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