Test Day

Today we took our test over the Chemical Bonding unit and overall, I thought it went ok. I didn't have much time to study due to soccer and other issues lately, so I didn't feel as prepared as I could have walking into it.

There were a few questions on the test about covalent/ionic bonds, as well as some about polarity that I'm pretty sure I got wrong.

Overall, I thought the test was pretty straightforward and I'm hoping that that equals a good grade...but most likely not!

Hybridization

One of the things that we learned this short unit was hybridization in correlation to chemical bonding.

The name is a lot scarier than the actual process, I promise you.

All you need to do is count the number of electron domains, or places where potential bonds can occur on the atom, and put that number in the basic electron configuration model (spdf).

If you remember, S has 1 orbital, p 3, d 5, and f 7.

So if a bond has 5 electron domains on the structure, it has a hybridization of sp3d.

If a bond has 7 electron domains it would be sp3d3.

As this picture demonstrates, the hybridization corresponds with the electron pair geometry of a bond, due to both being determined by the number of electron domains.

Here are some links if you want more info:

https://chemistry.boisestate.edu/richardbanks/inorganic/bonding%20and%20hybridization/bonding_hybridization.htm

http://chemwiki.ucdavis.edu/Core/Theoretical_Chemistry/Chemical_Bonding/Valence_Bond_Theory/Hybridization

http://www.mhhe.com/physsci/chemistry/carey5e/Ch02/ch2-3.html

Also, here's a chem pun because it's funny. Ha. Ha.

Lewis Structures

The entirety of this unit is based on the ability to create the BEST lewis structure for each bond.

There are a few guiding rules to doing that,  ut mostly it is just practicing and trial and error.

One rule is that electronegative elements do not like to carry positive charges, so if one does it is most likely not the BEST structure.

Another rule is that charges on elements prefer to be spread out over the entire bond, instead of all piled onto one elements. For example, if one O in NO3 has a -3 charge on it while the rest are at 0, there is a better Lewis structure possible.

One more thing to make certain is that the charges on each part of the bond correspond with the overall charge. You can't have 2, -1 O's on a bond and have the overall charge be +4.

Here's a picture demonstrating this concept:

Here are some additional links if you want some more help:

http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/lewis.html

http://www.chem.ucla.edu/harding/lewisdots.html

http://antoine.frostburg.edu/chem/senese/101/bonds/faq/simple-lewis-structures.shtml