Week 10 Reflection- Final Blog

        In this final unit of SG Chem 2A, our class learned about the concept of balancing equations. We learned that just as there are patterns in the way elements bond and the types of compounds they make, there are patterns in the way compounds are rearranged during chemical reactions. After hearing this, I decided that knowing these patterns must make it easier to predict or even determine the outcome of a chemical reaction, and I wanted to learn more about what exactly these patterns were. To begin our exploration of balancing equations, our class began with a reaction equations worksheet. Right away, I understood the idea that there must be the same number of atoms on one side of an equation as there is on the other, but I wasn't completely sure how to go about completing the problems. The class was then introduced to the help of diagrams, and I found that they really helped make the problems a lot easier to solve. After using the diagrams as an aid through the worksheet, I picked up on the process of the problems very quickly, and I was soon able to complete the equations easily without the diagrams. Attached are some examples of balanced equations:

When the class became used to balancing equations, we moved on to learn about different patterns of chemical reactions. The first reaction we learned about was a combination, or synthesis, reaction. When I first heard the words, I knew that the equation would deal with compounds combining in some sort of way. I soon learned that a combination reaction is a chemical equation with fewer product molecules than reactant molecules, meaning that the reactant molecules combined. We also learned about decomposition reactions, where the chemical equation has fewer reactant molecules than products molecules. Other chemical reactions we learned about were single replacement and double replacement. A single replacement reaction involves one element replacing another in a compound, and a double replacement reaction involves a trade of "partners" between aqueous ionic compounds. The last reaction we learned about was a combustion reaction, which is a combination with oxygen so that most or all of the products contain oxygen. All of these reactions made sense to me right away, and I think I would be able to identify them with ease. Overall, throughout this unit, I found that the concepts connect to other units we have learned this trimester in many ways. Chemical equations can include elements, mixtures, compounds, and pure substances, which were some of the main ideas we covered in the first unit. In addition, some of the chemical formulas we balance are ionic compounds, which means the signs of the elements still have to be taken into consideration in order for the overall charge of the compound to be zero.

        During the course of this trimester, I believe I have had various strengths, as well as various weaknesses. Overall, I feel like my greatest weakness was understanding the labs we completed. Since we often did labs at the beginning of each unit to introduce some of the main ideas, I found that I was always confused about what exactly we were doing. I always understood the procedure and the process, so I knew what we were doing in that respect, but I was never sure what the results were telling us. I think my strengths included the math portions of the concepts we learned, such as calculating molar mass and balancing equations, as well as understanding the main ideas of each unit. I also think my strengths included the blog reflections each week, and I think they were a great way to enhance my comprehension of each concept we learned. There were a few things I would have done differently in order to improve my grade in this class. At the beginning of the trimester, I would have spent more time adding my thoughts and connections to my blogs in order to earn the extra points. In addition, I would have studied harder for the Unit 5 Test, especially on the concept of finding empirical formulas because that was one of the questions on the test I messed up. Overall, I really liked the style of this class and the way it was organized. I believe the white boarding and group discussions were a huge help in enhancing my understanding of the various concepts, and I appreciated the time we got to spend with our table groups. Although I found group work to be a great help, one recommendation of improvement to this class would be to have more teacher-directed lessons on the concepts we learn. This would help the class make sure they are doing all of the problems right, because I know that often times myself I found that I wasn't sure if I was doing a certain problem in the right way or not. Some recommendations for the blogging assignment would be to maybe make them mandatory every other week, and every week in between could be extra credit. That way, if a student is too busy one week, they wouldn't have to worry about it. I really think that the blogs helped though, so they should definitely be a continued assignment for this course!

Week 8&9 Reflection

        To continue our exploration of Unit 6, our SG Chem 2 class spent over a week learning about the concept of naming compounds. We began at the beginning of last week through a packet that introduced us to ionic compounds. To start the packet, we were given the number of atoms of two elements, and we had to write the formula of the compound and draw a particle diagram. At first I was confused about why we were doing that packet; it was nothing different from what we had done before, and it seemed too easy. I felt like I was missing the point, and I soon found out that I was. A few questions later in the packet, it asked to look back at the patterns of the compounds we had formed. After examining the formulas for a while, I noticed that each compound was composed of one metal element and one nonmetal element. When we discussed the question in our small groups, I also came to realize that the elements had certain charges, and that the charge of each compound as a whole was always equal to zero. After the realization of this concept, the packet began to make a lot of sense to me, and I was able to predict the formulas of certain compounds and answer the remaining questions with ease. When the whole class was finished with the packet, we recollected as a class to go over what we had learned. As a group, we were able to conclude that ionic compounds are compounds composed of one metal element and one nonmetal element, and the charges of the elements are always balanced. Later on, we were introduced to the criss-cross method, which is a technique designed to help determine the chemical formula of an ionic compound if the charges of the elements are known. I found this method to be very helpful when naming ionic compounds, and a good way to make the work as simple as possible and eliminate errors. Below is an example of the criss-cross method and how it is used:

