10/21/13-10/25/13

This week in AP Chemistry we had the third test of the year. This exam covered the topics of covalent bonding, electron domain geometry, as well as and molecular geometry. To study for the test, I had to complete a task chain. We also went over a packet of review worksheets in class on these subjects. At home, I did the task chain multiple times and completed extra problems on the review worksheets. On Tuesday, I took the test and got an 80 percent on the multiple choice and a 98 percent on the free response. I was very satisfied with my performance on the free response but was a bit disappointed with my multiple choice test. I felt I had a fairly good grip on this subject matter (better than the material on the stoichiometry test in fact), and yet I got a score that I was not very satisfied with. Considering the fact that I have gotten relatively iffy scores on the last two multiple choice exams, I've surmised that I need to pay more attention to the multiple choice portions on tests and make a better effort to practice the types of the questions that may be found on them. All in all, I was a bit irked with my performance on the multiple choice. However, because of my results on the short answer and the fact that the multiple choice did not significantly affect my grade, I found some satisfaction in my performance on this test.

The day after the test was mole day. We were assigned an essay on the subject of polarity and hydrogen bonding as it related to an article we were given about paintball. I found the assignment to be difficult at times, as I found the prompt to be a bit vague. As only one section of the paintball article really related to polarity and hydrogen bonding, it was hard to find excerpts in the writing that supported the points in my essay. With that being said, I did learn about hydrogen bonding and how polarity relates to solubility. Hydrogen bonding is when a partially positive hydrogen atom bonded to an electronegative atom is attracted to a partially negative electronegative atom in a separate molecule. Unlike covalent bonding, there is no physical bond present, but there is attraction between the two atoms.

A hydrogen bond, represented by the dashed line.

After these two assignments, I was reintroduced to the topic of ionic bonding via a POGIL that I completed on class. Ionic bonds are bonds between a metal and a nonmetal. I learned that the size and individual charge of each ion has a profound effect on the strength of an ionic bond. The stronger the bonds in an ionic compound, the higher the melting point. For the most part, I felt I had a decent understanding of this topic. I was a bit confused about the sizes of cations and anions.

The final topic we covered this week was metals. Metals are elements that conduct heat an electricity well. They are typically malleable and ductile. They exhibit several other unique physical properties, such as luster. Metals bond together through metallic bonds. In metals, electrons called conduction bond electrons form a "sea" of electrons around atoms. These conduction bond electrons are d-orbital electrons that are promoted into the outer p-orbitals of atoms. Conduction electrons are the reason that metals conduct heat and electricity well. The more conduction electrons, the higher the melting point of the metal and the harder the metal is. Most ordinary uses of metals involve alloys. Alloys are mixtures of elements that have the properties of metals. There are two main types of alloys: substitutional alloys (solute particles take the place of solvent metal atoms) and interstitial alloys (solute particles fit in holes in between solvent metal atoms). To help learn about metals, I completed a POGIL on metals on Friday and watched a lecture on metals this weekend. I found the subject straightforward and easy to understand.

    

10/14/13-10/18/13

This past week in AP Chemistry was the most confusing so far. I was introduced to the concept of hybridization and sigma and pi bonds.

In covalent bonds, electrons can only be shared if the electron orbitals of two atoms overlap. Hybridization occurs when these orbitals combine to form new, hybrid orbitals. For example, the s and p orbitals combine to form an sp orbital. There are three widely accepted forms of hybridization: sp, sp², and sp³. Hybridization is a theory; whether or not it works beyond the second period is questioned. Additionally, hybridization is central to valence bond theory, a.k.a. sigma and pi bonding, which is explained later. 

sp hybridization.

I did not understand the concept hybridization well, which would explain my rather lackluster explanation. I do know that you can determine the hybridization of a molecule by simply counting the number of electron domains around the central atom. If there are two electron domains, the hybridization is sp. If there are three electron domains around the central atom, the hybridization is sp². If there are four electron domains, the hybridization is sp³. Whether hybridization occurs beyond that is hotly contested within the scientific world and does not concern me. However, conceptually I do not understand hybridization beyond the notion that orbitals can combine to make "hybrid" orbitals. To cover these concepts, I watched a lecture on hybridization and completed the Hybrid Orbitals POGIL for homework. While what I know may be enough to get by on the test, I still would like to understand the concept of hybridization more.


