SG Chemistry 2A Week 6 Blog

This week was the start of a new unit in SG chemistry.  This unit is unit 6, in which we are going to be focusing on the internal structure of an atom. As I heard that this is what we were going to be focusing on this unit I remembered back to previous years of chemistry about protons, neutrons and electrons and what the makeup of an atom consists of. As I remembered, I enjoyed learning about this before so I was excited to see what kind of labs we are going to be doing this week.

This week consisted of many different labs in which we explored mainly electrical charge of atoms and what causes this. The tape lab was our first introduction to it. In this lab we stuck two pieces of tape together and charged them by pulling them apart and then stuck them to a wooden stick as shown below:

After they were charged we observed how they were attracted to each other, and realized it was because of the static electricity, but at that point we didn't know what that was. I wondered what causes this static electricity and why both of the top pieces of tape were repelling as well as the bottom pieces of tape.

This is what we went over in our post lab worksheet. We learned how before the tapes stuck together they had equal charges (equal number of electrons), and after they were pulled apart, the top tape was positively charged (with one less electron) in the bottom tape was negatively charged (with one more electron). This really confuses me, I don't understand why or how electrons can jump between the atoms of the tapes, and I'm going to need more practice with this concept. We also debated over at what point it was that the electrons “jumped”. Was it when they first made contact, or was it the actual act of separation that caused this in the electrons? This is the question that Dr. Finnan asked us, and is one I still have unanswered.

In order to understand this concept of electrical conductivity, in our next lab we went around the room testing different substances for evidence of an electrical current running through it. Some of the substances are shown below, as well as their Molecular formula:

As it was to be expected all metals showed positive for this, but it was in the other substances that we tested where confusion started to come in. I wondered why sucrose (C12H22O11) conducted electricity as a liquid, but not as a solid or when dissolved in water. And so I had a question to answer, what makes a substance conduct electricity? After thinking long and hard about this as a class, we determined that if in any compound a metal is involved, it must conduct electricity. However this concept is still hard for me to understand, and I feel I still need more experience with it. And I still don't understand how sucrose could conduct electricity as a liquid, but not as a solid. What changes between these two phases?

To finish up the week we did a lab involving a U-tube and CuCl2(aq).  The setup is shown below:

 

The rods in the solution are graphite, a good conductor of electricity. This was good because an electric current had to be run through each of them, one positive and one negative. After running it overnight, we came back to something I found was very interesting.

 

What you see above is the separation of CuCl2(aq) into copper, which was collected on the positive side, and chlorine which was attracted to the negative side, and even the color changed into green. I found this very helpful to witness because I was actually able to see the separation of a compound into two separate elements. I know this is going to help me visualize future experiments, and it really helped me start to understand what causes a chemical change, electrolysis.

SG Chemistry 2A Week 5 Blog

This week in Chemistry we mainly white boarded worksheets in preparation for our test on Friday on Unit 5. By the end of unit 5 I found myself enjoying the concepts more. Unit 5 consisted much more of math as we did many problems that focused on finding the molecular and empirical formulas. I enjoyed this more because usually I am good at math and working with moles in chemistry has proved to come very naturally to me.

Over the weekend we worked on a worksheet where we mostly had to convert grams to moles. Because I was gone on the Friday before, this worksheet had really confused me. I couldn’t remember what labels to use when setting up the equations or even how to set them up, so it made it difficult for me.  However I was relieved on Monday in class when we were given time to work out the problems with our table groups and clarify the things I wasn’t sure about. As we talked as a class and reviewed concepts of converting from moles to grams or vise versa, I learned that we must always have units on our numbers, and that the correct way to set up one of these equations is to multiply fractions rather than divide in order to put more detail in and create opportunity for extra labeling. Once I actually started doing this instead of using “naked numbers” I was fascinated to see that I was doing better with the problems and it certainly made it easier to study when I reviewed for the test.

Continuing our Unit 5 review we did a worksheet in class that explored more of finding the empirical and molecular formulas.  Because we had cleared many things up the day before, I was ready to take on this higher level of chemistry. I understood the difference between empirical and molecular formulas right away, whereas I noticed other students were having trouble grasping it. Empirical formulas are based on experimental evidence and are the lowest whole number ratio for the elements in a compound, where molecular formulas are formulas containing the actual number of atoms in a compound, as shown in the following particle diagram below.

Sometimes the empirical formulas are the molecular formulas, but not always, which is why we have to be careful. The only way to determine which formula it is is by finding the compound’s grams per mole and seeing if it is consistent with the molar mass of the empirical formula. This makes me wonder if we are going to be doing a lab soon where we will have to find these formulas and determine what the formula is. However I am confused by how we would find the molar mass for the compound we would be working with unless it is given to us. And even though I understand this concept fairly well, I still feel I need more practice determining which one to use because it is not always clear.

At the end of the week we had our Unit 5 test which I feel went very well for me, thanks to our collaboration in class over the week, and I feel I am doing a good job of understanding the concepts. I still need to work on memorizing more of the atomic numbers, because it would make the work much quicker, and more practice using significant figures. Overall I enjoyed seeing how math is so vital in chemistry and I’m curious to see how we will apply these skills to real life situations such as in a lab.

