1. Unit 3 - The Modern Atom What is our model of the Atom? What is wrong with it? Homework: pg. 333-336 Q&P # 7, 8, 12-14, 20, 25, 31, 32, 36-39, 45, 50, 59, 70, 73, 75, 80, 81

2. Electromagnetic Spectrum review What are some parts of the spectrum? It propagates how? What are the parts of a wave? how are they related? frequency, nu ν wavelength, lambda λ c = λ λ  ν c = speed of light photons - packets of light energy

3. Electromagnetic Radiation  (lamba) = wavelength (m)  (nu) = frequency (Hertz, Hz or s -1 ) E = energy c = speed of light, 2.9979 x 10 8 m/s c =  they are inversely related Know the relative order of radiation in E, 

4. EM Spectrum

5. EM spectrum visual

6. Emission from atoms Atoms absorb or emit energy. Becomes “excited” when absorbs energy Can release energy by emitting a photon. Hydrogen Helium Carbon

7. Energy Levels of H atom The H atom has the experimental spectrum from the previous slide Only certain types of photons were produced! The H atom must have certain discrete energy levels. Levels are quantized. Model constructed by Bohr - like planets orbiting the sun. E = h x ν h = planck’s constant

8. Birth of Quantum Mechanics Black-body radiation - why does piece of iron change colors when heated? Planck’s explanation E = h x ν Photoelectric effect - light shining on a metal surface can emit electrons from surface. Einstein - light behaves like particle Bohr’s explanation of H-atom - little book pg 139 de Broglie - wavelength of a particle!!!! - little book pg 144

9. 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

10. Where is this going? Classical mechanics - visible objects at ordinary velocities Quantum mechanics - describes behavior of extremely small objects at velocities near the speed of light - (Dr. Quantum video) Heisenberg - Uncertainty Principle - impossible to know both position and momentum of object at same time - little book pg 148 (slit video)

11. Wave Mechanical Model Schrodinger solved wave problem mathematically ! - no physical meaning Orbits became orbitals - a region of space with a probability of finding the electron. Note: calculate the max probability for the H atom - you get 53pm, which is the same as calculated from Bohr’s model.

12. 12 Where are the electrons? These energy level rules tell us where the electrons are and how they are arranged in an atom. Principle Energy level n= 1,2,3,4,... Each Principle energy level has sublevels (orbitals) n=1 has 1 sublevel, n=2 has 2 sublevels, n=3 has 3 sublevels types of orbitals s, p, d, and f s has 1 orbital, p has 3 orbitals, d has 5 orbitals, f has 7 orbitals Now what? 12

13. 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.

14. Shapes of p and d orbitals

15. Electron Configuration rules 1. Electron’s occupy lowest energy level first - aufbau principle 2. 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 3. 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.

16. Periodic Table Practice some electron configurations – H 1electron - 1s 1 – He 2 electrons - 1s 2 – Li 3 electrons - 1s 2 2s 1 – Be - 1s 2 2s 2 – B - 1s 2 2s 2 2p 1 write e- configurations up to Ar – filling p orbitals - remember bus rule (unpaired e - ) also p has 3 orbitals so 6 total e- can fit in them Valence electrons (outermost e - ) -largest n – electron configuration of just valence e- – electron dot notation Review Chpt 4.8 - Parts of Periodic Table

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. Periodic Table 7 Periods and 18 Groups Group 1 - alkali metals - most reactive metals, Fr most Group 2 - alkali earth metals - next most reactive Group 17 - halogens - most reactive nonmetals, F most Group 18 - noble gases - colorless, odorless, nonreactive Valence electrons, normal ionic charge (oxidation number) gain or lose electrons when make ionic compounds electron configurations Staircase metalloids, metals and nonmetals Diatomic elements... Br I N Cl H O F Carbon has allotropes, also oxygen

19. Periodic Properties Most properties of elements follow a trend in the periodic table. 2 types of trends Group trend - properties change some as the elements go up or down a column. Period trend - properties change some as the elements go across a row.

20. Ionization Energy The ionization energy is the energy necessary to remove an electron completely from an atom. X --> X +1 + e - Group Trend - Electrons going into orbitals with larger principal energy levels(further from the nucleus) Period Trend - electrons going into orbitals with same principle energy level What else is going on across a period?

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

22. 22 Ionic Radius Cations larger or smaller than original atom? Group Trend? Period Trend? Anions larger or smaller than original atom? Group Trend? Period Trend? Mixed comparison 22

23. 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.

24. 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.