Week 7 Reflection

        Throughout this week, our SG Chem 2 class continued our exploration of Unit 6 through a variety of activities. We began on Monday by conducting a Sticky Tape Lab. In this lab, we took two pieces of tape and stuck them together on our table, one on top of the other so we had a top tape and a bottom tape. Then, we pulled them of the table and quickly pulled the two pieces apart from each other. We did this again with two other pieces of tape so we had two sets of top tapes and two sets of bottom tapes. At first, I was confused about why exactly we were doing this. I wasn't sure what pulling the tapes apart from each other would do, but when we put a top tape and a bottom tape next to each other, they were attracted and came together. Then, we put the top tape near another top tape, and we saw they moved away from each other; they seemed to repel. I tested this idea by putting my finger in between the two top tapes, and I could feel the repulsion taking place. The same thing happened with the bottom tape to the other bottom tape. After doing this, I concluded that pulling the tapes apart must have been giving them different charges. This made sense to me, because I know that opposite charges, like the top tape and the bottom tape, attract while the same charges, like the two top tapes or the two bottom tapes, repel. This was the conclusion the class was able to draw from the lab, and it helped to further our understanding of atoms and their charges.
        During the rest of the week, the focus of our class was on conductivity. On Thursday, we completed a Conductivity Lab to introduce us to the idea of what is conductive and what is not. Before doing this lab, I thought about what conductivity meant to me. It seemed that metals were usually conductive; I wasn't sure if all metals were or if only the majority were. I knew that all non metals were not conductive, but maybe if they were compounded in any state with a metal that was, they would be conductive as well. I kept these hypotheses of mine in my head throughout the experiment. For the lab, we went through fifteen stations, each with one or more substances. We used a cool tool to test each solution; if we put the tool on a substance and the light turned on, that meant that the substance was conductive, and if the light didn't turn on, that meant that the substance was not conductive. We tested a variety of different substances, as well as a variety of different substances in different states. We tested substances in the solid state, liquid state, molten state, and aqueous state, which meant the substance was dissolved in water. Below is a list of all the substances we tested:

My group made many observations throughout this lab, and we were surprised about some of the results. For example, I was surprised to find that NaCl in a liquid state was conductive, and I was also surprised that carbon was conductive in one allotrope, graphite, but not the other, diamond. On Friday, we went over this lab as a class, and each group white-boarded their results and thoughts. Attached is a picture of my group's board:

When we came together as a class after sharing our whiteboards, we were able to come to some conclusions. From our data, we were able to conclude that all metals are conductive, and that liquid, molten, or aqueous compounds that contain a metal are conductive as well. This means that all non metals are non conductive, as well as solid compounds with or without metal. These conclusions fit the data my group collected, and made a lot of sense to me. My hypotheses in the beginning were a little off, so this lab helped to improve my understanding of conductivity. All in all, our class spent the majority of this week exploring new concepts of chemistry that will help us improve our understanding of the subject overall.

Week 6 Reflection

        At the beginning of this week, we discussed empirical and molecular formulas to wrap up Unit 5. For a couple days, we worked on a worksheet to help us understand how to find these types of formulas. At first, I was confused what exactly the words empirical and molecular meant in terms of formulas, and I wasn't sure what the word empirical itself meant either. I later learned that empirical means something that is based on data. I also learned that an empirical formula is the simplest, whole number ratio of a compound, and a molecular formula is the exact formula of a compound, which would be a multiple of that compound's empirical formula. For example, the empirical formula of the zinc chloride from our previous lab was ZnCl2. If we had used twice the amount of zinc, the formula would be Zn2Cl4, which is the molecular formula of the compound. Thinking about these formulas in this way helped further my understanding of how they work, and made it easier for me to complete the worksheet we did in class. Below is a picture of some of the problems and worked out solutions on the worksheet:

Another concept from the worksheet was percent compositions of compounds. Finding percent compositions would tell us the make up of certain compounds, and how much of each element the compound is made of. For the problems on the worksheet, we were generally given the amount in grams of two or more elements in a compound, and sometimes the total mass of the compound as well. To find the percent composition of the elements, we had to divide the mass of the element by the total mass of the compound, and then multiply the answer by 100 to put the number in terms of a percentage. This part of the worksheet was very easy for me; I have been familiar with percentages in math for years now, so I understood the concept right away and knew exactly how to complete the problems. Attached is a picture of some of the percent composition problems from the worksheet:

