Battleship and Relations to the Periodic Table

We had recently learned how to write electron configurations and what they mean about an element. We know that the four quantum numbers are n, l (s,p,d,f), ml, and ms. These represent the principal energy levels, the sublevels, the orbitals, and the electron spin. The configuration, 1s2 2s2 2p6, for example, would be neon (Ne) because if you add the last number of each term (the individual electrons), then the sum is the total electrons; in this case 10. To practice writing these electron configurations, we split up into partners and played electron configuration battleship. Personally I had an intense game with Alex K, who put up a good fight, but in the end, couldn't hold up any longer as I sunk his battleship.

We also learned how to write the abbreviated version of an electron configuration. This involves writing the previous noble gas in brackets, followed by the configuration for just the next period. For example, the abbreviated electron configuration of Iodine (I) would be [Kr] 5s2 4d10 5p5. This is very helpful because the Kr takes care of the majority of the whole electron configuration (without it you would have to write: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5).

Orbital diagrams were next on dliebs's agenda. These are diagrams that have orbitals of each energy level that contain either an up arrow, a down arrow, or both. The s sublevels have one orbital so each s can hold only 2 electrons. P has 3 orbitals, d has 5 orbitals, and f has 7 orbitals. This means that as you progress from s to f, the number of electrons per sublevel goes up. Also, there is a certain way that electrons fill up an energy level. According to Hund's rule, each sublevel is filled up before another one starts to fill up. This means that 1s will have to be totally filled up before 2s or 2p. Also electrons will fill up so that they will be unpaired rather then paired up with another electron, meaning that each orbital isn't filled up at a time. Instead one-half of each orbital is filled in a sublevel before they are totally filled. After we learned this, dliebs told us that any electron in the top energy level can become "excited", and hypothetically, jump to any higher sublevel. This "excited" state explains why it releases energy, which is translated into light.

Finally we then had to determine how electron configurations and ions relate to the periodic table. Knowing how to interpret an electron configuration correctly will allow you to find any element on the periodic table. Also, ions can easily be explained when relating them to the periodic table and their electron configurations. Every element, wants to have their last energy level completely filled. To achieve this, their p sublevel has to be completely filled, and it can have no more electrons. All the noble gases (the family all the way to the right) have their last energy level completely filled. In this way, every other element wants to do the same, so they either have to gain or lose electrons to achieve their goal. For Na to have a completely filled p sublevel, it is easiest for it to lose one electron and have the same number of electrons as Ne. This why the Na ion looks like Na+, because it gains an electron. This works for many of the main-group elements, but not for the transition metals because they are rebels.

Ya, well I hope this post has been eye-opening and that you can now be a master of not only electron configurations, but also how they and ions relate to the periodic table. Also sorry for the lack of pictures but I wasn't able to take any.

Anyways..........The next scribe will be THE LOVELY, THE RESPECTED, THE MAGNIFICENT (NEED I SAY MORE)

Ethan Spalding

Electron Configuration

Electron configurations is determined by the sublevel energies and the element. An electron configuration is a short hand for showing where certain electrons are located on a certain atom of an element. Electron configuration is the assignment of quantum numbers to each electron in an atom of a given element. Electron configuration shows the number of electrons (which is the exponent or superscript after the letter) in each sublevel. For instance the book gives the example:

1s²2s²2p⁵

^{Which means that there are two electrons in the sublevel 1s and 2 in the sublevel 2s and 5 electrons in the sublevel 2p. This means that this particular element has 9 electrons which means it is Fluorine.}

When it is written as 1s² this refers to the electron being in the region of N = 1.Which means that the only "l" option that is possible is 0 (thats zero not "O" like in oh no!) or "s". Because l is 0 there are only two electrons that can be there, and since the 1s sublevel is full (which it must be to begin filling up the next sublevel) the superscript says two (meaning two electrons). If we continued writing a substance's Electron configuration we could go on for a long time but since this has only 9 electrons it would stop at 2p. Each letter ("s", "p", "d", "f") has a top level of electrons of which more than such it cannot hold. for instance, "s" can only hold 2 electrons maximum, "p" can only hold 6 electrons maximum, "d" can only hold 10 electrons maximum, and "f" can only hold 14 electrons maximum.

So for example when all are full it would look like this:

1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p⁶4d¹⁰4f¹⁴

^{Which has 60 electrons so this indicates that it would be the element Neodymium.}

^{Electron Configuration is a simple way of writing out the electrons positioning of an element, which is important to know for doing other things, and can tell us a lot about an element.}

^{soooo... yeah. thats Electron configuration,} Electron configuration, Electron configuration, Electron configuration which makes the tenth time I've said Electron configuration... 11 actually. So thats about it we will learn more on monday, i think, i dont actualy plan the lessons but whatever, the next scribe will be: Joshua D-D-D-EIN!!! (said in monster truck rally voice)

thankyou, and farewell, i will see y'all Monday.