1. Electron Configurations HW: read CH 11

2. Atoms and Energy Radiant Energy: knowledge led to refinements of atomic model A. Wave particle: light behaves as both a wave and a particle-strange! B. Electromagnetic Radiation (EMR): form of energy that exhibits behavior (light) as it travels through space (X-rays, visible light, radio waves) All EMR has a wavelength, frequency, and amplitude that determines its energy

3. Atoms and Energy Electromagnetic Spectrum Different types of radiation have different wavelengths Different colors have different wavelengths

4. Emission of Energy by Atoms When atoms receive energy from some source and become excited, they release energy by emitting light Light is different colors for different elements – must be because of various energy levels

5. Bohr Model Limitations Why is the Bohr model of the atom no longer accepted? 1. Does not explain behavior of more than 1 electron 2. NOT 3D representation 3. Impossible to know location of every electron https://www.youtube.com/watch?v= 8ROHpZ0A70I (4 min) https://www.youtube.com/watch?v= 8ROHpZ0A70I

6. Modern Theory: Atomic Orbitals Every energy level has sub levels: 1 st = s (lowest energy) 2 nd = s and p 3 rd = s and p and d 4 th = s and p and d and f s = 2 electrons p = 6 electrons d = 10 electrons f = 14 electrons

7. Let’s write this on our whiteboards

9. Orbital shapes Shapes indicate the region around the nucleus of an atom where an electron is likely to be found (90% chance) s = sphere/circle p = dumbbell d = dumbbells on 3 planes f = 2 dumbbells on 5 planes

11. Rules for electrons 1. Aufbau Principle: electrons are added 1 at a time to the lowest energy level until all electrons have been used. So 1s, 2s, 2p,... 2. Pauli Exclusion Principle: an atomic orbital may hold at most 2 electrons. Each electron must spin in opposite directions (arrows) 3. Hund’s Rule: orbitals of equal energy are occupied by 1 electron before any orbital is occupied by a 2nd electron (all must have 1 electron before having 2) Crash course https://www.youtube.com/watch?v=rcKilE9CdaA (12 min)https://www.youtube.com/watch?v=rcKilE9CdaA

13. Electron Config Examples (HINT: the superscript total should equal atomic number) Nitrogen – atomic number 7 1s 2 2s 2 2p 3

14. Electron Config Examples Silver – atomic number 47 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 9

15. Electron Config Examples YOU TRY: Magnesium – atomic number 12 YOU TRY: Gold – atomic number 79

16. Of course there will always be exceptions (fancy AP info) Chromium SHOULD be [Ar] 4s 2 3d 4 BUT is really [Ar] 4s 1 3d 5 Copper SHOULD be [Ar] 4s 2 3d 9 BUT is really [Ar] 4s 1 3d 10 This minimizes electron repulsions

17. Atoms get larger going down a group Atoms get smaller left to right across a period more protons as you move to the right; the more protons give more atomic pull. The strong attractive force shrinks orbitals so the atom is smaller Periodic trends – Atomic Radius

18. Recall: Ions are elements that gain or lose electrons More electrons the size becomes larger more repulsion's, spread out electrons, increase size Fewer electrons the size becomes smaller reduces repulsion's, electrons pulled closer to nucleus Periodic trends – Ionic Size

19. Energy needed to remove one of its electrons Atoms with high I.E. hold onto their electrons tightly Atoms with low I.E. easily lose electrons I.E decreases as you move down a group I.E. increases as you move from left to right think of octet rule for optimum number of electrons Periodic trends – Ionization Energy

20. Ionization Energy exceptions (more fancy AP info) It requires less energy to remove a p 4 than a p 3 due to increased repulsion with p 4 Ionization energy decreases a little between Be/B and Mg/Al and Ca/Ga because it goes from an s orbital (close to nucleus) to a p orbital (farther away from nucleus)

21. Ability to attract electrons, related to Ionization Energy Fluorine most electronegative (wants electron the most) Left side of table least electronegative (not want electrons) Periodic trends - Electronegativity