1. Electron Configurations
Now that we know electrons occupy orbitals with different shapes, or energy levels or orientations, we need to find a way to communicate this information:
This is what we call electron configurations!

2. Energy Levels - Rules Energy levels of lowest energy are filled first
When n increases, the energy of an electron increases
Within one energy level, energy increases in the subshells from s, p, d to f

3. Energy Levels - Rules
**As atoms become larger and the main energy levels become closer together, some sublevels start to overlap in energy.
Chart below helps to memorize proper order.

4. Filling Orbitals with Electrons
RULES FOR CREATING ENERGY LEVEL DIAGRAMS
1.The Aufbau Principal – electrons will always occupy lowest available energy level
2. Pauli Exclusion Principal – no two electrons have the same set of four quantum numbers
3. Hunds Rule – electrons remain unpaired for as long as possible
* Use boxes represent orbitals and arrows to represent electrons

5. Hydrogen Electron Configuration: 1s1
This is the hydrogen atom with its one electron in its ground state (lowest energy). The "1" in front of "1s1" means this is n=1. The "1" exponent means there's just "1" electron in that orbital. The "↑" indicates a +½ spin. Whether electrons actually spin is debated but it's one property of electron's configuration. When n=1, then l and m quantum numbers can only be zero. If the electron was boosted to the second shell (n=2), then the arrow below would be in the 2s box.

6. Helium Electron Configuration: 1s2
This is the helium atom. It has 2 electrons. These are able to share the same s orbital because of opposite spins. The "2" in "1s2" means there are 2 electrons in that orbital. "s" orbitals can only hold 2 electrons, so that's as high as you will ever see it. The 1 in "1s2" means this is the s orbital belonging to the n=1 shell.

7. Lithium Electron Configuration:
This is lithium. The third electron is in shell n=2.

8. Beryllium Electron Configuration:
This is beryllium. The fourth electron completes the 2s orbital. Again, it has to have opposite spin.

9. Boron Electron Configuration:
This is boron. The 2s orbital is full, so the expansion begins in the 2p orbital. "2p" means this orbital belongs to the n=2 shell. The fifth electron is the first p orbital electron. It is drawn as a particle in the upper image but actually has the dumb bell shape of a p orbital shown in the lower image. Notice the first p orbital occurs when the l quantum number is 1. So when you hear l=1, think p orbital. Also, note, that l starts at 0, and then counts up to n-1. Since this n=2 row, l started as l=0 for the s orbital and can only go up to l=1 (one less than the n value).

10. Carbon Electron Configuration:
This is carbon. It is adds another electron to the p orbital. This time I'm showing the electron cloud of the p orbitals. Notice the x, y, and z axes. The first 3 p orbital electrons will align with each axis. Carbon only has 2 p orbital electrons. I put the first one (pink) on the x axis and the second one (yellow) on the y axis. They don't share the same space by having opposite spins because that takes more energy than filling up an empty p orbital. Changing orientation keeps the electrons farther apart. The different orientations are indicated with the m quantum number. This is also called the magnetic quantum number. Notice the m quantum number starts with the negative of the l quantum number. Since we are at the l=1 quantum number (p orbital), the m quantum numbers start out as -1, then go up to 0 on the next electron, and +1 on the third electron. Each time it changes its orientation along a different axis.

11. Oxygen Electron Configuration:
This is oxygen. We skipped nitrogen, but nitrogen added one electron to the z axis of the p orbital. Oxygen has 8 protons and 8 electrons. The first 4 are in the first 2s orbitals. The last electron was forced to share one of the p orbitals with an existing electron. I put it with the one on the z axis, but they are all the same. Notice the m quantum number increased to +1 for the 3rd and 4th electrons aligning them with the z axis. However, they must have opposite spins to share that same space. By the way, the 2 unpaired electrons in m=-1 and m=0 orientations are responsible for the magnetic properties of oxygen.

12. Neon Electron Configuration
This is neon. We skipped fluorine, which added one more electron to fill 2p2y. At neon all of the p orbitals are filled. It has 3 orientations (x, y, and z) that are filled with 2 electrons each. This is a very stable configuration for the 8 electrons in the 2s and 2p orbitals. You may know it as the famous octet that many elements try to reach. So we see all of the p orbital electrons have the same n=2 principal quantum number and the same l=1 angular quantum number. Pairs of electrons of opposite spins are aligned along x, y, and z axes indicated by the m=-1, m=0, and m=-1 magnetic quantum numbers.

