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. 

12-16-13

This week, we reviewed compound naming, learned about thermodynamically favoring, and redox reactions.

The compound naming was a hotpot quiz we had that was due on Monday. It consisted of all of the compounds that we need to know (which we probably should've already known). I thought the hotpot was really helpful because I'm still not completely comfortable with all of the compounds, but I've definitely got a lot of them down.

A thermodynamically favored process is called spontaneous. It means it's a process that proceeds without any assistance from outside the system such as, iron rusting at the presence of oxygen and water. A thermodynamically unfavored process is called non-spontaneous. It means that a process requires assistance from outside the system in order to induce change such as, water does not freeze at 15 degrees Celcius. A process that is thermodynamically favored in one direction is not favored in the other direction. Exothermic processes are typically thermodynamically favored because nature tends to favor processes that cause a reduction in energy. In an exothermic reaction, the bonds in the products contain less energy than the bonds in the reactants and the excess energy is released as heat. Endothermic processes can be thermodynamically favored - evaporation and dissolving soluble compounds are thermodynamically favored. The first law of thermodynamics is the energy contained in the universe is constant. The second law of thermodynamics is the entropy of the universe is constantly increasing. If the combination of entropies is positive, the reaction is favored. Heat in an exothermic reaction leaves the system, and heat goes into a system in endothermic reactions. If enthalpy (H) is negative and entropy (S) is positive then the reaction is spontaneous at all temperatures. If enthalpy is positive and entropy is negative then the reaction is non spontaneous at all temperatures.

A reduction reaction is sometimes called a redox reaction. It's a reaction where electrons are transferred and one of the substances gets oxidized (loses electrons) and the other substance gets reduced (gains electrons). Most reactions are redox reactions except for double displacement and acid base.
I'm excited for the lab on Tuesday. I would give my understanding this week about an 8. For some of the  I understand the equations and the worksheets we worked on this week but I just think I need to work a little more on knowing which equation to use and when. I sometimes get the right answer to the question, but I have no idea how or why I came to that answer so I definitely need to look at that. I know I have a lot to study for the test this week.

11-11-13

This week we learned about vapor pressure, lattice energy, and reviewed for the test this Tuesday.

Gas particles hit surfaces with a certain amount of force. They all have mass and must stop when they hit a surface (which is acceleration). The more gas particles, the more force. Vapor is the gas state of a liquid at room temperature, and it exhibits a pressure just like any gas. As temperature rises, the fraction of molecules that have enough energy to escape increases (warmer liquid evaporate faster). The more molecules that escape, the higher the pressure they exert. A liquid reaches boiling point when the temperature at which it's vapor pressure equals atmospheric pressure. Vapor pressure increases as temperature increases because more and more molecules at the surface have enough kinetic energy to escape the surface. At higher altitudes, water boils at a lower temperature so cooking and baking lengths are longer. The higher the boiling point, the lower the vapor pressure and the lower the boiling point the higher the vapor pressure. Vapor pressure decreases with molecular weight, but boiling point increases. If the intermolecular forces increase, vapor pressure decreases.

Lattice energy is the energy required to completely separate a mole of a solid ionic compound into its gaseous ions. The energy associated with electrostatic interactions is led by Coulomb's law. Lattice is periodic and predictable because charge and ion size are periodic in nature. It increases with the charge of ions. It also increases with decreasing size of ions. As lattice energy increases, so does melting point. Smaller ions lead to increased lattice energy. Greater charge also leads to increased energy, and the effect of charge is greater than the effect of distance.

We also did an activity where we tested the conductivity of substances with LED conductivity testers that light up when a substance or solution is conductive. Through this activity we determined that steel was the only one that conducted electricity because the electrons are loosely held due to metallic bonding. Water, acetone, ethanol, and nonane were all poor conductors because they are covalently bonded, making the electrons unable to move around. Another activity we did was identifying six unknowns by comparing surface tension and viscosity. My group was fairly close and predicted 4/6 of the substances.

I'd give my understanding this week about an 8. I definitely understand lattice energy and surface tension and how they affect bonding and was able to explain it to my classmates when they asked. I'm still not quite totally sure about everything on vapor pressure and I especially noticed while doing the task chains. The easiest part to me is the water phase change diagram and identifying what's happening in the diagram at a certain point and when a change is endothermic or exothermic like in the task chains. 

11-4-13

This week we began learning about intermolecular forces, and did a POGIL on water.

Molecules attract each other, and the force of attraction increases as intermolecular distance decreases. In liquids, molecules are very close to each other and are constantly moving and colliding. In a gas, molecules are much further apart than in a liquid. Boiling points and melting points are largely determined by intermolecular interactions in the liquid. As molecular weight increases, intermolecular forces get much stronger as well. Intermolecular forces are also much weaker than intramolecular bonds. London dispersion forces are the weakest intermolecular force. They exist in every molecule. The larger the molecule the larger the polarizeablility of the molecule. Dipole induced dipole are the next weakest. It occurs in a molecule when it has a very small dipole moment. Next is dipole dipole. The strength of a dipole dipole interaction on the dipole moment and how closely the molecules approach one another. In a solid, molecules are held close together in a regular pattern by dipole dipole forces to minimize repulsions and maximize attractions. Dipole-dipole forces only occur if the molecule is polar. The strongest of the intermolecular forces is hydrogen bonding. Hydrogen bonding can only occur with Nitrogen, Oxygen, and Fluorine. These intermolecular forces are all called van der waals forces.

