2/10/14 Weekly Reflection

Equilibrium was the main focus of this past week. Chemical equilibrium is when the reverse and forward reactions of a given reaction are occurring at the same rate. It is referred to as dynamic equilibrium because both the forward and reverse reactions are happening at the same time. The equilibrium constant is represented by Keq, which is a ratio of concentrations or partial pressures of products to reactants, and changes based on temperature. Keq can be represented in terms of concentrations, when it is written as Kc, or in terms of pressure, when it is written as Kp. Le Chatelier principle describes how systems will shift equilibrium after a change in temperature, pressure, or concentration in order to establish a new equilibrium. If, for example, temperature is increased, equilibrium will move either to the right or left depending on whether the reaction is endothermic or exothermic. The Keq value for the reaction will also change accordingly, either increasing or decreasing depending on the reaction. It is important to remember that enthalpy values describe the amount of heat energy going into the system, so a negative enthalpy value means exothermic and a positive enthalpy value means endothermic. Then Q was introduced. Q describes an instantaneous ratio between products and reactants and is a good indicator of what the reaction will have to do to reach equilibrium from where it was at the instant the measurement was taken. If Q<Keq, then the forward reaction will have to progress faster than the reverse reaction until enough product is produced to attain equilibrium. If Q > Keq, then the reverse reaction will have to progress faster than the forward reaction until enough reactant is produced to attain equilibrium. Then we talked about RICE charts, which are a fairly simple way to help find either equilibrium values or quantities or initial values or quantities, depending on what you are given. They are mostly important for organizing information. We finished with delta G, or Gibbs free energy. The sign of standard state Gibbs free energy tells us which direction the reaction will have to proceed in at standard state conditions to come to equilibrium. If it is negative, the reaction will have to produce more products. If it is positive, the reaction will have to produce more reactant. The larger the value of delta G at standard state, the farther the reaction has to go to achieve equilibrium.

This week we did a lot of white boarding, and Dr. Finnan did a demonstration of the NO2 <->N2O4 reaction and showed us what happened when you changed temperature and pressure. The main thing that I don't understand about this section is how exactly Gibbs free energy relates. How does Gibbs free energy affect equilibrium? Or more properly expressed, what does a change in Gibbs free energy show about the reaction? From my notes, I don't really understand why the reaction can't be at equilibrium at 273K and 1atm. Unless you change temperature and are trying to find equilibrium at that new temperature from Gibbs free energy at standard state. I tried to participate in class, and in some areas I believe I have been very helpful to my fellow classmates, but generally speaking I need to be more focused and work harder in class. Overall, I understand the material fairly well, but I need to focus more and try to figure out the relation of Gibbs free energy to equilibrium.

1/21/14 Weekly Reflection

This week was entirely about gases, ideal and real. Ideal gases don't actually exist, but they follow the kinetic molecular theory, which is a model of what happens to gas particles as conditions change. Kinetic molecular theory has a couple of main tenets. According to KMT, gases consist of large numbers of molecules that are in continuous, random motion, the combined volume of the molecules of the gas is negligible (i.e. gas particles have mass but no volume), attractive and repulsive forces between gas particles are negligible, kinetic energy is conserved, and the average kinetic energy does not change with time as long as temperature remains constant. All of the properties of ideal gases and their relationships are illustrated with the equation PV=nRT, where P is pressure, V is volume, n is the number of molecules, R is the gas constant, and T is the temperature. This is actually a combination of relationships, which are detailed in Boyle's law (PV=k), Charles' law (V/T=k), Gay-Lussac's law (P/T=k), and Avogadro's law (V/n=k). Also important when dealing with gases are the concepts of effusion and diffusion. Effusion is the escape of gas molecules through some material, while diffusion is the spread of gas molecules throughout a space or materials. However, both rates are dependent upon mass in the same way. We also discussed partial pressures and how they are related to total pressures and mole fractions (which are the relative percent compositions by mole of a single component of the mixture, represented as decimals.) Then we talked about real gases and how they deviate from ideal gases, which allowed us to segue back into our previous discussions of IMF's and polarity. Both are extremely valuable in understanding how real gases deviate from ideal gases. Real gases deviate the most from ideal behavior under high pressure and low temperatures because that is when particles are the closest and moving the slowest, thus allowing the IMF's to have the strongest effect.

We spent a lot of this week whiteboarding and doing concept tests, although we also spent some time messing around with liquid nitrogen, in order to better understand the concept of gases. I understand this section much better than I did Entropy and Thermochemistry. The most trouble I have int his section is remembering back to the stoichiometry. I also have some trouble with remembering the smaller relationships between specific things. I have been trying hard to participate in class, but without complete comfort in my knowledge, I find it hard to voice those ideas in class. I tend to understand the correct answer once it is explained, or even once it is identified, but until then I have trouble choosing the correct response or being able to explain my response. I don't have any concrete questions, but I certainly need more practice. After this section, I have been wondering a lot about the process of respiration and how exactly breathing works on a chemical level. It is fascinating.