I found the concept of sigma and pi bonds to be more straightforward. Sigma bonds are bonds with head to head overlap and cylindrical symmetry of electron density about the internuclear axis. Pi bonds are bonds with side to side overlap. Electron density is above or below the internuclear axis. In a single bond there is one sigma bond. In a double bond there is one sigma bond and one pi bond. In a triple bond there is one sigma bond and two pi bonds. I was confused about this concept until I completed the task chain quiz that involved sigma and pi bonds. When I got an answer wrong regarding a sigma and pi bond, the information window that popped up gave a good explanation of sigma and pi bonds. To cover sigma and pi bonds, I watched a lecture on the subject. This website also gave me some clarity on the subject.

There is a pi bond between the blue p orbitals, as there is side to side overlap between them. There is a sigma bond between the green orbitals, as there is head to head overlap between them. 

Throughout the week in class I worked on a lab/activity involving the program Web MO in which I modeled molecules from the VSEPR Theory lab with my table group. Using the program, I determined the bond angles, polarity, dipole moment, and more for several molecules. The activity really helped me better understand the concepts of bond and molecular polarity, as well as how unbonded electron pairs affect the shape of molecules.

10/7/13-10/11/13

This week in AP Chemistry we continued work on Lewis structures and VSEPR theory. Additionally, we learned the concepts such as formal charge and polarity.

Near the beginning of the week, we finished up the balloon/gumdrop activity, which pertained to VSEPR theory. While Lewis structures help to understand the composition and covalent bonds of molecules, VSEPR theory (short for "valence shell electron pair repulsion") allows us to predict the shape of molecules. There are two main categories in VSEPR theory: electron domain geometry and molecular domain geometry. In predicting a molecules electron domain geometry, we assume that electron pairs are placed as far apart as possible. Electron pairs are referred to as electron domains; one electron pair is equal to one electron domain. Double and triple bonds also only count as one electron domain. To determine electron domain geometry, one first counts the number of electron domains and then chooses the corresponding shape. Pictured below are several electron domain geometries.  

Different electron domain geometries. For two electron domains, the shape is linear. For three domains, the shape is trigonal planar. For four domains,  the shape is tetrahedral. For five domains, the shape is trigonal bipyramidal. For six electron domains, the shape is octahedral.  

Molecular domain geometry is slightly different. It is defined by the positions of only the atoms and not the nonbonding electron pairs. Here are some examples of molecular geometries:

When there are two electron domains, the following molecular geometries are possible:

• AB2 - linear

When there are three electron domains, the following molecular geometries are possible:

• AB3 - trigonal planar
• AB2E - bent
• ABE3 - linear (this is a rare occurrence)

When there are four electron domains, the following molecular geometries are possible:

• AB4 - tetrahedral
• AB3E - trigonal pyramidal
• AB2E3 - bent
• ABE3 - linear

When there are five electron domains, the following molecular geometries are possible:

• AB5 - trigonal bipyramidal
• AB4E - seesaw
• AB3E2 - t-shaped
• AB2E3 - linear

When there are six electron domains, the following molecular geometries are possible:

• AB6 - octahedral
• AB5E - square pyramidal
• AB4E2 - square planar

To help learn the concept of molecular and electron domain geometry, I completed a POGIL in class and made models of molecules using gumdrops and toothpicks. I also watched a lecture on VSEPR theory. While I feel I have a decent understanding of this subject, I feel I could use more practice determining the molecular geometry and electron domain geometry. I sometimes get confused with differentiating between the two kinds of geometries, as the electron domain geometry involves lone pairs of electrons and the molecular geometry does not.

We also covered the concept of formal charge in class.  Formal charge is equal to the number of valence electrons in a free atom minus the number of bonds plus the number of nonbonding electrons. The sum of each individual formal charge of each atom in a molecule must be equivalent to the total charge of the molecule. If there are multiple structures for a molecule, the molecule with the lowest formal charge is preferred. If there is a negative formal charge, it should be on the least electronegative atom if there is a choice. If the central atom is in period three or higher, multiple bonds are possible if it will reduce the formal charge of the molecule. To help me learn formal charge, I had to view a lecture on formal charge. In class, I completed the Lewis Structures III POGIL and the Lewis Structures IV POGIL, both of which dealt with formal charge. I felt I had a relatively strong grasp of this concept, although I sometimes have trouble finding which Lewis structure has the lowest formal charge when there is more than one possible structure.