14. Explaining the Periodic Table
The maximum number of electrons in the s, p, d, and f orbitals corresponds exactly to the number of columns of elements in the s, p, d, and f blocks in the periodic table.
For example, the transition elements are elements filling the d energy sublevel with electrons (the 5d orbitals can accommodate 10 electrons, and there are 10 elements in each transition-metal period)

15. Below it all comes together to show how these quantum numbers actually define the Periodic Table of the Elements (see image below). The first two columns on the left are when l=0, which are s orbitals. "m" is only zero, so there is only one orientation (spherical). The +1/2 and -1/2 spins gives us two elements in this block. The right green block represents "l=1" which is the p orbitals. Since "m" has 3 orientations and there is a + and - spin for each orientation, there is a progression of 6 electrons that cover these combinations. That's why this block is six elements wide. The middle yellow block is for "l=2", which is the d orbitals. As mentioned, m goes from -2 to +2 giving us 5 orientations. With the two spins, it takes 10 steps to fill up the d orbitals, which is why there's a run of 10 elements that go across this middle section. Finally, the Periodic Table always shows a separate block of elements. This is for "l=3", which is the f orbitals. Since m can go from -3 to +3, that's 7 steps. With both spins, that means there it takes 14 electrons to fill up the f orbital and why there are 14 elements in that row. 15

16. Periodic Table
After the 2p electron orbital is filled, the elements start filling the 3s orbital. The filling is just like it did in row 2. So if you look at rows 2 and 3 on the Periodic Table, you will see they look the same. In other words, 2 electrons fill the s orbital and 6 electrons fill the p orbital. The only difference in the quantum numbers is the n=3, which is reflected in the electron configuration notation as 3s23p6. Note: I left off the "s=" below to save space. The spins of +½ and -½ are obviousThe purpose of learning these electron configurations is to know what electrons are available for bonding. This gives you insight in how elements will bond. In the next chapter when you learn more about bonding, you will see the usefulness of these electron configurations.

17. Shorthand Electron Configuration
Can use the preceeding noble gas in brackets

18. Explaining Ion Charges
Examples:
1) Zn: [Ar]4s23d10
Zn2+: [Ar]3d10
2) Pb: [Xe]6s24f145d106p2
Pb2+: [Xe]6s24f145d10
Pb4+: [Xe]4f145d10
Electrons from highest energy levels are lost first.

19. Some Unexpected Electron Configurations
Examples: Cu and Cr
Expected
Actual
Cr: [Ar]3d44s2
[Ar]3d54s1
Cu: [Ar]3d94s2
[Ar]3d104s1

20. Explanation…
in each case an electron is borrowed from the 4s subshell and placed in 3d subshell
The Cr 3d subshell becomes half-filled
The Cu 3d subshell becomes full
Half-filled and fully filled subshells tend to be more stable
The progression is pretty consistent except for a few exceptions. For example, look at chromium (group 6, period 4) and copper (group 11, period 4). Chromium shows 4s13d5. Instead of having 2 electrons in the 4s2 and 4 electrons filling the first 4 electrons in 3d orbital (3d4) as predicted, it took an electron from 4s2 and leaving it as 4s1 and filled the first 5 electrons in 3d orbital (3d5). There still are just 6 electrons (group 6) but they are not split 2 in the 4s orbital and 4 in the 3d orbital. The 4s13d5 means 1 is in the 4s orbital and 5 are in the 3d orbital. So every once in awhile, you will see these anomalies.  It's not fully understood why this happens, but it's probably because that arrangement is more stable (has lower energy).

21. The Whole Periodic Table
Below is the entire Periodic Table which shows the electronic configuration for all elements. This was tough to create, but I got it done. The progression is pretty consistent except for a few exceptions. For example, look at chromium (group 6, period 4) and copper (group 11, period 4). Chromium shows 4s13d5. Instead of having 2 electrons in the 4s2 and 4 electrons filling the first 4 electrons in 3d orbital (3d4) as predicted, it took an electron from 4s2 and leaving it as 4s1 and filled the first 5 electrons in 3d orbital (3d5). There still are just 6 electrons (group 6) but they are not split 2 in the 4s orbital and 4 in the 3d orbital. The 4s13d5 means 1 is in the 4s orbital and 5 are in the 3d orbital. So every once in awhile, you will see these anomalies.  It's not fully understood why this happens, but it's probably because that arrangement is more stable (has lower energy).

22. Homework Read p181-188 Answer p187 #13-15, 17; p188 #12
Worksheet – Electron Configuration