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.