Biodiesel Links

Yesterday we began researching biodiesel for a competitive video. After reading, I found that it is much cleaner and healthier for the environment compared to petroleum diesel. In addition, it can improve the economy by providing jobs and making America more dependent on itself. Here are links with additional information:

Biodiesel Benefits 

Other Biodiesel Benefits

Biodiesel

http://www.alternative-energy-news.info/technology/biofuels/biodiesel-fuel/

http://butane.chem.uiuc.edu/pshapley/GenChem2/B11/3.html

Helpful Study Links

While studying for the chemistry test tomorrow, I found these videos most clear in explaining all the concepts I found most difficult to understand in the chapter. I hope you  find these helpful!

Formal Charges: Calculating Formal Charge
Drawing 3 D Structures 
Dipole Moment
Formal Charge and Dot Structures 

Bond Polarity and different bonds

In our last lesson over the unit, we learned about polarity and different types of bonds. With polarity, electrons are not always shared equally between two atoms. It is often the case where one atom may exert more force on the electron cloud than the other. The polarity of an atom is determined by the difference in electronegativity between the atoms. A coavlent bond is nonpolar if the difference between the two is only a 0.2-0.5 difference. A covalent bond in polar if the difference is ~0.5-1.6. All other bonds with a difference higher than 1.6  are most likely ionic. Covalent bonds regularly occur between two nonmetals and sometimes a nonmetal and a meatalloid. An ionic bond in normally between a metal and a nonmetal. Here is an image that displays this concept.

http://2012books.lardbucket.org/books/principles-of-general-chemistry-v1.0/s12-09-polar-covalent-bonds.html

Also, here are some links that further explain the topic:

study.com

polar and nonpolar bonds

Shape of Molecules

In class today, we learned about resonance and molecular shape. To have resonance, a compound must have multiple bonds and when you move a bond, the atom you move it to must be able to take the bond. Resonance equalizes bond length and bond strength. Below is an image exemplifying this concept:

http://www.chem.ucla.edu/harding/tutorials/resonance/draw_res_str.html

Next, discussed the shapes of molecules. The valence shell electron pair repulsion theory is used to predict molecular shape. Something to remember is that shape determines function. So, in this concept, the way the molecules arrange themselves in a space depends on the number of lone pairs and bonded entities present.

There are 5 shapes that we are focusing on. First, a tetrahedral molecule has 4 bonded entities around one central atom. A triagnol pyramidal molecule has three bonded entities and one lone pair of electrons on the central atom. Next, a bent molecule has two bonded entities with two lone pairs of electrons on the central atom. A linear molecule has two bonded entities to the central atom that does not have a lone pair. Lastly, a triagnol planar molecule has three bonded entities to the central atom and the central atom does no have any lone pairs. Here is a picture that shows a few of these:

http://cnx.org/contents/d5d1d182-3eb0-419a-bfda-3615e56fafea@1

Here are a couple helpful links that further explain these new ideas:

Molecular Geometry
Resonance

First Lesson

Today in class, we learned our first lesson in our new unit: Chemical Bonding. We first discussed the Lewis Dot Concept to illustrate chemical bonding. In one of these diagrams, you simply place electrons around the element symbol, making sure to fill each side of electrons first before pairing them up. Here is a chart that shows how this is done with different elements:

http://edtech2.boisestate.edu/lindabennett1/502/Compounds%20and%20Naming/Lewis%20Dot.html

The octet rule technically determines how many electrons are to be placed in the valence shell. An atom can not have more that 8 electrons total in their outer shell. However, there are exceptions. Hydrogen and Helium only have up to two valence electrons, boron requires 6 electrons to be stable, and beryllium only needs 4 electron to be stable. 

Another concept we learned today was how to solve electron dot formulas of molecules with the HAVE NEED SHARE method. Below is an example:

Here are a couple websites to further explain and practice these new concepts:

Purdue

Lewis Dot Structures

Practice Problems

Periodic Trends

In our last lesson of the unit, we learned about periodic trends. There are 4 trends that mainly focus on the S and P blocks on the periodic table. First, we learned about atomic size. As you move across a period from left to right, atomic size decreases because there is an increase in protons. Since the protons increasingly attract the electrons, the electrons pull toward the protons more and more, making the size smaller as you go to the right. Here is a chart that displays this trend:

https://www.boundless.com/chemistry/textbooks/boundless-chemistry-textbook/nonmetallic-elements-21/properties-of-nonmetals-147/atomic-size-569-7507/

Next, we discussed ionization energy. Ionization energy is the energy needed to remove an electron from a gaseous state. As you move up and to the right of the periodic table, the ionization energy increases. Here is a chart that displays this concept;

https://en.wikibooks.org/wiki/High_School_Chemistry/Ionization_Energy

Then, we learned about trends in electron affinity, which is how easy it is to add another electron to an atom. Electron affinity also increases as you move up and to the right of the periodic table. Here is a chart the exemplifies the trend:

https://en.wikibooks.org/wiki/High_School_Chemistry/Electron_Affinity

Lastly, we discussed trends in electronegativity, which is the tendency of an atom to draw electrons toward itself when chemically combined with another element. Elements with larger electronegativity tend to pull electrons to themselves when bonded to other elements. It increases as you go up and to the right of the periodic table too. Here is a chart that displays this trend:

http://chemteacher.chemeddl.org/services/chemteacher/index.php?option=com_content&view=article&id=91

Here are a couple links that further explain some of these trends:

Chemwiki

Practice Quiz