, 6.63x10-34 J•s
: frequency (Hz)

12. Energy & Frequency Example
EX: Determine the wavelength of a photon whose frequency is 3.55x1017 Hz.
GIVEN:
E = ?
v = 3.55x1017 Hz = 3.55x1017 s-1
h = 6.63x10-34 J•s
WORK:
E=hν
E=2.3510-16 J

13. Energy & Frequency Example
EX: Find the energy of a red photon with a frequency of 4.57  1014 Hz.
GIVEN:
E = ?
 = 4.57  1014 Hz
h = 6.63  J·s
WORK:
E = h
E = (6.63  J·s) (4.57  1014 Hz)
E = 3.03  J

14. Some vocab… Ground state-lowest energy state of an atom
Excited state-a state in which an atom has a higher potential energy that it has in its ground state
So, how do atoms transition between their ground state and their excited state???

15. Bohr Model
e- exist only in orbits with specific amounts of energy called energy levels
Therefore…
e- can only gain or lose certain amounts of energy because they can exist only at certain energy levels
only certain photons are produced (because only certain amounts of energy is released)

16. Line-Emission Spectrum
excited state
ENERGY IN
PHOTON OUT
ground state
Emission spectrum of H2 gas

17. Line Emission Spectra
Classical theory-atoms would be excited by any amount of energy added to them, and should give off a continuous spectrum of EM radiation.
Attempts to explain further developed the quantum theory of the atom, and led to the Bohr model of the hydrogen atom

18. Bohr Model Energy of photon depends on the difference in energy levels6
Energy of photon depends on the difference in energy levels
Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom 5 4 3 2 1

19. Other Elements
Each element has a unique bright-line emission spectrum.
“Atomic Fingerprint”
Helium
Bohr’s calculations only worked for hydrogen! 

20. Electron Configuration

21. Where are electrons Found???
AKA “electron cloud”
Region in space where there is 90% probability of finding an e-
Radial Distribution Curve
Orbital

22. Atoms are three dimensional…
Not “flat” like Bohr’s model with electrons circling the nucleus like the planets circle the sun 2s
2pz
2py
2px 22

23. 11/6 TEST TODAY Turn in HW Pickup paper from side table
You need a calculator and your colored periodic table
We will review for half the period and then test

24. Electron Configuration
Specifies the “address” of each electron in an atom
UPPER LEVEL

25. General Rules Pauli Exclusion Principle
Each orbital can hold TWO electrons with opposite spins.

26. General Rules Aufbau Principle
Electrons fill the lowest energy orbitals first.
“Lazy Tenant Rule”

27. General Rules WRONG RIGHT Hund’s Rule
Within a sublevel, place one e- per orbital before pairing them.
“Empty Bus Seat Rule”
WRONG
RIGHT

28. Representing e- configurations
3 ways to represent e- configurations
1. Orbital notation
ex:      
1s s p s
2. Electron configuration notation
ex: 1s2 2s2 2p6 3s2
3. Noble gas notation
ex: [NE] 3s2

29. O Orbital notation 8e- 1s 2s 2p
Orbitals are represented by lines or boxes that are labeled with the energy level and sublevel
e- are represented as arrows
↑=+ ½ e- spin
↓ =- ½ e- spin
**Remember Hund’s Rule!!** O
8e- 1s 2s 2p

30. Periodic Patterns s p d (n-1) f (n-2) 1 2 3 4 5 6 7 6 7
© 1998 by Harcourt Brace & Company

31. Orbital notation practice
Draw the orbital notation for He Li Na Si Ne Ca

32. e- Configuration Notation
Energy levels and sublevels are represented by numbers and letters
Electrons are represented by superscripted numbers 1s 2 2s 2 2p 4 O
8e-

33. e- Configuration Notation practice
Write the e- configuration notation for the following elements: Be Ge S Cd

34. Some vocab…
Valence electrons-electrons that are found in the highest occupied energy level
Highest occupied energy level-electron-containing main energy level with the highest principal quantum number
Inner shell or core electrons-electrons that are not in the highest occupied energy level

35. Noble Gas Notation “Abbreviated” form of e- configuration notation
Use the noble gas (Gp. 18) that immediately precedes the element of interest as shorthand for the majority of the notation
Write the e- configuration notation for Ne and compare it to S. What do you notice?

36. S 16e- 1s2 2s2 2p6 3s2 3p4 Ne 10e- 1s2 2s2 2p6 S 16e- [Ne] 3s2 3p4
Noble Gas Notation
S 16e-
1s2
2s2
2p6
3s2
3p4
Ne 10e-
1s2
2s2
2p6
Only difference is the valence electrons in sulfur. Use [Ne] as an abbreviation for all parts that are the same!
S 16e- [Ne] 3s2 3p4

37. Noble Gas Notation practice
Write the noble gas notation for the following elements: Cl B Sr I

38. Periodic Patterns Period # A/B Group # Column within sublevel block
energy level (subtract for d & f)
A/B Group #
total # of valence e-
Column within sublevel block
# of e- in sublevel

39. 1s1 Periodic Patterns 1st column of s-block 1st Period s-block
Example - Hydrogen
1s1
1st column of s-block
1st Period
s-block

40. Periodic Patterns p s d (n-1) f (n-2) Noble Gas Notation
Core e-: Go up one row and over to the Noble Gas.
Valence e-: On the next row, fill in the # of e- in each sublevel. s
d (n-1)
f (n-2) p

41. Periodic Patterns
Example - Germanium
[Ar]
4s2
3d10
4p2

42. Stability Full energy level Full sublevel (s, p, d, f)
Half-full sublevel

43. Exceptions Copper Copper gains stability with a full d-sublevel.
EXPECT: [Ar] 4s2 3d9
ACTUALLY: [Ar] 4s1 3d10
Copper gains stability with a full d-sublevel.

44. Exceptions
Chromium
EXPECT: [Ar] 4s2 3d4
ACTUALLY: [Ar] 4s1 3d5
Chromium gains stability with a half-full d-sublevel.

45. Configuration of Ions 1+ 2+ 3+ NA 3- 2- 1- Ion Formation
Atoms gain or lose electrons to become more stable.
Isoelectronic with the Noble Gases. 1+ 2+ 3+ NA 3- 2- 1-

46. Configuration of Ions O2- 10e- [He] 2s2 2p6
Writing Ion Electron Configuration
Write the e- config for the closest Noble Gas
EX: Oxygen ion  O2-  Ne
O e [He] 2s2 2p6