1. Electronic Structure

2. Wave Nature of Light Electromagnetic Radiation
Gamma Rays, Visible Light
Moves Through Vacuum at 3.00x108m/s (c)
Wavelength (λ, m) = distance between successive peaks or troughs
Frequency (f, s-1) = how often a wave passes through a particular point
c = f · λ, or v · λ

3. Wavelength Practice
The brilliant red colors seen in fireworks are due to the emission of light with wavelengths around 650nm when strontium salts such as Sr(NO3)2 and SrCO3 are heated. Calculate the frequency of red light of wavelength 6.50x102nm.
A FM radio station broadcasts electromagnetic radiation at a frequency of 103.4MHz. Calculate the wavelength of this radiation. (1MHz=106s-1)

4. Quantize Energy and Photons
Wave model explains much of the behavior of light but not all:
Black body radiation – Emission of light from hot objects
Photoelectric Effect – Emission of electrons from metal surfaces
Emission Spectra – Emission of light from excited atoms

5. Black Body Radiation

6. Hot Objects and Quantization of Energy
When objects are heated they emit light
Red hot (cooler) → white hot (hotter)
Light only emitted at certain wavelengths
Max Planck declared that energy can only be emitted or absorbed in packets (quanta, photon)
E= h · v
h = Planck's Constant 6.63x10-34J·s
Energy emitted at whole number multiples of hv
Think walking up and down stairs
Why don't we notice this? Why does energy seem to flow continuously for us?

7. Photoelectric Effect and Photons
When light shines on an object, electrons are emitted
Light has to have a specific energy, frequency and wavelength in order for e- to be emitted
Photons are absorbed
Too little energy – nothing happens
Just right amount – electrons are emitted
A little too much – electrons are emitted and excess used as kinetic energy

8. Calculation Practice
The blue color of fireworks is often achieved by heating copper (I) chloride to about 1200°C. Then the compound emits blue light having a wavelength of 450nm. Calculate the frequency and quantum of energy that is emitted at 4.50x102nm by CuCl.

9. Spectra Radiation emitted from a source contains various λ's
When separated into its different λ's a spectrum is formed
Two types of spectrum
Continuous
Line

10. Continuous Spectra

11. Line Spectra

12. Bohr Model Assumed electrons orbit nucleus in circular patterns
To move between levels energy is absorbed or emitted
Ground State – electron at lowest energy
Excited State – electron is at higher energy state
Bohr model only accurately explains hydrogen

13. Wave Behavior of Matter
Lights is both a wave and a particle
Louis de Broglie believed matter could have wave properties
Applied idea to electrons
λ=h/m·v
Works for all matter so why don't we observed this in our everyday lives?

14. Uncertainty Principle
If matter can act as a wave we should be able to calculate position and velocity
Heisenberg determined that we cannot know both position and velocity of subatomic matter.
Solving Schrödinger's equation gives use probabilities of location.
The solutions correspond to the orbitals

15. Quantum Numbers n – principle quantum number
Whole numbers – 1,2,3,....
l – azimuthal quantum number
From 0 to n-1
Determines shape of orbital
0 → s = shape
1 → p = principle
2 → d = diffuse
3 → f = fundamental

16. Quantum Numbers Cont. ml – magnetic quantum number
Goes from -l to l including 0
Determines orientation of orbital
ms – spin quantum number
Two values +1/2 and -1/2
Determines spin of electron
No two electrons can have the same four quantum numbers – Pauli Exclusion Principle

17. Possible Quantum Numbers

18. Quantum Numbers Example
Which of the following sets of quantum numbers are not allowed? For each incorrect set, state why it is incorrect.
n = 3, l = 3, ml = 0, ms = -1/2
n = 4, l = 3, ml = 2, ms = -1/2
n = 4, l = 1, ml = 1, ms = +1/2

19. Quantum Numbers Practice
Which of the following sets of quantum numbers are not allowed? For each incorrect set, state why it is incorrect.
n = 2, l = 1, ml = -1, ms = -1
n = 5, l = -4, ml = 2, ms = +1/2
n = 3, l =1, ml = 2, ms = -1/2

20. Quantum Number Practice
What is the designation for the subshell with n = 5 and l = 1? How many orbitals are in this subshell? Indicate the values of ml for each of these orbitals.

21. Representation of Orbitals
S – orbital is a sphere

22. Representation of Orbitals
P Orbital

23. Representation of Orbitals
D Orbital

24. Atoms with Multiple Electrons
Shapes of orbitals remain the same
Energy of orbitals varies
For a given n, energy increases as l increases

25. Electron Configuration
Governed by three rules
Pauli Exclusion
Hund's Rule – for orbitals with same energy, lowest energy is attained when the number of electrons with the same spin is maximized
AUFBAU – Energy shells are filled from lowest energy to highest energy

26. Orbital Diagram

27. Where Orbitals are Filled

28. Filling Order

29. Electron Configuration Example
Write the electron configuration for the following elements:
Silicon
Chromium
Iodine

30. Electron Configuration Practice
Write the electron configuration for the following elements:
Copper
Sulfur
Tin

31. Exceptions to Filling Rules
Chromium
Copper
Silver
Molybdenum

32. Homework
2, 6, 10, 14, 20, 22, 26, 46, 52, 54, 60, 68

33. Electron Configurations