Another Exciting Post From The Mind of Ethan Spalding About the Wonderful Topic of: Ionic Bonding

In my second post of a three part series, I am going to go over the beginning of Ionic Bonding. We did not go over a ton in class, as we went over the test. So I'll quickly go through our notes.

First of all, the key component of this unit is the idea of Valence Electrons. These our the outermost electrons on the highest energy level. Every element wants to have 8 of them. Thus, they try to gain, lose or share them with other atoms.

There are two easy way to determine how many valence electrons an element has. One is to count how far it is from the left in the main block (Li 1, Be 2, and so on). The other is by the written electron formation. The number of S and P electrons in the final electron level combined is the number of valence electrons.

Ionic Bonding is the bonding between a metal and a non-metal. It is the "loss or Gain" of an electron from another.

Again, the next poster will be me :(

Periodicity (sorry for delay)

Ok so this will be the first of three posts today... I'm obviously just an idiot who cannot get a post done on time, and it has made me have to do 5 total posts :(.

The periodic table is really organized in a cool way. Atoms in the same rows and columns relate to each other in certain ways. This is called periodicity.

We discussed two characteristics today that follow patterns (and touched on a third). They are as follows:

Atomic Radius: Atomic radius is a manner of measuring atomic size. This is measured by taking the distance between nuclei of two atoms of the same type, and dividing that by two. So which ones are the biggest? That can be determined by where they located on the periodic table. By a rule of thumb, elements get smaller when you go across the periodic table, before jumping back up when moving on to the higher energy level. Why? When going up an energy level, the valence set is farther away, so that is why it increases as going down. When going across, the size depletes because with more protons, the positive charge pulls stronger.

First Ionization Energy: This is the amount of energy needed to get take an electron away from the element. It takes less energy to take elements on the bottom and left of the table, which is pretty easy to understand. Left elements only need to "lose" one element to become at a full valence level, so they naturally give them up easy. The ones on the bottom are even easier, because the valence level is so far away from the nucleus. That is obvious the reason why alkali metals (especially Francium) react so violently:

ElectroNegativity: This is the opposite of Ionization energy: How likely it is to gain an atom. It is based on a scale of 1-4, with 4 being the highest number. The lower the energy level and the farther right an element is (not counting the Noble Gases, which do not gain an electron), the higher the number. The reasons for this are obvious if you've read the rest of this post: the lower the energy level, the more proton pull, and elements on the far left want to gain an electron for a full valence level. Only one element has an electronegativity rating of 4: fluorine. So what happens when you put fluorine and francium together? Big big big boom.

So now you know about some periodic characteristics. Now, I will start to write on some more chemistry. The next scribe is: ME!!!!!!!!!!!!

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.

CHEM IS COOL

A couple of things to take care of:

Mateo has one "t".

The upside down "y" thing that is a symbol for wavelength is called a "lambda", and the "v" looking thing that is a symbol for frequency is called a "nu". In case you were wondering, which I was.

If you still do not understand the electro-magnetic spectrum, here's a quick song to ingrain it into your brain:

So today, we did some stuff. Like chemistry stuff. Like really complicated electron orbital stuff.

We began with the rather disheartening concept of not really being able to know where a specific electron is at a given point in time. What we do know is that there is a certain area around the nucleus where we will "probably" find the electrons. However, as Kaitlin astutely observed, there is a small small chance one of our poor electrons are lost somewhere, searching desperately for their owner.

From there, we moved on to Quantum Mechanics.

Mr. Lieberman stressed the difference between electron orbits and electron orbitals. In fact, electron orbits don't actually exist. This means that everything you learned about how atoms look is wrong. Everything that our middle school teachers taught us is wrong. And above all, Jimmy Neutron was wrong.

So now that everything you have learned has been upset, we are learning a new way to define an electron's location: orbitals. As stated previously, we cannot know for sure where an electron is in relation to its nucleus, or in relation to anything at all, so scientists decided to create a thing called an orbital. Basically, an orbital is an area where there is great er than a 90% chance of finding an electron. In pictures it looks more like a fuzzy sphere kind of thing. However, Bohr's electron is much more conducive to the imagination and is thus still taught in elementary and middle schools today.

Instead of equating the location of an electron with a sort of solar system looking model, scientists like to look at the location as sort of a seat in a stadium, or concert hall, or auditorium. Except in an altogether different fashion. Tickets typically have four pieces of information on them in regards to where your seat it, the gate you enter in, the section, the row, and the seat. Each one getting more specific, and no two tickets to the same seat.

Much like the ticket analogy, we "locate" electrons with a series of QUANTUM NUMBERS. Each number gets sequentially more specific until we can know the relative energy of the orbital, its shape, its orientation in space, and which direction its spinning.

The first number we must concern ourselves with is the value "n". Thankfully, "n" comes only in integer form (that is, 1,2,3,4,5,6,7,8,9,10,11...) basically, anything positive that is not a decimal or a fraction. It aids in determining the energy level of the electron; the larger the "n" value, the higher the energy. Unfortunately, we have only really found up to 7, though I guess we are fairly close to obtaining the eighth.

The next number is represented by the letter "l" (thats "L", only lower-case). This, fortunately, also comes in integer values. These are limited however, to a range extending from 0 to n-1. "n", if you recall, is the relative value of the energy of an electron. "l" is the energy sublevel of an electron, sort of like a more accurate way of defining an electron's energy. If you think about it, the number of sublevels is the same as whatever energy level it is in. If that doesn't make sense, the book says it like this:

In the nth principal level, there are n different sublevels.

So scientists think that it would be easier to memorize letters than numbers and gave us this:

sublevel: 0 | 1 | 2 | 3 | 4.....

letter : s | p | d | f | g ... and from here it goes on in alphabetical order. Furthermore, when defining a sublevel (using the s-p-d-f system), you put the principal (the n) value before it. So for example, if you were in level 4, sublevel 2, you would have 4d. If you were in level 5 sublevel 6, WAIT THATS NOT POSSIBLE. Its not possible because the sublevel values are confined between n-1 and 0. HA. But if you were in level 5 and sublevel 4, you would say 4g, like sprint.

Next variable: m_l. m_l is also made up of integer values, and its range extends from "l" to -"l". That's positive "l" to negative "l". This defines the orientation of the orbital. The higher the sublevel, the higher the number of orientations. For example, in sublevel 3, you can have m_l values of 3, 2, 1, 0, -1, -2 , -3. and they all represent different orientations of the electron cloud. Look at this picture if you still don't get it.

The final number is represented by m_s. Basically, it has nothing to do with the other values. It just shows which way the electron is spinning. It either has a value of +1/2 or - 1/2. Or sometimes referred to as "spin up" and "spin down". Each orbital can have 2 electrons, each one spinning a different direction.

This is a fun song, if you are working out, are just walking, or just sitting down, and it has a rather fitting title, considering it's the end of the post.

This is a tough unit, so don't worry if you don't get it right away. I'd be willing to help if you need it.

Anyway, stay awesome period six.

Next scribe is going to be.. peter w.... PPEEUTTUUHHRR DUUUEEHHBBLE-YUEEWW