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.

12/16/13 Weekly Reflection

Thermodynamics and its applications in chemistry has been the main focus of the week. We looked at  entropy, enthalpy, precipitation reactions, and redox reactions. We started with the internal energy of a system, knowing that it is the sum of the heat transferred into(+) and out of(-) the system and work done on the surroundings by the system or vice-versa. We examined this equation by looking at gas in a cylinder. Enthalpy was the next point of discussion. Enthalpy is the same thing as heat at a constant pressure. We learned how to find Enthalpy of a reaction using bond energies (BE bonds broken + BE bonds formed), using calorimetry (the energy change in the surroundings to find the energy change in the system), using enthalpies of formation (sum enthalpies of formation of products - sum enthalpies of formations of reactants), and finally using Hess's law (where the overall change in enthalpy is measured using several steps as opposed to one). In Hess's law, reactions can be flipped around but you must change the sign - a good example of the whole process can be found here. Entropy, however, was by far the most difficult concept so far. Entropy is defined as a statistical measure of the number of most probable distinguishable microstates available to the system. A high, positive number of distinguishable microstates is favorable or spontaneous. One of the ways to determine whether or not a system is favorable is to use the Gibbs Free Energy equation, which takes the enthalpy change and subtracts the product of the system temperature and the entropy change. If the result is negative, the reaction is favorable. Then we looked at precipitates (reactions where two solutions are added to each other and react forming a solution and a solid) and did a logic puzzle using our knowledge of precipitates to determine which compound was which. Lastly we looked at redox reactions, or reactions where atoms change whether or not they take or receive electrons (are either reduced or oxidized).

While I understand the overall concept of entropy, I don't really understand how it works or why it is relevant. How is the entire universe working against entropy?Or is it actually not? And how is it that when energy is released from a system, that system loses entropy? How are entropy and energy connected? I also don't understand the purpose of Gibbs Free Energy. Is the rest of the energy locked in bonds? What does the energy in a system that can be used to do work have to do with Entropy? I understand Enthalpy well enough now, and I definitely understand precipitation reactions. I really liked the logic puzzle we did, and it helped to solidify and apply the knowledge we already had on precipitation reactions. I tried hard to help my group to understand the material to the best of my ability(excluding the material I didn't understand), and I strove to completely understand as much material as possible. Learning this section now, during the start of winter has put a brand new spin on making hot chocolate, especially the coffee cup calorimeters we will be using for the lab - I keep wondering how much energy my hot chocolate is releasing and how I could calculate it.

11/11/13 Weekly Reflection

This week we started by delving back into IMF's, but this week we looked more at how they can affect and how they explain phenomena, such as surface tension, boiling point, evaporation rate, and viscosity of a liquid. We learned about how cohesive and adhesive forces affect all of these things and how they are affected by IMF's. We also looked at solids, including the two main groups of solid - crystalline and amorphous- as well as a couple of groups contained in those two - Ionic crystals, Covalent-network solids, and Molecular solids. We then looked at vapor pressure. Zach had a really good definition of vapor pressure, saying that it was about how many gas particles were moving at a given time and with what force. Thus, as vapor pressure increased, more molecules were moving in a gaseous state at a given temperature, so the higher the vapor pressure, the closer the substance is to boiling. Basically, the higher the temperature, the more particles have enough energy to escape the liquid and hang about as a vapor. Then we worked with lattice energy, which is the energy required to separate a mole of a solid ionic compound into its gaseous ions. It explains why the reaction forming an ionic compound is exothermic, and it is related to the charge and size of the ions. On Friday, we used everything we had learned about vapor pressure to try to figure out which compound was which given five chemical formulas and five liquids in dropper bottles. We used surface tension, viscosity, and evaporation rate to determine which compound was in which bottle.

I had a couple of questions. Why is it that you can pump something up as far as necessary but it can't be pulled up without a state change? What does enthalpy mean and why is it relevant? There was a question on the task chain for the test that had a phase graph - what was up with that graph? Which line were we supposed to be focusing on?  I tried hard to participate in class this week. My group has changed, so my style of explanation has to shift as well, which has forced me to reexamine how I think about concepts. I understood everything fairly well in this section, although keeping the various relationships in order in my head was probably the most difficult part. I still need to work on my method of remembering which relationships correspond in which ways. This section has been fascinating, and I'm now trying to figure out why various substances have higher or lower surface tensions than others.

11/4/13 Weekly Reflection

Ionic compounds and metals were the first topic of the week. Specifically, we worked on metallic solids, alloys, and the electron sea model. Next, we dealt with van der Waals intermolecular forces, including H-bonding, dipole-dipole interactions, induced dipole-dipole interactions, and London dispersion forces. H- bonding is when hydrogen is in an extremely polar bond, causing it to be nearly entirely positive and attract other, negatively charged ends of molecules in which H-bonding is present. This is the most powerful of van der Waals forces. Dipole-dipole interactions are when molecules with dipole moments are attracted to the oppositely charged ends of other molecules with dipole moments. This is the next strongest. Next is dipole-induced dipole interactions, where the dipole moment in one molecule will cause the electron cloud of another molecule to shift, inducing a dipole moment in that molecule and causing some attraction. London dispersion forces (LDF's) are the weakest intermolecular forces and are present in every substance. Because electrons are constantly moving, there will occasionally be an instant when the electrons are shifted to one side of the molecule, causing a temporary dipole moment which is then passed around the substance. We spent the rest of the week white boarding and working with water models.  