        The day after completing our empirical/molecular formulas and percent composition worksheet, we worked on our review guide in class, and the next day we had the Unit 5 test. I think the test went well for me; I understood the concepts of moles and molar mass, as well as finding the relative mass of different substances. I also knew how to convert grams of a substance to moles, and moles of a substance to grams. I remember the test had problems where you had to find the number of particles of a substance, and I used Avogadro's number, 6.02 x 10^23, to find the answers. There were a few percent composition problems, and I had no trouble on those either. I took my time on all the problems so I could avoid making small errors, and I made sure to show all my work and used units wherever I could. My only issue on the test was with one of the empirical and molecular formula problems. The compound was made up of three elements, so I first used my conversions to find the number of moles for each element, and I got 0.2, 0.2, and 0.3. If I simplified these numbers down further, my ratio would be 1:1:1.5, and since 1.5 is not a whole number, I concluded that the ratio must be 2:2:3, which helped me find my empirical formula. After doing this, I knew that I would have to compare the molar mass of the empirical formula to the molar mass of the actual substance (which was given in the problem), so I calculated the molar mass of the empirical formula and got approximately 74.0g. The molar mass of the actual substance was 90.0g, and I was very confused because when I compared these two masses, the ratio was 1:1.2. This didn't make any sense to me, because the molecuar formula would then have to be 1.2 times the empirical formula, which would be 2.4:2.4:3.6, which isn't right because the ratio would have to have whole numbers. I must have made an error somewhere throughout the course of this problem, and I hope that I will be able to find out what my error was. Other than that problem, I didn't have trouble on any of the others, so I think that the test went well for me overall.
        The day after the test, on Friday, we began our exploration of Unit 6 through the Black Box activity. Each student was given a box with a certain pattern inside, and we had to figure out what the pattern was without being able to see inside the box. I thought this task was very difficult; it was a challenge to rely on senses other than sight to be able to determine what was inside the box. The main idea of the activity was to get us thinking about the unknown and how we can determine the unknown without actually seeing it. This is significantly relative to atoms; scientists have been able to determine the structures of atoms though they are extremely tiny. This made me wonder how they were able to do this; how were scientists able to find out so much about something so small? This concept will be one of our main ideas throughout Unit 6, and I'm hoping that some of my questions will be answered throughout the course of the next couple weeks.

Week 4&5 Reflection

        Over the past two weeks, our SG Chemistry 2 class discussed a variety of different concepts and ideas that expanded our knowledge of the subject. During the first week, we finished Unit 4 and took a test on the material. On Tuesday that week, we completed a worksheet on the expression of the ratio of the mass of certain elements to the total mass in a sample. This worksheet was simple to me, and I understood the concept right away. Later on, however, we got to a problem that was difficult for me to comprehend at first. We were given three compounds (A, B, and C) and each compound had a specific mass of nitrogen that combined with 1.00 grams of oxygen. After finding the ratio of nitrogen to oxygen for all three, I noticed right away that compound A had twice the amount of nitrogen as compound B, meaning that compound B had one particle of nitrogen and one particle of oxygen, whereas compound A had two particles of nitrogen and one particle of oxygen. The part that confused me was compound C, because its ratio of nitrogen to oxygen was half of compound B, which was stated as having a chemical formula of NO. If compound A had twice as many nitrogen particles than compound B, then would compound C have half a nitrogen particle? How could a compound have half a particle? After a class discussion of the worksheet, I was able to realize that compound C would not have half a nitrogen particle after all, but two oxygen particles instead. Having the discussion as a class was very helpful in this situation, and I was able to fully understand the entirety of the worksheet.
        The new couple days were dedicated to review. We worked on several worksheets in class to help us prepare and study for the test. Attached is a picture of the review guide we white boarded with our groups and then discussed as a class:

        After our test on Thursday, we had a small introduction of Unit 5 for the rest of the hour. Our class had to examine a large bag of styrofoam peanuts and see if we could come up with a way to determine the number of peanuts in the bag without counting. My immediate thought was to do some sort of volume calculation; maybe measure the length, width, and height of the bag to calculate its volume, then measure the length, width, and height of the peanut to find its volume, and then divide the volume of the bag by the volume of the peanut. Other people in the class seemed to have other ideas, however, mainly dealing with mass. We ended up deciding to find the mass of the whole bag and the mass of a peanut, and then dividing the mass of the bag by the mass of the peanut. I suppose that was an easier method than measuring the bag and peanuts and finding their volumes. In conclusion, our calculations told us that there were approximately 1500 peanuts in the bag. I was surprised and doubtful; by the looks of it, there seemed to be no more than 500 inside, max. Maybe this is how scientists felt when they discovered the atom; they must have been uncertain about their discoveries, but in the end, they just have to trust their data.
        On Friday, we began our exploration of relative mass through a container and hardware lab. My first thought when I heard the term "relative mass" was ratios of one mass to another. This would give the mass of one object in relation to another. For the lab, we started with four containers: one with 25 washers, one with 25 hex nuts, one with 25 bolts, and one with nothing inside. We then found the mass of each of the four containers and were able to find the mass of the 25 pieces by themselves by subtracting the mass of the empty container from the original mass of each. After finding these values, we had to answer the questions on the bottom in order to finish filling out the table. The questions were confusing to me; it was hard for me to logically try and figure out what exactly the questions were asking. In addition, it was often tricky to know when to subtract the weight of the container or the box, and when barrels came into the problems as well, it all became even more confusing. Discussing and white boarding the questions as a group helped me comprehend the problems more than when I was on my own, but they still proved to be difficult in my mind. Attached is a picture of my group's white board from this lab:

        The following Monday, our class continued the concept of relative mass by completing a POGIL activity. For the first part, we compared the ratio of the mass of chicken eggs to quail eggs, and found it to be 16:1. This first section was simple to me, and I understood the concepts of all the questions. By the beginning of the second part, I realized that we were learning about moles. I had never heard of a mole before, and I was slightly confused about what exactly they are. Then, I was told to think of a mole as I would think of a dozen; a dozen is a unit, and each dozen contains 12 objects. These objects can be anything, from a dozen tennis balls to a dozen pennies. Even though a dozen always consists of 12 items, the mass of the dozen can vary quite distinctly. This same characteristic applies to moles; just as a dozen pennies has a lower mass than a dozen tennis balls, a mole of carbon has a lower mass than, say, a mole of oxygen. Making this comparison significantly improved my understanding of moles, and has aided my comprehension of the worksheets and activities we have done since then in class.
        Over the next couple days in class, we worked on our empirical formula lab. The goal of this experiment was to react zinc with hydrochloric acid and come up with a chemical formula for the product, zinc chloride. The first thing we did was find the mass of our beaker and then the mass of our beaker with the zinc inside. Then, we subtracted the mass of the beaker from the mass of the beaker and zinc to find the mass of the zinc by itself, which would help us with later calculations. After reacting the zinc and hydrochloric acid together, we had to wait overnight to continue our experiment. Below is an image of our beaker while the contents were reacting:

The next day, we found that the substance left over in the beaker was solid and white, and covered the bottom of the glass. We then had to heat up the substance, zinc chloride, before we calculated the new mass of the beaker. At first, I wasn't sure why heating up the zinc chloride was necessary, but I soon found out that we needed to evaporate the left over hydrogen. Attached is an image of the heating process:

After heating the beaker twice and obtaining the same mass result each time, we were able to continue our calculations. To find the mass of the zinc chloride on its own, we subtracted the mass of the beaker we found earlier on. Then, to find the mass of the chlorine, we subtracted the mass of the zinc from the mass of the zinc chloride. Afterwards, we had to find the number of moles of both chlorine and zinc. I recalled that in order to convert from grams to moles, we had to multiple the grams by one mole over the molar mass of the substance. Using the periodic table, I was able to find the molar mass of both zinc and chlorine, and then used my conversion factors to calculate the number of moles each contained. My ending result was 0.055 moles of zinc and 0.11 moles of chlorine. The goal of the lab was to determine a chemical formula for zinc chloride, so my thinking was to compare the ratio of moles of zinc to moles of chlorine. Since 0.0055 moles : 0.11 moles is approximately 1:2, I was finally able to conclude that the chemical formula for zinc chloride is ZnCl2. Attached is our class data for the lab:

        Overall, the past two weeks of SG Chemistry 2 have been filled with significant concepts of the subject that will help guide us to a full understanding of the class as a whole.