In the water POGIL we learned that covalent bonds occur in a single molecule of water. These bonds are intramolecular. We used the femto beaker of water and the molecules with magnets to represent water. Unlike other molecules, when water freezes the volume increases. When pressure is applied to ice, the volume decreases and it becomes liquid. When you apply pressure to ice, the structure breaks and you melt ice due to pressure, so you're able to skate on ice.

We also learned about a fifth type of force - ion dipole interactions. The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents. If cation-anion attractions are stronger than ion dipole attraction, the compound will not be soluble.

I'd give my understanding this week about an 8. I definitely understand the different types of intermolecular forces and how melting and boiling point goes up as molecular weight goes up. I also understand that the higher molecular weight means more polarizability, which means a higher boiling point. I understand that every molecule has London dispersion forces. I'm still a little unclear on how to determine which gas is more soluble, but I think if I go over the Powerpoints and lectures I'll be able to understand it.

10-28-13

This week, we started out with review for our test. We asked a lot of questions, especially about the Lecture Chemical Bonding packets, and used class time to review.
After taking the test, it was Mole Day. We had really good cookies and hot chocolate.

  

We received a packet on Paintball and wrote about hydrogen bonding and polarity. We learned about how water’s polarity is due to the differences in electronegativity between oxygen and hydrogen. In water there is a region of partial negative charge on the side of oxygen, and a partial positive charge on the side of hydrogen. The molecules shape of bent and the polar bonds make the molecule polar overall. Hydrogen bonds occur when a hydrogen atom attaches to a small and highly electronegative atom, in this case Oxygen, in the vicinity of an atom with nonbonding electron pairs. Hydrogen bonds are the strongest of the intermolecular forces (but not as strong as covalent or ionic bonds). Hydrogen bonds are about 1/15th the strength of a covalent bond. The hydrogen bonds in water are what hold the molecule together.

We then began learning about ionic bonds, which was mostly a review. Ionic bonds are formed between two atoms when the atoms involved transfer one or more electrons to produce two charged species - positive (cation) and negative (anion). Atoms with loosely held electrons tend to form positive ions, but those who can hold additional electrons relatively strongly tend to form negative ions.

We learned about metals as well. Some properties of metals are that they have a shine or luster, can conduct heat and electricity, they're ductile, and they are malleable. Nonmetals do not have these properties - they're typically poor conductors of heat or electricity, and they aren't malleable or ductile. Electronegativity is much lower for a metal than for a nonmetal as well. In metals, the bonding is different from both covalent and ionic bonding. The electrons in their bonds are localized meaning they either are shared by a pair of atoms or they are associated with one of the two species involved in the bonding interaction. Valence electrons on a metal atom are shared with many neighboring atoms, not just one. These valence electrons are delocalized. The force of attraction between the positive metal ions and the sea of mobile negative electrons forms a metallic bond that holds these particles together.

This week, I'd rate my understanding of our topics at about a 9. It's mostly review with the ionic and covalent bonding and I feel like I understand the metals so far. I was able to help some of my classmates at my table especially with the Ionic Bonds POGIL because most of it was information that I already knew. Although, I was surprised when we were looking at the model of NaCl, that the atom for Na was the smaller particles and not actually the bigger ones. So far, I'm enjoying this topic and I hope I do better on this test than I did on the last one.

10-21-13

This week, we started out with a lecture quiz on sigma and pi bonding, then learned about hybrid orbitals, and had time in class to work on our WebMO.

Sigma bonds are characterized by head to head overlaps and cylindrical symmetry of electron density about the internuclear axis. Pi bonds are characterized by side to side overlaps and electron density above and below the internuclear axis. Single bonds are always sigma bonds because sigma overlaps are greater, which makes a stronger bond and more energy lowering. Double bonds contain one sigma and one pi bond. Triple bonds form with 2 pi bonds and 1 sigma bond. Resonance structures are used to more accurately reflect the structure of the molecule or ion.

We also learned about hybrid orbitals. The electron pairs around the central atom in a molecule are said to be in a set of orbitals (domains) that are hybridized from the usual set of atomic orbitals (s,p,d) for the atom. If one s-orbital and one p-orbital are used-2 hybrid orbitals called sp hybrid orbitals are made. With sp hybrid orbitals there are 2 hybrid orbitals in each set, one s, and one p, 180º between orbitals, and they are linear. With sp2 hybrid orbitals there are 3 hybrid orbitals in each set, one s, and two p, 120º between orbitals and they are trigonal planar. With sp3 hybrid orbitals there are 4 hybrid orbitals in each set one s, and three p, 109.45º and they are tetrahedral. 

We worked a lot on our WebMO diagrams making each of the 13 molecules on the website, finding their angles, dipole moment, and electrostatic potential map. 

My understanding of the two new topics that we learned this week is probably a 9.5. I definitely understand sigma and pi bonds, and I understand hybrid orbitals and could teach them to my classmates if needed. For the test on Tuesday, throughout the week I felt very lost and confused with each of the subjects we had previously learned (such as molecular shapes, bonding, and polarity), but I feel that after doing the WebMO project and the VSEPR Balloon lab report I really understand more of it. The hotpot quizzes really helped, and I took a lot of screenshots of the questions I didn't understand. After studying and working on the projects this weekend I'm beginning to feel a little bit more comfortable about the test on Tuesday. I think that the more I study Monday night, I'll be completely ready for the test on Tuesday....hopefully.