2-10-14

This week we learned about equilibrium calculations and thermodynamics and equilibrium.

We learned Le Chatelier's principle which is, "if a system at equilibrium is disturbed by a change in temperature, pressure, or the concentration of one of the components, the system will shift its equilibrium position so as to counteract the effect of the disturbance".

Kc=[C]c[D]d      [A]a[B]b

This equation is used to find the equilibrium constant of molar concentration in an equation. 

Kp=(C)c(D)d(A)a(B)b

This equation is used to find the equilibrium constant of partial pressures of each gas in an equation. 

Kp=Kc(RT)Δn

This equation is used to convert from Kc to Kp. The delta n means the number of moles on the product side, minus the number of moles on the reactant side. Q gives the same ration as K, but for a system that is not at equilibrium. Q is calculated using the initial concentrations of the reactants and products. If 

Qc < Kc, there is more reactant and less product in the initial conditions than at equilibrium. The reaction will then move towards the products, so to the right. If Qc > Kc, there is less reactant and more product in the initial conditions than at equilibrium. The reaction will then move towards the reactants, so to the right. If Q = K, the reaction is already at equilibrium under initial conditions so it doesn't shift. 

Another way used to calculate equilibrium is using a RICE chart. RICE stands for Reaction, Initial, Change, and at Equilibrium. With some reactions, they may end up getting into cube roots. However, if x is smaller than 5% of initial concentrations, it can be left out and the quadratic can be avoided. 

A reaction is favorable when ΔG° is less than 0, and is called exergonic. A reaction is unfavorable when ΔG° is greater than 0, and is called endergonic. ΔG° for a reaction measures the difference between the free energies of the reactants and products when all components of the reaction are present at standard-state conditions. When it's negative, the reaction would have to shift to the right converting some of the reactants into products. There is only one value of ΔG° at a given temperature, but an infinite number of values of ΔG. The smaller the value of ΔG°, the closer the standard state is to equilibrium. The larger the value of ΔG°, the further the reaction has to go to reach equilibrium. 

This week I'd give my understanding about an 8.5. I was able to do the lecture quizzes without any confusion, and although I missed Monday I completed the equilibrium I worksheet. This unit I've also paid a lot of attention to the lectures and in class, and I'm really trying to do as well on this test as I did on the last one. I think the worksheets have really been helping understand the concepts and Concept tests as well. I hope I do even better on this test than on the Gases unit because I think I really understand and like the topic of equilibrium. I find it interesting and the demonstrations at the beginning of class are intriguing and useful to learning the topic. 

1-20-14

This week we finished up our gasses unit by learning about ideal and real gasses.

We started with learning about the Kinetic-Molecular Theory (KMT) which is a model that aids our understanding of what happened to gas particles as environmental conditions change. A measure of the average kinetic energy of atoms and molecules in a system is temperature. The Kelvin scare is proportional to this, so when kinetic energy doubles the Kelvin temperature doubles. The main parts of the KMT are - gases consist of large numbers of molecules that are in random motion, the combined volume of all the molecules of the gas is negligible relative to the total volume in which the gas is contained (even though they have different masses), masses are considered volume less points of mass (they have mass but no volume), and there are no attractive or repulsive forces between gas molecules. Collisions between gas particles are elastic and kinetic energy is conserved. Average kinetic energy does not change with time, as long as temperature stays constant. Molecules with lower mass go faster than those with a higher mass.

Average kinetic energy of molecules is also proportional to the absolute temperature. Effusion is the escape of gas molecules through a tiny hole into an evacuated space. A smaller molar mass means faster effusion. Diffusion is the spread of one substance throughout a space or throughout a second substance. The average distance between collisions is the mean free path. As pressure increases, mean free path decreases. 

In the real world, the behavior of gases only conforms to the ideal gas equation at relatively high temperature and low pressure. Real gases are not ideal because they have attractive forces and gas particles have real volume. In the equation, "a" means the interaction part of the molecules and "b" is the volume component of molecules. Attractions between gas molecules reduces pressure and pulls particles away from the walls where pressure is measured and they're not hitting as hard. IMF forces increase as particles get closer which leads to condensation at low temperature or high pressure. Condensation causes temperature to decrease and kinetic energy to decrease and can cause particles to stick together. Condesation also causes pressure to increase and causes higher rate of collisions. Gases  that are closest to ideal gases are non polar, have no hydrogen bonding, smaller molecules, high temperature, smallest LDFs, spaced far apart (low pressure), and have less volume. Gasses furthest from ideal gases have low temperatures, low kinetic energy and more time to intreat, high pressures, and the particle count is high. 

This week, I'd give my understanding of our new subject about a 9. I think that I understand the topic well enough that I was able to help my classmates with questions that they had. This unit I've really paid attention during class and watched the lectures really closely, which is really different from what I did last unit. I think that doing the Gases II worksheet really helped, as well as doing the Concept tests at the beginning of class. I hope that I do better on this test than thermodynamics unit, which I know I will because I've really been trying to understand everything in the lectures and in class.