        After learning about naming ionic compounds, we continued on to a different type of compound; molecular compounds. With the help of another packet, the class was walked through the process of naming molecular compounds, and learned how the process differed from naming ionic compounds. At first, I couldn't see the difference between ionic and molecular compounds, and I was confused about why molecular compounds had a special way to be named. After going through a few of the questions in the packet, it came to my attention that all the molecular compounds I had seen were a combination of a nonmetal element and another nonmetal element. When I realized this, it made sense to me for a few minutes, but then I was confused again. I wondered how nonmetals could combine with other nonmetals because the charge of a nonmetal is negative, so the charge of the compound as a whole would not equal zero. Then I remembered that ionic and molecular compounds were two different things, so I came to the conclusion that this must be one of their differing factors. After establishing this separation between ionic and molecular compounds in my head, the packet went on to talk about the prefixes used in naming molecular compounds. The concept of prefixes made sense to me, because they basically showed what the chemical formula of a molecular compound was. I also soon came to realize that prefixes are used because of the fact that the charges in a molecular compound don't equal zero, so a form of clarity is necessary when naming the compound. In an ionic formula, prefixes are not needed due to the ability of adding the charges to determine the formula of a compound. The realization of these concepts brought me a greater understanding of both ionic and molecular compounds as a whole, and helped guide my way throughout the remainder of the unit. Attached are some examples of naming molecular compounds:

        Towards the end of this past week, our class took the Unit 6 Test. I thought it was definitely the most challenging test we've had so far; I felt well prepared going into it, but during the test I thought differently. The multiple choice section was fairly tricky; there were some questions I was stuck between two answers and not sure which one to put down. The naming section was difficult as well, and I feel like I probably made a few errors here and there throughout. In some questions, part of the compound contained a polyatomic ion, but at times I wasn't sure if I was supposed to name it with the name of the ion or in just a basic way. After the naming section, I was stuck on the problems with the tape for a while because I was confused on the way they were set up. It was only after a while that I realized the tape was on the right hand side, not the paper and the metal, and that helped to clarify the diagrams. Now that I think about these problems, I think I forgot to mention that the electrons move in the metal because it is a conductor of electricity, whereas the electrons in the paper are attracted to the tape, but not a conductor of electricity, so they only shift to the side of the atom. During the test, I remembered that we had talked about problems just like those, but at the time I had a hard time remembering what exactly we had said. The last part I had trouble on in the test was towards the end with drawing a representation of a compound, and then showing how it splits into individual ions. I had no problem splitting the compound into its separate ions, but I did have trouble drawing the compound in the beginning. I was confused on how exactly the compound should be represented; since it contained a polyatomic ion, I wasn't sure if I would supposed to draw each element that made up the ion individually, or one atom that contained the ion as a whole. Overall, I definitely think the test was challenging, but I gave it my best effort, and the mistakes that I made will only be ones that I will be able to grow and improve from.

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.

Week 3 Reflection

        This week, my SG Chemistry 2 class explored a variety of concepts that helped us to further expand our knowledge on the subject. On Monday and Tuesday, we completed and discussed a worksheet on Avogadro's Hypothesis. To start, we had to look at a picture of two containers of gas particles held at the same temperature, volume, and pressure and come up with a conclusion about the number of gas particles each container held. Because the pressure, temperature, and volume were all constant, we were able to conclude that each container must hold the same number of gas particles. This was the same conclusion Avogadro came to long ago, making it possible to deduce the formulas of compounds formed when these gases react. For the next part of the worksheet, we had to draw representations of hydrogen and oxygen particles and react them together to form water molecules with no gas left over. Since two volumes of hydrogen react with one volume of oxygen to form water, we knew we had to draw twice as many hydrogen particles as oxygen. Our final product had two molecules of hydrogen attached to one volume of oxygen and we called it H20. We followed this process to react different molecules together in the next problems, and these were our results:

