1. Modern Atomic Theory Notes

2. Electromagnetic radiation – energy that travels through space as waves.
Waves have three primary characteristics:
Wavelength (- lambda) – distance between two consecutive peaks or troughs in a wave. Unit = meter
Frequency (  = nu) – indicates how many waves pass a given point per second. Unit = Hertz (Hz)
Speed – velocity (c = speed of light = 3 x 108 m/sec) - indicates how fast a given peak moves in a unit of time
c = 

3. Tell me what you know… Which wave has the greatest wavelength?
Which wave has the greatest frequency?
Which wave has the greatest speed?

4. Electromagnetic radiation (light) is divided into various classes according to wavelength.

5. Tell me what you know… Which color has the greatest wavelength?
Which color has the shortest wavelength?

6. Wave- Particle Theory – Light as waves & Light as photons
Photon/quantum – packet of energy OR a “particle” of electromagnetic radiation

7. Change in Energy of a photon = (Planck’s Constant) x (frequency)
Energy - (E – change in energy) – Unit Joules (J)
Planck’s Constant –
(h = x J * s)
Ephoton = h
Change in Energy of a photon = (Planck’s Constant) x (frequency)
c =  & Ephoton = h →
Ephoton = hc 
Ex: What is the wavelength of light with a frequency of 6.5 x 1014 Hz? What is the change in Energy of the photon?
E = hc 
ΔE = 4.3 x J
Given
= 6.5 x 1014 Hz
 = ? ΔE = ?
= c 
= 3 x 108 m/sec
6.5 x 1014 Hz
= (6.626 x J x s)(3 x 108 m/s)
4.6 x 10-7 m
λ = 4.6 x 10-7 m

8. Tell me what you know…
What does ΔE tell us about a photon?

9. Excited State – atom with excess energy
Ground State – lowest possible energy state
Wavelengths of light carry different amounts of energy per photon
Only certain types of photons are produced (see only certain colors)
Quantized – only certain energy levels (and therefore colors) are allowed

10. Emission and Absorption Spectra
Emission Spectrum – bright lines on a dark background. Produced as excited electrons return to a ground state – as in flame tests.
Absorption Spectrum – dark lines in a continuous spectrum. Produced as electrons absorb energy to move into an excited state, only certain allowable transitions can be made. Energy absorbed corresponds to the increase in potential energy needed to move the electron into allowed higher energy levels. The frequencies absorbed by each substance are unique.
Nucleus
Nucleus

11. Tell me what you know…
How can we tell elements apart using emission spectra?

12. Bohr Model – suggested that electrons move around the nucleus in circular orbits
Only Correct for Hydrogen
Wave Mechanical Model – Described by orbitals gives no information about when the electron occupies a certain point in space or how it moves *aka – Heisenberg's Uncertainty Principle

13. Parts of the Wave Mechanical Model
1. Principle Energy Level (n) – energy level designated by numbers 1-7.
-called principle quantum numbers 1 2 3 4 5 6 7
2. Sublevel – exist within each principle energy level
-the energy within an energy level is slightly different
-each electron in a given sublevel has the same energy
-lowest sublevel = s, then p,
then d, then f s p d f

14. Tell me what you know…
Write the sublevels in order of highest energy to lowest energy.

15. Parts of the Wave Mechanical Model cont.
3. Orbital – region within a sublevel or energy level where electrons can be found
s sublevel – 1 orbital
p sublevel – 3 orbitals
d sublevel – 5 orbitals
f sublevel – 7 orbitals
- ** No more than two electrons can occupy an orbital**
-an orbital can be empty, half-filled, filled

16. Tell me what you know…
How many total electrons are in each sublevel (s, p, d, & f)?

17. Sulfur = 1s2 2s2 2p6 3s2 3p4 Cd = 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
Electron Configuration – arrangement of the electrons among the various orbitals of the atom Ex: 1s22s22p6 = Neon
Sulfur =
1s2
2s2
2p6
3s2
3p4
Cd =
1s2
2s2
2p6
3s2
3p6
4s2
3d10
4p6
5s2
4d10
Na =
1s2
2s2
2p6
3s1 Ne Na

18. Principle Energy Level
Summary
Principle Energy Level
# of sublevels
# of orbitals present
s p d f
Total # of orbitals
Maximum # of electrons 1 1 1 2 2 2 4 8 3 3 9 18 4 4 16 32

19. Shapes of orbitals All s orbitals are spherical
as the principle energy level
increases the diameter increases.
All p orbitals are dumbbell or figure-8 shaped – all have the same size and shape within an energy level

20. 4 of the d orbitals are 4-leaf clover shaped and the last is a figure-8 with a donut – all have the same size and shape within an energy level

21. f orbitals are complicated!!!!!

22. Electron Spin Spin – motion that resembles earth rotating
on its axis– clockwise or counterclockwise
Pauli Exclusion Principle – two electrons in the same orbital must have opposite spins
Hund’s Rule – All orbitals within a sublevel must contain at least one electron before any orbital can have two
Orbital Diagram – describes the placement of electrons in orbitals
use arrows to represent electrons with spin
line represents orbital (s=1, p=3, d=5, f=7)
____ full ____ half-full ____ empty

23. Orbital Diagrams Neon = 1s__ 2s__ 2p__ __ __ Ex: Carbon = 1s__ 2s__
Zinc =
1s__
2s__
2p__ __ __
3s__
3p__ __ __
4s__
3d__ __ __ __ __
Gallium =
1s__
2s__
2p__ __ __
3s__
3p__ __ __
4s__
3d__ __ __ __ __
4p__ __ __

24. Tell me what you know… Summarize each in four words or less: Spin
Pauli Exclusion Principle
Hund’s Rule

25. Aufbau Order
Aufbau Order – Tool to predict the order in which sublevels will fill
OR use order on Periodic Table

26. Core Electrons – innermost electrons – not involved in bonding
Noble Gas Configuration – Shorthand configuration that substitutes a noble gas for electrons
Ex:
Valence Electrons – Electrons in the outermost (highest) principle energy level in an atom, electrons use in bonding
Core Electrons – innermost electrons – not involved in bonding
Valence Configuration – shows just the valence electrons
Na = 1s22s22p63s or [Ne]3s1
Sn = 1s22s22p63s23p64s23d104p65s24d105p2 or [Kr]5s24d105p2
Na = 3s1 3rd Shell/1valence electron
Sn = 5s25p2 5th Shell/4 valence electrons