SG Chemistry 2A Week 4 Blog

This week in Chemistry 2A we began by learning what a moles are and the Empirical formula that prepared us for later in the week when we began our two day lab where we reacted zinc and chloride. So that we would better understand the lab we would be doing later in the week, we finished up calculations from the Relative Mass activity. This was where we had taken the masses of boxes of different types of screws and were working with these numbers, but now we connected this idea and the calculations we had came up with and applied it to chemistry and working with different elements. We discovered that each of the different nuts and bolts represented a different element, and in this case the washers represented carbon. Using the concept of relative mass, we found the ratios of the bolts and nuts to washers, and used these ratios to find out what elements the bolts and nuts represented. After multiplying each by 12.02 (atomic mass of carbon) we discovered that the bolts, with a mass of 32.9, were sulfur, and the nuts, with a mass of 16.33, were oxygen. Doing this activity where we first used objects that we were familiar with, instead of jumping right into atomic masses and elements, really helped me to visualize what we were trying to find and was a really great transition into the lab, where we began taking real life measurements.

Before jumping into the lab, we began the “Relative Mass and the Mole” worksheet where we practiced putting our newfound knowledge of the mole into context. We learned that because we can’t count every single atom in a measurement, chemists had to come up with a unit, this unit they use is called a mole. A mole is the amount of grams of an element equal to its atomic mass. An equal number of atoms is in every mole, this number is 6.02e23. I was very surprised to discover this. Dr.Finnan had us try and guess how many atoms were in a jar that he showed us, containing one mole of sulfur. I guessed way too low, only four billion. Some other students even guessed only one million! We were all very surprised.

Before we began our lab on Wednesday, Dr.Finnan made sure to demonstrate what we would be doing. Using zinc and HCL, he combined the two in a glass beaker. We noticed that it began bubbling right away and emitting a gas. I was curious as to what this gas was, and Dr.Finnan had us guessing. Someone suggested we do the flaming splint test. I predicted that the gas was hydrogen, because we know that the zinc and HCL form zinc chloride, so that means the hydrogen from the HCL must escape at some point. The flaming splint test confirmed my hypothesis.

When we did this on our own (after having took the mass of the empty beaker and the zinc) the zinc and HCL reacted the same way, fizzing and bubbling, and we were able to observe hydrogen being emitting at a closer look also, just like in the picture below.

Overnight Dr.Finnan put the beakers with the zinc and HCL on hot plates, and the next day the substance was solidified and now was called zinc chloride.

That day in the lab we had to work with bunsen burners in order to get the water out of the zinc chloride to get a more accurate measurement, and my group was a bit scared, as we had never used them before. When the beaker was being heated, it turned into a brown liquid, and gave of a strange smell.

After all the classes measurements were taken and put on the board we found that the empirical formula for zinc chloride is ZnCl2. This lab felt very empowering because we got to experiment as real chemists do in the field and work together and collaborate as a team.

SG Chemistry 2A Week 3 Blog

This week in chemistry was a wrap up of unit four and and introduction to unit five. With the unit four test on Wednesday, we whited boarded and reviewed Monday and Tuesday in order to prepare for the test. The last concept we learned for unit four was mass and atomic proportions of atoms. We finished learning about this on monday when whiteboarding our answers to Unit 4 worksheet 4. Before discussing this and talking about it I was very confused about mass ratios. I was mainly confused about what the number represented, but after whiteboarding and also receiving after school help I realised that a mass ratio represents a proportion of one element to another in a compound (mass A/mass B). These ratios are also a good way to compare compounds to each other in terms of mass and element proportions. I was very glad to have this figured out once the test came around because there was a question on it.

On Tuesday groups discussed and white boarded the entire Unit 4 review, which was a very productive and useful class period for me.  On the review I understood very well how to distinguish between mixtures, pure substances, elements and compounds. This was good because my group members were having trouble with this and we were able to help each other because of it. The one question on the review we didn’t understand was when it asked if the molecules in a previous question were diatomic how many volumes of product would there be. We had a misconception that there would just be the same amount of molecules but each molecule would be twice as big and therefore still only one volume of the product. However we discovered after working with the class that this is not the case. Instead there will be two volumes of the product and twice as many molecules. Working as a class helped us to straighten this confusion out before the test, which I was thankful for.

Once we finished the test we did an activity where we had to guess how many packing peanuts were in a bag as our introduction to Unit 5. Dr.Finnan demonstrated that we can either count every single peanut, or we can take the total weight and divide that by the average weight of one. This is similar to how we find a number of atoms in a given space. This demonstration helped me to visualise what we will be doing in our next unit and was a good way to introduce the concept.

The next following days we did activities and worksheets that began to introduce us to the concept of relative mass. Because we just started learning about this, it’s still confusing for me. Working with the nuts, washers and bolts helped me to visualise what we were doing and I understood how to find the individual masses but once we had to start converting from individual pieces to boxes and barrels I started to get lost. I’m also still trying to figure out the proper way to set up a mathematical equation that involves differents measurements. But I’m sure once we get more into the unit I will begin to really grasp the concept better.