For the final part of the worksheet, we learned that occasionally, one volume of gas can react with another volume of gas to produce two volumes of a gaseous product. Avogadro came to the conclusion that when this occurs, the molecules of some gaseous elements must contain two atoms as opposed to only one. In the following problems, we had to draw representations of multiple volumes reactions of gases, and these are the results we came up with: 

        On Wednesday, we completed an online activity to further our understanding of reactants and their products. After measuring the mass and volume of various substances before and after they were heated or burned, we were able to come up with three main ideas. The first idea is that some substances are composed of discrete amounts of two or more other substances. This means that elements react in defined proportions to create a product. Our second main idea is that the total mass of the products of a chemical reaction is exactly equal to the mass of the reactants. This means that mass is never created or destroyed during the process. Our final idea is that elements combine in specific, defined ratios during chemical reactions. All three of these concepts helped us to further comprehend the process of chemical reactions and the behaviors of the reactants throughout the procedure.
        On Friday, our class started a worksheet on compounds and different hypotheses. My group found the ratios of the element's masses easily, but had a hard time sketching the particle diagrams. We brainstormed our ideas until we came to a consensus and then drew our thoughts on the white board. As a class, we then discussed the first part of the worksheet and my group edited our answers after listening to what everyone else had to say.

Week 1&2 Reflection

        To begin our exploration of SG Chemistry 2, my class started out by reviewing many of the principle concepts learned throughout SG Chemistry 1. First, we completed three unit study guides which highlighted the main ideas in each unit. In the Unit 1 study guide, we defined mass as the amount of "stuff" an object contains and volume as the amount of space an object takes up. Generally, mass is measured in grams (g) or kilograms (kg) and volume is measured in centimeters cubed (cm^3) or milliliters (mL). We also recalled that centimeters cubed and milliliters are directly related; 1 cm^3 is equal to 1 mL. This conversion makes it simple to transition from one unit of volume to the other when necessary. A final idea we reviewed in the Unit 1 study guide was that the mass of an object divided by its volume is equal to the density of that object. Density can be referred to as the "compactness" of a substance and each individual substance has its own, unique level of "compactness". This means that by looking at an object's density, you can figure out what substance that object is made up of. After completing and discussing the Unit 1 study guide as a class, we moved on to Unit 2. While reviewing the Unit 2 study guide, we recalled that particle movement is a result of, or results in, changes in pressure, volume, or temperature. Any factor that affects the number of collisions in a substance impacts the pressure, so pressure, volume, and temperature are all interrelated As the temperature of a substance increases, the faster the particles move due to the addition of heat, causing more collisions between the particles, and an increase in pressure. This means that temperature and pressure are directly related. As the volume of a substance decreases, there is less room for the particles to move, causing an increase in the number of collisions between the particles, and resulting in an increase of pressure. In contrast to temperature and pressure, volume and pressure are inversely related. The final unit we reviewed was Unit 3, which we discussed as a group after our Unit 2 study guide. We recalled the concept of energy, which is defined as a "substance-like" quantity that can be stored and transferred in physical systems. Energy stored by particles is called thermal energy, energy stored due to particle arrangement is phase energy, and energy stored due to particle attractions is called chemical energy. We used these ideas of energy to help us review our energy bar charts and our heating/cooling curves. The review of Unit 1, Unit 2, and Unit 3 of SG Chemistry 1 was a helpful was to jog our memories of the significant concepts that will help us to better understand SG Chemistry 2.
        To start off our first unit of SG Chemistry 2, we had to begin by defining some key vocabulary words. In small groups, we worked together to clarify the meanings of molecules, atoms, particles, compounds, mixtures, pure substances, and elements. Attached is an image of the definitions we came up with:

At the end of the week, we also talked about the properties of matter. Matter has many properties; texture, shape, size, etc. However, the most important properties of matter are the ones that we can measure. In our small groups, we brainstormed the properties of matter that can be measured, and added to our list as the class shared their ideas. Here is the final list we came up with:

        Though our first two weeks of SG Chemistry 2 were mainly a review of SG Chemistry 1, we discussed many significant and prominent concepts of chemistry that will guide us to success during the course of the trimester.