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. 

10-14-13

This week, we started out by finishing up the VSEPR structures, then more Lewis structures, and Formal charges. We use formal charges to make correct Lewis structures which can then be used to identify VSEPR structures.
We started out finishing up the VSEPR theory lab. We used balloons to determine the shape and angles of molecules. Electron domains are a region where electrons are most likely to be found. Our group used balloons to make SF6 and BrF5. There are two types of of VSEPR structures. One is Electron Domain geometry which is based off how many electron domains there are per central atom. We learned that the 5 possible shapes that can be formed by electron domains: Linear, Trigonal Planar, Tetrahedral, Trigonal Pyramidal, and Octahedral. Linear structures contain two electron domains, trigonal planar have three electron domains, tetrahedral have four electron domains, trigonal bipyramidal have five electron domains and octahedral have six electron domains. The second type of structure is molecular domain geometry. Molecular domain geometry structures are based on how many actual bonded atoms there are connected to the central atom. If all of the electron domains of a central atom are bonded to another atom, then the molecular domain geometries are exactly the same as the electron domain geometries. 

We used balloons to create models of electron domain geometries and used gummies to create molecular geometries. These structures made out of the different items made it a lot easier to visualize the shape of the molecular structure, but I feel like the balloons were a little easier to use because the gummies kept melting out of place. 

We also learned a lot about how to make Lewis structures by using formal charges. Formal charges are assigned to atoms in molecules according to a set of rules. 1. Nonbonding electrons are assigned to the attached atom. 2. Shared electrons are evenly divided between the bonded atoms. A Lewis structure is not complete unless the formal charges are indicated (except "0" is not used). Formal charges by doing formal charge = core charge - number of assigned electrons. The best Lewis structures are the ones with the lowest formal charges. Formal charges greater than positive or negative 1 are never found in good Lewis structures. We also learned that atoms in the 3rd, 4th, and 5th periods can have extended octets, meaning that the sum of the bonding and lone-pair electrons can be greater than 8. 

This week I'd give my understanding on this subject about a 7. I really understand the idea of formal charges, and how to find them. It was easy for me to explain to my classmates on how to find it and how to use it. I also understand the extended octets on the 3rd, 4th, and 5th periods, and helped to remind my tablemates that when we were whiteboarding some of the atoms could have more than eight atoms. The reason my understanding of the topics is at a higher rating this week is because of the molecular and electron domain geometries. I'm still unsure about how to identify the different shapes of each molecule. My questions for this week would be how to identify the angles in the electron and molecular domain geometries? I hope this week I can understand electron domain and molecular geometries.

10-7-13

This week we started learning more and more about Lewis structures and molecules. We started with bond order and bond strength. Bond energy is defined as the energy required to sever the bond that holds two adjacent atoms together in a molecule. It's usually expressed on a molar basis. Bond order tells the type of bond between two atoms. For example, two atoms with a single bond has a bond order of 1. Two atoms with a triple bond have a bond order of 3. The higher the atomic number, the higher the bond order. Triple bonds have the highest amount of strength, while single bonds have the lowest amount of strength. Bond length is defined as the distance between the nuclei of two bonding atoms. The shorter the bond, the stronger the bond because there is more energy the shorter the bond.

We also did a lab to try and determine the percentage by mass (mass percent) of copper in brass screws or any other brass item. It was important during the lab to use extreme safety because of the concentrated nitric acid. Concentrated nitric acid is corrosive and will attack and destroy metals, plastics, and proteins. Skin contact with the acid would discolor the skin for days. The gas that forms from it is a toxic, reddish-brown gas of NO2. The nitric acid addition was performed under the fume hood and we all made sure we wore goggles. The nitric acid was added to the brass screw and left in the beaker over night. The next day, after the screw had dissolved, our group used the visual comparison test to determine the concentration of the unknown solution. The two test tubes of the reaction solution, and the unknown solution were placed on top of a white sheet of paper and then the intensity of the color of each solution was compared and the solution was removed or added until the colors finally matched. Then, after measuring the depth of each solution we used the equation (Molarity1)(Depth1) = (Molarity2)(Depth2). 

On Friday, we started the VSEPR Theory Lab and used balloons to determine the shape and angles of molecules. Electron domains are a region where electrons are most likely to be found. Our group used balloons to make SF6 and BrF5. This week, I'd give my understanding on this subject about an 8. I really understand the relationship between bond order, bond length, and bond energy. I really enjoyed the lab with watching the brass dissolve, and the colored gas come from the reaction between the two elements. The VSPER lab (that's still not done) was also really fun with the balloons, although I probably got a little carried away with the balloons. I still have questions on how to name the shapes that came from the gum drop models? I don't quite understand the octahedral and square pyramid part, so I'm hoping this week I'll get a better understanding of it. 