I tried very hard to participate in the learning process this week because I understood all of the material very well. I especially liked the idea that the electron cloud around any given molecule is not fixed, and that they can all be on one side temporarily, spreading an induced dipole throughout the substance. I need to work more on interacting with my group - everyone in my new group is kind of quiet, and we don't know each other that well, so it's a little awkward, and I need to work to iron that out. I was curious about surface tension. Specifically, I was wondering why, if water has a stronger adhesive force with glass than a cohesive force with itself, things can rest on the surface tension of the water. Is it just because there is a different and weaker adhesive force between water and, for example, water striders than between water and glass? Also, in the PowerPoint on cohesive and adhesive forces you said that trees could only pull water up thirty-three feet. How are trees taller than this if they can only bring water up thirty-three feet? This section was fun because now I get to consider the possible cohesive, adhesive, and intermolecular forces operating all around me.  

10/28/13 Weekly Reflection

There were two major events this week (actually, there were three). First we had the covalent bonding test with a day of review beforehand. Then, it was mole day, when we got to eat mole cookies and drink hot chocolate, listen to the mole song which drove Jackson crazy, and write an essay on hydrogen bonding and polarity. Lastly, we had the AP pre-test and the next day started the ionic and metallic bonding unit. The difference in electronegativity between hydrogen and oxygen results in the electron of hydrogen being pulled slightly towards the oxygen atom. This causes the hydrogen to have a slightly positive charge and the oxygen to have a slightly negative charge. Thus, in water, the slightly positive hydrogen's are attracted to the slightly negative oxygen's of different molecules. This attraction is called hydrogen bonding. It is not an actual bond, however, and is thus notated by a dashed line. Ionic bonding is a bond between a metal and a nonmetal in which the nonmetal takes the metals atoms and the two atoms are then attracted by the difference in charge. In metallic 'bonding', the cations are arranged in a lattice structure with all of the valence electrons being fairly free to move about the substance. This explains why metals are such good conductors of electricity.

So far, I understand this new section on ionic and metallic bonding very well. I was wondering why metals are ductile and malleable, although I think we will probably cover that very soon. I was also wondering why this section was called, "May the force be with you", as it doesn't seem relevant yet. I kind of hope it will be some pun on something we learn later, so halfway through the section I will be able to look back at it and go, 'Hah! I get this now.' I tried to be very engaged and a part of the learning process this week, despite the fact that most of this week was testing. I understood everything about this section so far, although I said that about the hybridization too and that didn't turn out so good. Anyways, I need to work on studying, especially with old material. I need to refresh my memory with everything, and in my free time (ha ha) I should do practice problems. I really didn't know about the sea of electrons, so now whenever I touch a metal I get this weird sensation of 'oh my, this is actually vibrating at speeds too fast for my comprehension, and the valence electrons are just kind of floating around all over the place...this is a bit worrying'.

10/21/13 Weekly Reflection

Most of this week was spent on VSEPR models and WebMO. Hybridization and the test on bonding coming up next week were also covered. However, the VSEPR and WebMO report was the center of attention for this week. WebMO is a really cool website that a couple of chemists came up with that you can use to diagram whatever molecules you would like. You can change settings for comparative purposes, draw ball and stick and space filling models, show dipole moments, and examine molecular diagrams for all sorts of fun stuff. It is a program that costs money and I think it was originally intended for colleges, but Dr. Finnan got us hooked up (thanks again, by the way). To better understand the WebMO and get more practice with it, we did a report in which we have to give the molecules we made and a table of information about them, as well as a paragraph explaining why each of these molecules is how it is. We also spent some time on hybridization because so many people were confused by it. Hybridization, as far as I can tell, is something that occurs constantly for all atoms that it is capable of occurring for (unlike hybridization), and in these elements the s and p orbitals are combined to form a hybrid sp orbital.

I have a few questions from this week, and I will do my best to explain them as clearly as possible. I was curious as to why hybridization occurs in any period excluding the second. For example, why does sulfur hybridize, but selenium and phosphorus do not? I had a couple of other questions, but as I was writing them out I answered them by myself, so I decided against writing them. For example, I had been wondering about why hybridization even happens, and I realized that it was simply a way of explaining why all of the bonds in, for example, tetrahedral molecules exhibit the same properties. I started the week with a confident, complete understanding of hybridization, but then I realized that my understanding of it was wrong. That put a damper on things, because now I had to get rid of some of what I had learned before and relearn it, which is frustrating because then you are never quite sure what is new and what were your preconceived, incorrect notions. I think I figured it out though. I completely understand VESPR models, and I tried very hard to participate as much as possible. Throughout this whole hybridization reimagination process, however, I realized that I had to work on my turnaround time for new, contradictory ideas. It took me way too long to change my thinking quickly and effectively on that particular topic. I will definitely be considering that problem and how to fix it.