Week 7 Blog

Ariana Trossen
Dr. Finnan
SG Chem 2
10/29/15

     This week in class we continued learning about Unit 6, which is all about the internal structures of particles. The structures of particles is how they get their names and this is what we studied this week in class. POGIL after POGIL we have learned all we need to, to be prepared for our test on Monday. There are different methods for naming different types of compounds. Ionic compounds have a way and molecular compounds have another. Writing the formulas for ionic compounds are a little more complicated to figure out than molecular compounds, so we spent most of the week with those and spent only a little time on molecular compounds. Then we ended the week by reviewing for our test on Monday.

     Ionic compounds contain a metal and are held together by ionic bonds. They have specific names for every ionic compound, but there is an easy pattern to figure it out from the formulas of the ionic compounds. If you are given NaCl, for example, you can see that there is a sodium atom and a chlorine atom. The metal is always listed first and the name stays the same, so sodium is still written out as "Sodium." Then instead of writing chlorine as "chlorine" it is written as "chloride." So the whole name of NaCl is "Sodium chloride." Another example is CaS, this is calcium and sulfur, which is written out as "Calcium sulfide." The pattern that you can see here is the the second atom in the compound, the nonmetal, its ending is changes to -ide. This part of naming ionic compounds is very simple. However, there is more; there are polyatomic ions.

     Polyatomic ions are a group of atoms that carry a charge. Some examples of polyatomic ions are NO3, CO3, PO4, and C2H3O2. The names of these are nitrate, carbonate, phosphate, and acetate. We were not required to memorize all of the polyatomic ions and were provided with a sheet that had them listed for the test. If you had an ionic compound with a polyatomic ion in it, it is pretty simple to name. For example, BaSO4 this is barium and sulfate compounded together, and the name is simply written as "Barium sulfate." Naming ionic compounds is not difficult, but writing their formulas from  their names is a little trickier.

This picture is of three problems that our group white boarded out. It shows two different atoms and we were directed to write out the formula and name of the ionic compound formed by the two given atoms in each problem. We also drew out underneath the formula units and what the compound would look like.

     It is important to remember that the net charge of the compound has to equal zero. We were not required to memorize the charges of each atom and polyatomic ion and were again provided with a sheet that listed them for us to use on the test. An example of writing the formula of an ionic compound from it's name is this, you are given  Potassium permanganate, the symbol for potassium is K and the symbol for permaganate is MnO4. The formula for this one is simple  because the charge of potassium is 1+ and the charge for permanganate is 1- so they cancel out and equal a net charge of zero, the formula is KMnO4. Another example is Copper (II) nitrate, the symbol for copper is  Cu and the symbol for nitrate is NO3. This one is a little more complicated because copper has a varible charge and the roman numerals in parentheses indicates what charge it has, in this case copper has a charge of 2+ and nitrate always has a charge of 1-. This means that the formula must be written this way: Cu(NO3)2, to make the net charge of the ionic compound equal to zero. The same thing is done with writing the formulas of molecular compounds.

     Molecular compounds are made up of all non-metals held together by covalent bonds. When writing out the names and formulas of molecular compounds it is fairly straight forward. When naming molecular compounds prefixes are added depending on the number of the atom in the compound. An example is IF7, which is written out as "Iodine heptaflouride." Just like in ionic compounds, the second atom in the compound ends in -ide. Another example is P4O10, and is written out as "Tetraphosphorus decoxide." An impotant thing to notice and remember is that when a prefix is added to oxygen the vowel on the end of the prefix is dropped and then added to the front of oxygen. All of the prefixes are mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6), hepta- (7), octa- (8), nona- (9), and deca- (10). This is important to know when writing the formula of a molecular compound when given its name. An example of this is Tribromine octoxide, the formula for this would be Br3O8. Another example is Chlorine monoflouride, the formula for this would be ClF. After doing these, we ended the week by reviewing for the Unit 6 test.

This picture is of two problems from one of the POGILs that our group white boarded out. In problem 20, we analyzed the compound (NH4)3PO4 by looking at the number of each element in the compound. In problem 21, we broke down the compound in the first part, then in the second part we talked about what was wrong with the student's formula that was given.

     While reviewing we looked back on the experiments of J.J. Thomson and out Sticky Tape lab. We also continued practicing naming and writing the formulas for ionic and molecular compounds. Among some other things we looked at these all in preparation for our Unit 6 test on Monday, and I feel confident that I did well.

In this picture, Dr. Finnan is holding up a flame to a sheet of copper to show the change in the metal as the heat is applied. The color of it changes and the flame appears to be green, but when the flame is pulled away and heat is no longer being applied the color of the copper sheet returns to how it was originally.