9-30-13

This week we finished up our stoichiometry unit and had a test on Wednesday. Stoichiometry is the mass and amount relationships between reactants and products in a chemical reaction. After the test, we started learning about Lewis structures. Lewis structures are also known as electron dot diagrams. They are named after Gilbert N. Lewis who published his essay "The Atom and the Molecule" in 1916. Lewis structures are used to provide a simple way for chemists to represent molecules that allow reasonable predictions to be made about the structure and properties of the actual molecule. The properties of a molecule really depend on how the electrons are distributed throughout the molecule. http://www.chem.ucla.edu/harding/lewisdots.html (this website gives a description on drawing Lewis structures which really helped me as well).
Two methods the atom Ne could be drawn is

The diagram on the left shows the Lewis structure. The amount of dots surrounding the atom gives the core charge. For example, the atom Ne has a core charge of +10. 

Molecules can also be shown with a Lewis structure. Covalent bonds are shown by drawing lines. Covalent bonds are the sharing of two electrons in the valence shell of both atoms. 

 This diagram shows sulfuric acid and the covalent bonds connecting Sulfur and Oxygen and Oxygen and Hydrogen. Hydrogen must always share two electrons known as a bonding pair. Also, the sum of the shared or bonding electrons and the lone pair electrons for carbon, nitrogen, oxygen, and fluorine atoms must be eight (an octet). The other elements usually follow the octet rule as well.
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. I also think that doing this POGIL really helped my understanding, and doing the whiteboard activities and listening to other people's ideas really helped as well. I hope that I understand this unit better than the stoichiometry unit and hopefully do better on the test.

9-23-13

This week, we learned more about stoichiometry, limited and excess reactants, reaction particle diagrams, yields, and empirical formulas. Stoichiometry is the mass and amount relationships between reactants and products in a chemical reaction. You start with the grams given, use molar mass of the given molecule to convert it to moles of the given molecule. Then you use a mole ration from the balanced equation to get moles of the unknown molecule. Using that, you use molar mass of the unknown molecule to convert it to mass of the given molecule and get grams calculated at the end. During a chemical reaction, if there are fixed amounts of reactants to work with, one of the reactants may be used up first. This prevents the production of more products. To do this in class, we used the example of assembling a race car and having a fixed set of parts to make one race car, and there was always a part that there wasn't enough of to make an entire race car, so that was the "limiting reactant". We also learned that the reactant with the smaller number of moles isn't always the limiting reactant. It all really depends on the ratios of the reactants, and the one with the smallest ratio is the limiting reactant.
We used reaction particle diagrams to show before and afters of reactions. Theoretical yield is what you're supposed to get, and it rarely gets 100%. Actual yield is what you get (lower than theoretical) and it's the measured amount or produced experimentally. The percentage yield formula is
(actual yield/theoretical yield)X100%.
An empirical formula is a formula that represents the simplest ratio among the elements of a compound. http://www.chem.tamu.edu/class/majors/tutorialnotefiles/empirical.htm (really explains the empirical formula).
                                                                                   

Reaction particle diagrams 

 

I really understand the stoichiometry problems and finding the limited and excess reactants. I think that I need a little more practice with empirical formulas, but I understand the idea of them and what they mean. 

9-16-13

This week, we learned about stoichiometry and the relationship between concentration of a solution and transmittance of light. Stoichiometry is the mass and amount relationships between reactants and products in a chemical reaction. In simpler terms, it refers to the relative proportion of components. It's also important for equations to be balanced. The molecules on one side of an equation must be equal to the molecules on the other side, and the masses (in both amu and grams), and the amount of moles on one side must add up to the mass on the other side of the equation. Mole ratios are the ratio of the coefficient of one molecule or formula unit in a balanced chemical equation to the coefficient of another one in the equation. I learned that in an equation such as, 4NH3(g) + 5O2(g) → 4NO(g) + 6H2O(g)  the mole ratio of ammonia (NH3) to oxygen is 4:5. All stoichiometric equations must have a mole ratio, otherwise they can't be completed. You can also use equations to find the amount of moles needed in a reaction. (http://www.esf.edu/efb/schulz/Stoichiometry.htm)

In class, we also did an experiment that helped show the relationship between the concentration of a solution and the amount of transmitted light that will pass through. We used blue #1 dye to make each solution. Each solution contained a certain amount of mL of the dye, and a certain amount of mL of deionized water. We also tested Powerade and Gatorade. From this experiment, I learned that the higher the concentration of the solution, the lower the amount of transmitted light. From this I learned that concentration and transmitted light are inversely related, and concentration and absorption are directly related. We used a colorimeter to get these results.

I understand the ideas we learned in class this week, especially from the lectures and the worksheets. With a little more practice on setting up the equations like those on the stoichiometry worksheets, I think I'll fully understand it.