Build the Big Ones Lab

Today in class we finally got to a a lab! It was called Build the Big Ones. The big ones refers to molecules with more than one central atom. The molecules we worked with were ethanol, acetic acid (vinegar), serine, and styrene. We had to build each molecule using spheres, wooden sticks , and metal springs. The wooden sticks represented a single bond and the metal springs represented double bonds. As explained in the directions, a black sphere is carbon, a yellow sphere is hydrogen, a red sphere is oxygen, and a light blue sphere is nitrogen. Here are the pictures for each molecule:


ethanol


acetic acid


serine


styrene

For each molecule we had to determine the central atom, the geometry of the atom, an example would be tetrahedral, the bond angle for that geometry and the hydridization. Here is what me and my partner Katie I. cam up with:

Lastly, we had to do our conclusion for the lab and if any of you had trouble, this is the conclusion Katie and I came up with:

I hope our data table and conlcusion help! Well that was all we did in period 6 today.
The next scribe is......................................Kaitlyn Y.

Continuation on Electron Pairs and Start of Hybridization

To start off class today, we first discussed the types and names of the electron pairs we learned yesterday (In Ethan's Post). Next, we moved into looking at 4,5, and 6 areas of concentration. These are the highest levels of concentration that can ever happen in an atom.

The first molecule that we looked at for the day was PF5. PF5 has 40 valence electrons and using this, we constructed the Lewis structure. Phosphorus can break the octet rule, which is the only reason why this Lewis Structure works. Then we used that to help us form the angles at which the balloons in the demonstration were forming(90,120,180). PF5 is an example of a trigonal bipyramid molecule. We then did the same steps for the Octahedron molecule of SF6. Sulfur can also break the octet rule.

During the demonstrations of the shapes of how these molecules would look like if they were much bigger, we also saw how the non-bonding pairs of atoms continuously took away an equatorial angle, instead of a linear angle. This happens because this is the most stable configuration.

The Last thing we learned about for the day was hybridization. Hybridization is the mixing of one s orbital and one p orbital in order to produce two sp orbitals. We looked at a video that explained the different formations as well as the actual action of what it would look like. . That is all we learned.

The next scribe is...Deena M.

My. Last. Post. Heck. Yeah. It's on 3D Molecular Design.

So we thought we were done. The Lewis Structure was the life, being able to show how atoms were bonded. But guess what, my friends? You don't live in a 2D world made by 2D atoms. Otherwise, you'd look like this:

So how does this work? Well atoms form bonds, as I covered yesterday. They do this in a 3D manner. Let's begin with what we learned today:

Linear Bonds: are bonds that use all the electron pairs in the valence level, there are no "free" electrons. These have a 180 degree angle.

Next, we have the Trigonal Planar Bond, which is the same as above except with three bonds. The angle measure is 120 degrees:

Sorry for the large image size. Next, we have Tetrahedral Bonds, with 4 bonds and an angle measure of 109.5 degrees:

Finally, we have the Bent Bonds, which have electron pairs in place of bonds. This causes the angles to change, and the shape to "bend". The molecules with 2 bond pairs have 120 degrees, the ones with 3 have around 107 degrees.

Now, the moment you've all been waiting for.... BEN A. will be my replacement after my long tenure as writer. Thank you and have a nice day.

an old friend

just wanted to say hello. hello.

Resonance (My Second to last post!!!:))

So we all get covalent bonding now? As I detailed in my last post, there some easy ways to draw the Lewis Dot Diagram (click the link or scroll down the blog for the deets). However, just when we think everything is hunky-dory, Liebs decides to throw a big ole monkey wrench in the mix. And while Resonance isn't rocket surgery, it can still seem a bit complicated. To make it less complicated, I will give you a great song about a great dancer that has absolutely no revelance to chemistry whatsoever:


Keep the jeep rolling people! Now, back to resonance. For an overview, the Wikipedia is a good starter:

In chemistry, resonance or mesomerism [1] is a way of describing delocalized electrons within certain molecules or polyatomic ions where the bonding cannot be expressed by one single Lewis formula. A molecule or ion with such delocalized electrons is represented by several contributing structures

A few of these words may be a bit complicated. Therefor, I will expand upon it in a simpler manner:
Basically, Resonance happens when there are two different bond types when between 3 or more atoms of 2 elements. For example, let's say you have the element NO2-. Well, if you did the Lewis dot structure, you would get this result:

Notice that there is one double bond, and one single bond. Well, for reasons unbeknown to me, this is not how the element really exists. Instead, they exist in a way that is "in between" (Liebs described it as "one and a half" so it's easier to visualize, though it is not technically one and a half). Therefor, this is how they should be drawn:

Well, that's what we learned today. For the remaining time, we were generously given time to work on our worksheets. Whether you used it wisely (and you can assume if I did or not) is up to you. So basically yeah. I'll be back tomorrow. Stay Classy HCPER61011.

Covalent Bonding

On Thursday, we learned about Covalent bonding. Simply, this is the "sharing" of electrons to reach a full valence level. Covalent Bonding only happens between two nonmetals.

Covalent bonding can occur in single, double or triple bonds. It cannot occur in quadruple bonds. That is why carbon cannot form a C2 compound to fill its valence level.

Compounds are diagrammed in Lewis Dot Structures. These diagrams can be simple:

Or very complex:

Just throwing elements together can really be complicated if you do not do it right, however. But Mr. Lieberman, being the great teacher he is, gave us a few of guidelines to draw them:

1. Count the # of valence electrons (total) in the compound. Remember to keep ions intact.

2. Predict the arrangement of atoms in the compound. The first atom in the formula is generally in the center, unless it's Hydrogen.

3. Form single bonds, and see how many valence electrons are still left over.

4. Place electrons around the exterior elements until they have a full valence level, then do the same for the center atom. Remember, their can only be 8 total electrons around each element, unless it is on a row with transition metals.

5. If more electrons are needed, draw double and triple bonds.

And then you've got a Lewis Diagram!!!! (That was fun, wasn't it). The next scribe will be: MEEEEE :(