Weekly Blog Post 2/3-9

This week we were learning equilibrium. We spent most of the week either doing calculations or doing ConcepTests, both of which are very helpful. Equilibrium is a fairly simple concept, and for me hasn't been too hard to understand. LeChatalier's principle is a fundamental principle when it comes to equilibrium, and it says that if you change a reaction in some way, it will always try to come back to equilibrium. For example, if you have the reaction A+2B<--->3C, and you add more A to the reaction, the reaction will shift towards the product to balance it. The same works for pressure. If the reaction was put under higher pressure, the pressure wouldn't shift in this case because there are equal numbers of moles on each side of the equation. However, if there were four moles of C rather than three, the reaction would shift towards the reactants because there would be a higher pressure in the products, C, that needed to be balanced out. The unit K tells us the equilibrium constant for a reaction, and can either be Kc or Kp depending on whether we are measuring pressure or concentration of the reactants/products. To calculate these we would divide the concentration/partial pressures of the products by those of the reactants at equilibrium to find the value for K. Here is a good website for calculating K. There is also the letter Q which represents the K value for any specific point during a reaction not at equilibrium, calculated in the same way. 

(I accidentally pressed the quote button and that's why it's indented like this... don't know how to fix it. Sorry about that.)

We also did a great demo with a syphon and two graduated cylinders that visually represented the equilibrium process very well. When more of the liquid was added to one cylinder, it went through the syphon to the other "side" and came to equilibrium when the two water levels were at the same height. The demo was pretty helpful in explaining the basics of equilibrium.

I think my grasp of the concepts we learned this week is pretty solid, but I have a feeling the hard parts are just beginning so I need to stay on top of the material we have now so I don't fall behind for the next week. Other than that, I have really enjoyed equilibrium.

Last week we began our gas law unit. We started with ideal gasses and how they behave. We worked a lot with the ideal gas law, PV=nRT where P= pressure in atmospheres, V=volume in litres, n=number of moles, R=.0821 or the universal gas constant, and T=temperature in degrees kelvin. Derived from this ideal gas law equation are five other gas laws, named after the scientists who contributed to them. Gay-Lussac's law states that pressure and temperature are directly related through the relationship P1/T1=P2/T2. Charles' law states that volume and temperature are directly related through the relationship V1/T1=V2/T2. In Boyle's law, volume and pressure are inversely related, P1V1=P2V2. Dalton's law says that you can add partial pressures to get the total pressure. Avogadro's law states that volume and number of particles are directly related, V1/N1=V2N2. If you combine Gay-Lussac's Charles' and Boyle's laws, then you get the combined gas law, (P1V1/T1=P2V2/T2) which is good to remember because there's no need to remember the individual other laws, just cover up the constant and you have the equation you need.

I understand ideal gasses pretty well, sometimes the math can be a bit confusing, but only the stoich parts.

Real gasses take into account the attractiveness of particles to each other, and also their volume. If the particle is bigger, it will make the total volume bigger and also it will be more polarizable, making the pressure go down due to the particles pulling themselves from the walls. The volume for ideal gasses is adjusted up by a and the pressure for ideal gasses is adjusted down by b due to these phenomena.

I understand gas laws pretty well, but not perfectly. I still have work to do. My major problem with any test has seemed to have been the math part, so I need to get some extra practice with that tonight and tomorrow.

Last week we had a major focus on entropy of reactions and how entropy of reactions work. This week was probably the most difficult to understand so far. Entropy is a difficult concept for me to understand, especially with all the equations needed to calculate values. The basic definition of entropy is the number of distinguishable micro states in a given system. Evaluating if entropy increases or decreases in a reaction is relatively easy, if the number of particles or the energy of the particles increase, then the entropy increases. Entropy always increases in the universe.

In addition to entropy, we also worked on several other thermodynamic principles, one being Gibbs free energy. This free energy, as it is called, is energy in a reaction that can be used to do work on the system or for the system to do work on the surroundings. The equation for Gibbs free energy is ∆G=∆H-T∆S, where ∆H is enthalpy, T is time, and ∆S is entropy. Enthalpies and entropies are usually given in a table, but can be calculated using this equation if the free energy was known I suppose. This ∆G, or free energy, can be used to determine if a reaction is thermodynamically favored or not. If ∆G is negative, then the reaction is favored and will "go" without any outside energy needed. For example, dry ice sublimes at room temperature because it is favored. But water does not boil at room temperature because the ∆G of that reaction is positive.

Another thermodynamic principle is Hess' law. Hess' law states that in a reaction, the sum of the enthalpy of the steps of a reaction can be added up to to equal the total enthalpy of the reaction. My understanding of this is that if you add up the bond enthalpies in each different step and them cancel out the ones that are on both sides and eventually you will come to the total heat of reaction or the enthalpy of the reaction. Here's a link to a site on Hess' law.

Overall, my understanding of this unit so far is not where I would like it to be. Hess' law is confusing, although it seems simple. My major problem with all of these principles is knowing when to apply them. I think all I need to do is work more problems and I'll know whats going on. I think the lab we are doing next week will help a lot.

Weekly blog 11/4 thorugh 11/10

This week, we learned about vapor pressures, lattice energies, boiling points, and how intermolecular forces affect each of these. The higher the boiling point of a substance, the lower the vapor pressure. This means that the more intermolecular forces in a substance there are, the lower the vapor pressure in that substance. Lattice energy increases as ion size decreases and charge increases. This is a result of the coulombic relationship in lattice energy. Lattice energy is actually correlated with melting point, because ions in a lattice are harder to break from that lattice as lattice energy goes up. This means that when there is more lattice energy, the melting point rises. Here is a link to some good review on lattice energy.