Finally, we touched on the concept of polarity in a lecture that was assigned last week. An molecule is polar when its electrons are not shared equally. Shared electron pairs in polar covalent bonds are not shared equally, while electron pairs in nonpolar covalent bonds are shared equally. Between the atoms of molecule, the greater the difference in electronegativity, the more polar the bond. Molecular polarity is possible and is calculated by adding up the individual bond dipoles. One thing I did not understand about polarity is the "dipole moment." From the lecture and from various sources online I was unable to gather a clear definition. I know it pertains to electrical charge and electrons, but I feel I need a concrete definition in my head to best understand what a dipole moment is.

9/30/13-10/4/13

This week in AP Chemistry, I continued to learn about Lewis structures. Over the course of the week I was introduced to the topics of bond order, resonance, hypervalency, and covalency. Bond order is the number of chemical bonds between a pair of atom. A single bond has a bond order equal to one, a double bond has a bond order equal to two, and so on and so forth. To help learn the concept of bond order, I completed the Lewis Structures Part 2 lecture at home and completed a POGIL on Lewis structures in class.

Resonance is a term used to describe a situation in which there are more than one valid Lewis structures for a given molecule. Initially I was unsure of the exact definition of resonance, but this website gave a nice concise definition. The Lewis Structures II POGIL and the Lewis Structures Part 2 lecture dealt with this topic.

Possible resonance structures of NO3. The double headed arrows indicate interchangeability (all three structures are valid). 

If there are multiple valid Lewis structures, the bond order is affected. This was a point of confusion for me. Initially, after watching the Lewis Structures Part 2 lecture, I came under the impression that the bond order was equivalent to the number of bonds divided by the number of possible structures for a molecule. However, after starting the POGIL on Monday I discovered that this was certainly not the case. I learned that bond order between two atoms was simply equivalent to the number of bonds present. I am still confused as to how resonance affects bond order. I have searched online to find an answer, but I have not found one that makes it clear to me.

Hypervalency is a situation in which there is more than eight electrons around an atom in a molecule. When making a Lewis structure, if you satisfy the octet rule for all of the outer and central atoms and still have extra valence electrons, distribute them about the central atom. This is only possible if the central atom is in period 3 or greater.

In the PCl5 molecule, the Phosphorus atom is hypervalent. It exceeds the octet rule and has 10 valence electrons.  

When considering hypervalency, size is important. The larger the central atom, the larger the number of electrons that can surround it. Expanded octet most often occurs when the central atom is bound to atoms with high electronegativity such as Fluorine, Chlorine, or Oxygen. To learn the concept of hypervalency, I watched the Lewis Structures Part 3 lecture. The concept was also part of the VSPER Theory POGIL we started in class on Friday. I found most of this topic easy to understand. With that being said, I thought the portion in the lecture that explained how the expanded octet is possible was very confusing.

Covalency was the final new topic we covered this week. In the lecture on Covalency, I learned that molecules with covalent bonds stay together because the force of repulsion between protons is weaker than the force of attraction between protons and electrons. Electrons are said to be "paired" when they have opposite spins and enter the domain of the other. Additionally I learned that as bond order increases, strength increases.

Overall I though I understood most of the concepts covered this week well, with the exception of the relationship between resonance and bond order as well as the part of hypervalency that I explained earlier.

In addition to learning these concepts, we did a stoichiometry-related lab in class on Wednesday and Thursday. The purpose of the lab was to find the percentage mass of copper in a brass screw. In the lab, we took a brass screw and dissolved it in nitric acid under the fume hood, which produced both toxic NO2 gas and a liquid, Cu(NO3)2.

The brass screws react with nitric acid to produce a blue liquid, Cu(NO3)2, and a brown gas, NO2.

Then, half of the groups made dilutions with the Cu(NO3)2 liquid, measured their absorbance, calculated concentration, and made a calibration curve. The other groups (mine included) tested the visual comparison method of calculating concentration. In this process, you filled two beakers with solution. One beaker had stock solution and the other had the product of the brass screw/nitric acid reaction, known as the "mystery solution". The beaker with mystery solution was filled until its color matched that of the stock solution. The depth of each solution was then measured and then the concentration of the mystery solution was calculated with the equation:

(Molarity1)(Depth1) = (Molarity2)(Depth2)    

The lab was fairly straight forward. I think my work on it is an improvement over my work on the previous lab. Now that I have a little bit of experience doing labs, I am more comfortable and the work in my lab notebook is much tidier. I am unsure when the lab is due and would like to know so I know when to start working on the post-lab questions.