1. Chpt 7 - Atomic Structure
Electromagnetic Radiation
Atomic Spectrum - Bohr Model
Quantum Mechanical Model
Orbital Shapes and Energies
Electronic Structure & Periodic Table
Periodic Trends
HW: Chpt 7 - pg , #s 23-27, 37-43, 54, 62, 65, 67-68, 70, 71, 74, 76, 82, 86, 98, 100, 102, 104, 115, 116, 119, 126

2. Electromagnetic Radiation
λ (lamba) = wavelength (m)
ν (nu) = frequency (Hertz, Hz or s-1)
E = energy
c = speed of light, x 108 m/s
c = λν they are inversely related
Know the relative order of radiation in E, λ, ν

3. 1900s Death of Classical Physics
Black Body Radiation
Planck’s hypothesis… energy is quantized
E = hν or ΔE = nhν n = integer
h = 6.626x10-34 J.s
Photoelectric effect
Einstein proposed EM radiation is quantized
A stream of “particles” called photons
E = hν = hc/λ
deBroglie λ = h/mv (wavelength of a particle) velocity in m/s mass in kg - so units cancel with J

5. Photoelectric Effect
Light with frequency lower than a specific threshold have no electrons emitted (no matter how intense it is)
Light with frequency greater than threshold emits electrons and number of electrons increases with intensity

6. Diffraction Pattern in a Crystal
Electron beam is diffracted off of a crystal.
Electron exhibits wave behavior!!!
Davisson Germer experiment - They shared Nobel prize with GP Thomson which did similar type experiment.

7. Continuous vs Discrete Spectrum
Continuous spectrum vs. discrete spectrum (line spectrum)
Absorption vs emission spectrum
Only certain energies are allowed for the electrons in any atom

8. Hydrogen Atom
The observed spectrum was explained by Bohr by proposing the electrons move around the nucleus in certain allowed circular orbits.

9. Bohr Energy Expression
Calculated from hydrogen atom spectrum
E = x10-18 J (Z2/n2)
Z = atomic number, 1 for hydrogen
n = orbital that the electron is located
ultimately only good for hydrogen atom spectrum

10. Quantum Mechanics
Schrodinger solved the problem mathematically (no real physical significance) treating electrons as waves.
Hψ = Eψ
ψ is the wave function of the electron’s coordinates in 3 dimensions
Heisenberg - uncertainty principle
Δx * Δ(mv) >= h/4π
position momentum
See Heisenberg laser slit video

11. Orbital shapes and Energies
Orbitals are simply then a probability distribution of where the electron could be found.
(left) probability function for s-orbital
(below) Radial probability function for s-orbital

12. Shapes of p and d orbitals

13. What can we know about electron?
4 Quantum numbers describe the electron in an orbital.
n is principle quantum number - relates to size of the orbital, n = 1, 2, 3, 4,…
l is angular momentum q.n. - relates to shape of orbital, l = 0, 1, 2, …, n - 1
s-orbital is l = 0
p-orbital is l = 1
d-orbital is l = 2
f-orbital is l = 3
ml is magnetic q.n. - relates to orientation in space ml = -l,…,0,…, +l
ms is electron spin q.n. - relates to spin of electron ms = - 1/2 or +1/2 (called spin up & spin down or clockwise/counter clockwise)

14. Quantum numbers
Examples of valid quantum numbers for various orbitals. In addition, spin +/- 1/2 for each individual orbital.

15. Energy Levels of orbitals
As we keep adding energy levels, we see as the principle quantum number, n, increases the number of sublevels (types of orbitals) increases. In addition the energy spacings get closer together 1s - 2s - 3s - 4s - etc. So the energy of the 4s orbital comes lower than the 3d. The order need not be memorized because the elements in Periodic Table shows it with its s,p,d,f blocks.

16. Electron Configuration rules
Electron’s occupy lowest energy level first - aufbau principle
Maximum of 2 electrons in any orbital - Pauli exclusion principle
If 2 electrons occupy the same orbital they have opposite spins. +1/2 or -1/2 also called spin up / down or clockwise / counter-clockwise
For degenerate orbitals (the same energy like the three p, five d, or seven f) use Hund’s rule, also known as the bus rule - only pair up the electrons if necessary.

17. General s,p,d,f blocks
The periodic table clearly shows that after the 3p orbital, the 4s fills before the 3d. Likewise, 6s 4f 5d 6p is the order when the lanthanides start.

18. Electron Configurations
A couple of exceptions Cr and Cu groups in the transition metals promote an s electron to achieve a half-filled and fully-filled set of d-orbitals because they have more stability.

19. Mendeleev’s Original Periodic Table
Organized by increasing atomic mass and put in columns by similar properties and reactivities
Left spaces for undiscovered elements together with predicted properties - these were confirmed by experimental results!!!

20. Periodic Table Trends some are found in Chpt 8
Ionization Energy
Electron Affinity
Atomic Radius
Ionic Radius
Electronegativity

21. Ionization Energies
The ionization energy is the energy necessary to remove an electron completely from an atom. X --> X-1 + e- The 2nd ionization energy would remove the next electron, etc.
Notice the trends in this chart a) across the period - general and detailed b) 1st ion E, 2nd ion E, etc. large jumps associated with core electrons.

22. 1st Ionization Energy Chart

23. Electron Affinity
EA is the energy change with adding an electron to an atom
X + e- --> X-1
This energy is correlated to thermodynamics, thus atoms that have a high EA (like to gain e-) the associated E change is negative (exothermic) the higher EA the more exo it is.
Generally, EA increases up a group and across a period.

24. Atomic Radius Radii are estimated from actual
spacing in metals or molecules

25. Ionic Radius Trends
Ionic radius of most common ion reported in picometers.
The size typically decrease across the period with a large jump when going from anion to cation.
Also, cations are smaller than their atoms and anions are larger than their atoms.

26. Electronegativity Trends
Electronegativity is the ability of an atom to attract electrons to itself in a chemical bond. It generally increases across a period and decreases down a group.

27. Alkali Metals Periodicity
The alkali metals are shown below with various physical properties. These are expected trends for other groups of metals as well.