Bond Polarity

For molecules to be polar, it means that there is a difference in electronegativity and that their dipole moments don't cancel each other out. This is caused by one atom exerting more of a force on the electron cloud than the other. An electric dipole is the unequal sharing of electrons within a bond. Partial charges are marked by delta plus and delta minus. This is shown below:

Source

If the only bond in a molecule is polar, the molecule is polar. If the bond is nonpolar, the molecule is nonpolar. If it is nonpolar, there is no difference in electronegativities, the dipole moments cancel, and they have a symmetric shape. In molecules with more than one bond, both the shape and bond polarity determines if it's polar or not. 

Here are a couple links to practice problems for what is explained above:

-Practice 1

-Practice 2

Shapes of Molecules

There are five shapes of molecules that we have learned about so far:

Tetrahedral:
Have four bonded entities around a central atom. Its shape and geometry are both tetrahedral. These molecules are not flat, symmetrical, non-polar, and have identical bond angles.

Trigonal Pyramidal:
Have three bonded entities and one lone pair of electrons on the central atom. Its shape is trigonal pyramidal, while its geometry is tetrahedral. These molecules have no symmetry and are polar.

Bent:
Have two bonded entities and two lone pairs of electrons on the central atom. Its shape is bent, while its geometry is tetrahedral.

Linear:
Have two bonded entities and no lone pairs of electrons on the central atom. Its shape and geometry are both linear.

Trigonal Planar:
Have three bonded entities and no lone pairs of electrons on the central atom. Its shape and geometry are both trigonal planar. These molecules are non-polar.

Here is a picture of a few of the shapes listed above, and an example of their Lewis Structure:

Source for above picture

Here is a website that I found that explains these shapes more in-depth, and even has practice games on it: Molecular Shapes

Molecular Model Lab

This past week in chem class, we went to the library (for the first time ever), to do a lab about molecular models. We were given 10 molecules, and using the "have, need, share" method, we had to determine their Lewis Formula. Here is a link to a website that I found that explains how to draw Lewis Structures: Lewis Structures. We also had to sketch their 3-D Geometry, after we built the molecule. This lab really really helped me better understand how to use/do the "have, need, share" method, which I think will be super beneficial when it comes time for the test. Also, I enjoyed this lab because the change of scenery was nice, and it was really fun getting to write on the dry-erase tables. I've never seen anything like that! Here is a picture of what the table looked like with all of our writing on it, and some of the pieces we used to build the models:

Electronic Structure Test Reflection

After taking the test for this unit on Friday, I am feeling very confident about it. I've never felt this good about a test in Chemistry! It's probably partially due to the fact that this is just an easy unit, but I think it is also because I studied a lot. I did all four online practice tests, printed out all of the other study tools and completed all of them before the quiz and re-did them before the test. I actually did not do very well on the quiz for this unit, so that also kinda freaked me out and made me realize that I needed to study more in order to redeem myself on the test. The only thing on the test that tripped me up a little bit was just one question having to do a math calculation, and I didn't know what formula to use on it. Other than that, I actually knew how to do everything! For the future, I now know that if I want to do well on a test, I need to study as hard as I did for this unit. I'm just hoping that the next unit will be as pain-free as this one was!

Periodic Trends

Here are the four periodic trends that we learned about last week in class:

Atomic Size: As you move down a group and from right to left on the table, the atoms tend to get larger. This is due to electrons being added to larger orbitals and shielding. Here is a picture that helped me better visualize what the sizes of each elements atoms are:

Source of above picture

Ionization Energy: As you move up a group and from left to right on the table, the ionization energy increases. This means that it needs more energy to remove an electron from a gaseous atom. Here is a website that I found to help explain what ionization energy is: Ionization Energy

Electron Affinity: As you move up a group and from left to right on the table, electron affinity increases. This is the ease with which an electron may be added to an atom, forming an anion. Some electron affinities can even be negative.

Electronegativity: As you move up a group and from left to right on the table, electronegativity increases. This is the tendency of an atom to draw electrons toward itself when chemically combined with another element. Here is a website that helped me to understand what this means: Electronegativity