1. Atomic Orbitals and Electron Configurations

2. Quantum Mechanics Better than any previous model, quantum mechanics does explain how the atom behaves. Quantum mechanics treats electrons not as particles, but more as waves (like light waves) which can gain or lose energy. But they can’t gain or lose just any amount of energy. They gain or lose a “quantum” of energy. A quantum is just an amount of energy that the electron needs to gain (or lose) to move to the next energy level. In this case it is losing the energy and dropping a level.

3. Atomic Orbitals The energy levels in quantum mechanics describe locations where you are likely to find an electron. Orbitals are “geometric shapes” around the nucleus where electrons are found. Quantum mechanics calculates the probabilities where you are “likely” to find electrons.

4. Atomic Orbitals An electron can be found anywhere. Scientists agreed to limit these calculations to locations where there was at least a 90% chance of finding an electron. Orbitals as sort of a "border” for spaces around the nucleus inside which electrons are allowed. No more than 2 electrons can ever be in 1 orbital. The orbital just defines an “area” where you can find an electron. What is the chance of finding an electron in the nucleus? Zero. There aren’t any electrons in the nucleus.

5. Energy Levels Quantum mechanics has a principal quantum number. It is represented by a little n. It represents the “energy level” similar to Bohr’s model. n = 1 describes the first energy level n = 2 describes the second energy level Etc. Each energy level represents a period or row on the periodic table. Red n = 1 Orange n = 2 Yellow n = 3 Green n = 4 Blue n = 5 Indigo n = 6 Violet n = 7

6. Sub-levels = Specific Atomic Orbitals Each energy level has 1 or more “sub-levels” which describe the specific “atomic orbitals” for that level. n = 1 has 1 sub-level (the “s” orbital) n = 2 has 2 sub-levels (“s” and “p”) n = 3 has 3 sub-levels (“s”, “p” and “d”) n = 4 has 4 sub-levels (“s”, “p”, “d” and “f”) There are 4 types of atomic orbitals: s, p, d and f Each of these sub-levels represent the blocks on the periodic table.

7. Orbitals In the s block, electrons are in s orbitals. In the p block, the s orbitals are full. New electrons fill the p orbitals. In the d block, the s and p orbitals are full. New electrons fill the d orbitals s p d

8. Orbitals Complete the chart in your notes as we discuss this. The first level (n=1) has an s orbital. It has only 1. There are no other orbitals in the first energy level. We call this orbital the 1s orbital. Energy Level Sub- levels Total OrbitalsTotal Electrons per level Total Electrons n = 1s1 (1s orbital)22 n = 2spsp 1 (2s orbital) 3 (2p orbitals) 2626 8 n = 3spdspd 1 (3s orbital) 3 (3p orbitals) 5 (3d orbitals) 2 6 10 18 n = 4spdfspdf 1 (4s orbital) 3 (4p orbitals) 5 (4d orbitals) 7 (4f orbitals) 2 6 10 14 32

9. Where are these Orbitals? 1s 2s 3s 4s 5s 6s 7s 3d 7p 6p 5p 4p 3p 2p 5f 4f 6d 5d 4d

10. Electron Configurations The electron configuration is the specific way in which the atomic orbitals are filled. Think of it as being similar to an address. The electron configuration tells me where all the electrons “live.”

11. Rules for Electron Configurations In order to write an electron configuration, we need to know the RULES. 3 rules govern electron configurations. Aufbau Principle Pauli Exclusion Principle Hund’s Rule Using the orbital filling diagram at the right will help you figure out HOW to write them Start with the 1s orbital. Fill each orbital completely and then go to the next one, until all of the elements have been accounted for.

12. Fill Lower Energy Orbitals FIRST The Aufbau Principle states that electrons enter the lowest energy orbitals first. The lower the principal quantum number (n) the lower the energy. Within an energy level, s orbitals are the lowest energy, followed by p, d and then f. f orbitals are the highest energy for that level. Each line represents an orbital. 1 (s), 3 (p), 5 (d), 7 (f) Low Energy High Energy

13. No more than 2 Electrons in Any Orbital…ever. The Pauli Exclusion Principle states that an atomic orbital may have up to 2 electrons and then it is full. The spins have to be paired. We usually represent this with an up arrow and a down arrow. Since there is only 1 s orbital per energy level, only 2 electrons fill that orbital. Quantum numbers describe an electrons position, and no 2 electrons can have the exact same quantum numbers. Because of that, electrons must have opposite spins from each other in order to “share” the same orbital.

14. Hund’s Rule Hund’s Rule states that when you get to degenerate orbitals, you fill them all half way first, and then you start pairing up the electrons. What are degenerate orbitals? Degenerate means they have the same energy. The 3 p orbitals on each level are degenerate, because they all have the same energy. Similarly, the d and f orbitals are degenerate. Don’t pair up the 2p electrons until all 3 orbitals are half full.

15. Electron Configurations ElementConfigurationElementConfiguration H Z=11s 1 He Z=21s 2 Li Z=31s 2 2s 1 Be Z=41s 2 2s 2 B Z=51s 2 2s 2 2p 1 C Z=61s 2 2s 2 2p 2 N Z=71s 2 2s 2 2p 3 O Z=81s 2 2s 2 2p 4 F Z=91s 2 2s 2 2p 5 Ne Z=101s 2 2s 2 2p 6 (2p is now full) Na Z=111s 2 2s 2 2p 6 3s 1 Cl Z=171s 2 2s 2 2p 6 3s 2 3p 5 K Z=191s 2 2s 2 2p 6 3s 2 3p 6 4s 1 Sc Z=211s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 Fe Z=261s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 Br Z=351s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 Note that all the numbers in the electron configuration add up to the atomic number for that element. Ex: for Ne (Z=10), 2+2+6 = 10

16. Electron Configurations ElementConfiguration H Z=11s 1 Li Z=31s 2 2s 1 Na Z=111s 2 2s 2 2p 6 3s 1 K Z=191s 2 2s 2 2p 6 3s 2 3p 6 4s 1 This similar configuration causes them to behave the same chemically. It’s for that reason they are in the same family or group on the periodic table. Each group will have the same ending configuration, in this case something that ends in s 1.