On friday we applied all of these concepts we have learned in an activity where we had to determine if specific substances were or were not conductive, and also we had to identify unknown elements based on their properties. The conductivity test was fairly simple; we put a probe in each of the substances and it told us whether or not it was conductive. We discussed why some substances were conductive only when in a liquid state, like salt. Salt breaks into ions which can then conduct electricity when in liquid form. As for the identifying unknown elements, it was fun. It was interesting to see how some molecules were more viscous than others and to be able to tell which molecule was which just based on things like solubility, viscosity, and vapor pressure. Even though they all had virtually the same viscosity except glycerin, we could still figure it out. The key was testing the solubility with water.

My understanding of what was going on in class is pretty good, I think that I could still understand better why some things are conductive as a liquid but not as a solid. It seems counterintuitive to me, why something that has less density could be more conductive than something with more density. However, I think I'll be able to figure it out. Hopefully Dr. J will have another chat room open on monday night, because I found that enormously helpful when studying for the test last time.

Weekly Blog 9/28 through 10/3

We began this week focusing on metallic bonding. We learned that in metallic bonding, there is essentially a cluster of cations with a "sea of electrons" flying around the cations, not associate with any particular atom. Also, alloys are made in two different ways: interstitial or substitial. Interstitial is when smaller atoms fill the space in between larger atoms, making the alloy more dense, and substitutional alloys are where two similarly sized atoms are mixed, which does not increase density. Here, substitutional is on the left and interstitial is on the right.

Alloys are essentially solutions of metals.

We also learned about the intermolecular forces, or IMF's, that exist such as hydrogen bonding, dipole dipole bonding, and LDF's, or London Dispersion Forces. These are all van der Waals forces. Ion-dipole bonding can also happen, but this is not a van der Waal force. Ion-dipole forces are the strongest, then hydrogen bonding, then dipole-dipole, then the LDF's (induced dipole-dipole and induced dipole-induced dipole). We learned that LDF's are present in all molecules, polar, non-polar, ionic, everything. They are, however, very weak and don't compare much to hydrogen bonding and ion-dipole bonding. Here is a good site about LDF's.

Another thing we learned was about the hydration spheres of ions in a solution, We saw that if you dissolve K+Cl- in water, the K and Cl will separate and get surrounded by water molecules. The water molecules will align themselves so that the positive end is towards the negative ion and the negative end is towards the positive ion. The sphere will also include hydrogen bonding because of the two hydrogens in water bonded to the oxygen.

I think I have a good understanding of the ideas we covered this week. I'm looking forward to seeing what role entropy plays in these molecules and also learning more about what entropy really is. I've heard of it before, but I get the impression that I don't know anything significant about it.

Weekly Blog Post 10/20 through 10/27

On Monday, we began the week by reviewing for the exam on bonding. We reviewed amongst ourselves and went over a few problems from the review worksheet with the class. This was very helpful in getting to understand the concepts. That night, Dr. J put up a chat room in which we could ask questions of other students as well as Dr. J, which was extremely helpful. Being able to ask people questions you had rather than trying to figure it out yourself is super beneficial. Plus, it's nice to answer other people's questions to boost your confidence.

The next day we took the test, which was not extremely difficult but it had it's challenges. It was mostly conceptual, and I am fairly good at picturing things in my head so I did OK on it. I definitely mastered the concepts better than I thought I would be able to. 

Wednesday was mole day!

Dr. J made some delicious cookies and we had some hot chocolate while we talked for the whole class. We were assigned an essay about paintball and the role of polarity and hydrogen bonding in the paint balls and how scientists have made the paint water soluble without using water in the paint. 

We then started to learn about ionic bonding, which was an extended review of what we began to learn last year and had some more of it over the summer. It's not very difficult, but the POGIL we did covered a lot of material. I think I have a good understanding of ionic bonding so far. It's essentially just knowing what the tendency of each atom is; to become positive or negative. Just know that metals bond with nonmetals in ionic compounds. 

Last, we began a worksheet on metallic bonding but didn't get too far. I'm still pretty unsure about the whole thing, considering I only got about three questions done on the POGIL we got. The lecture and lecture quiz will help my understanding. 

My understanding of the material this week is good besides the metallic bonding. I learned a lot over the weekend (or, rather came to understand a lot) about covalent bonding through studying and the chat room. Those are very helpful. 

Weekly Blog 10/7 through 10/13

This week was a very busy week. We started it off with learning about vsepr models and how molecules are shaped by the bonded and non-bonded pairs of electrons. The vsepr models demonstrated why water has an irregular shape and why other atoms have their shapes. It also taught us how you can figure out the shape of the molecule knowing only the mount of bonded and unbonded pairs of electrons and the number of atoms around the central one. Here is a diagram of some vsepr structures.

We also learned about formal charges, which are used in Lewis dot diagrams in order to denote the charges of atoms. The formal charges are a result of polarity in a molecule, where one atom has more pull on the shared electron than the other. This can affect the vsepr structure of the molecule because of the way bonds are formed when an atom is polar. A molecule can be non-polar even if some of the bonds are polar. If the dipole moments all add to zero, then the molecule will have a net dipole moment of zero and will be non-polar. Here is a picture of a molecule and it's dipole moments.

At first I didn't understand polarity and dipole moments at all, but after I watched the lectures several times it began to make more sense. I have a much better understanding now than I did when I first watched the lecture. I still think that I need a lot more practice with it, however. Also, whiteboarding the Lewis structures and figuring out the molecular and electron domains was extremely helpful and I now know what I'm doing with those. Overall, my understanding of the material is good, but it can definitely be better.