Week 6 Blog

Ariana Trossen
Dr. Finnan
SG Chem 2
10/25/15

     We finished last week with our Unit 5 test, so this week we began Unit 6, which is all about the internal structures of particles. We looked at J.J. Thomson and his experiments with cathode rays. Then, we did our own experiment with sticky tape to understand the interactions it had with other materials. After that we tested the conductivity of different elements and compounds. Then, we went deeper by studying what was happening to the sticky tape as it interacted with the other materials and how its charge was affected. We also observed the electrolysis of CuCl2. To end the week we started looking at the composition of compounds that are solids that don't conduct electricity and found patterns in their equations.

     J.J. Thomson did experiments with cathode rays in the late 1800s. Cathode rays are beams of electrodes emitted from the cathode of a high-vacuum tube. We looked at three different experiments that he performed. In the first one, Thomson directed a beam at an electrometer and tried to separate the evidence of charge from the path of the beam. He found that cathode rays deposited an electric charge, and he found that he could bend the rays with a magnet. From this first experiment, he ultimately discovered that the charge was negative and cannot be separated from the rays. In the second experiment, Thomson tried passing the cathode ray through an electric field. When the ray beam passed through the electric field it bent. From this he concluded that cathode rays are charges of negative electricity carried by particles of matter. In the third experiment, Thomson did some careful measurements on how much the path of the cathode ray was bent in a magnetic field and how much energy was carried. Thomson was able to describe the mass/change ratio of the cathode particles. Thomson also came up with an amazing result from this experiment, he found that the mass-change ratio of cathode rays is far smaller than that of a charged hydrogen. This means that cathode rats either carry an enormous charge or are extremely light relative to their charge. From these experiments Thomson presented three hypotheses, only two of which were accepted as true. From those two a model of the atom was made, know as the Plumb Pudding model. This model will help us to understand the internal structures of particles.

     In the sticky tape lab we took a piece of tape and folded it, making a bottom half and then doing the same to make a top half. Then, we stuck the pieces of tape together and pulled them apart quickly to give them a charge. After that, we observed how the different pieces of tape interacted with each other and with paper and aluminum foil. We found that the top and bottom pieces were attracted to each other and the top and bottom tape both attracted the foil and the paper, but the top tape repelled itself and the bottom tape repelled itself. From this we determined that the top tape was positively charged, the bottom tape was negatively charged, and the paper and foil were both neutral. Ultimately from this lab we learned that positive and negative things are attracted to each other, two positive things repel each other, to negative things repel each other, and both positive and negative things are attracted to neutral things.

     This week we used Dr. Finnan's homemade conductivity testers to test the conductivity of different elements and compounds. We found that all metals and compounds with metals in them conduct electricity. Then, we discussed as groups about what we thought the reason was for the different things that do conduct electricity. Our group concluded that metals conduct because they have free electrons that flow easily and ionic compounds (a metal and a non-metal bonded together) conduct because they have free ions that move easily.

     To end the week we started looking at different compounds and their particle drawings and formulas. We were supposed to then look at patterns between the different compounds, but that's where we ended and I couldn't and still don't really understand what kinds of patterns I'm looking for. We also watched the electrolysis of CuCl2 and wrote down our observations as bubbles formed on the positive side, it smelled like a pool, and the bottom of the copper rod on the negative side began turning pink. After a little longer it continued to smell like a pool, the pink spread, and there were bubbles on both sides. Overall, this unit is very interesting and so far I like what we're learning.

Week 5 Blog

Ariana Trossen
Dr. Finnan
SG Chem 2
10/17/15

     This week in chemistry we did a lot of practice to prepare us for our test which we had at the end of the week on Friday. Our test covered Unit 5 which is all about counting particles. Last week we mastered the basics and were introduced to some more complicated concepts of counting particles. This week we mastered the more complicated concepts and showed our understanding on our tests.
Some of the more complicated concepts include calculating how many moles, molecules, or atoms are in an element or compound, empirical formulas, and molecular formulas. At the beginning of the week I felt a little unsure about these concepts, but now I feel a lot more confident about them.

     To begin the week we took a look at the packet from last week that asked about the number of atoms, molecules, or moles in an element or compound. Each group was assigned a different problem from the packet to whiteboard. Our group was assigned problem #7, which asked us to find the number of atoms in Bottle 7. Bottle 7 contained some iron nails. In order to solve the problem, we have to start by finding the mass of the bottle with the nails inside, then we have to subtract the Tare Weight (T.W.) from that mass. The T.W. is the mass of the bottle when it is empty. By subtracting the T.W. from the mass of the bottle with the nails we find the mass of the iron nails by themselves. After that you have to find what 1 mole of iron is by looking at the periodic table. Then you divide the mass of the iron nails by that number. This will equal how many moles of iron the iron nails are. With this new number you multiply it by Avogadro's number, which is 6.022×10²³, and you will find the number of atoms in the iron nails contained in Bottle 7.

     After more practice with problems solving for the number of atoms, molecules, and moles in elements and compounds, we began to look at empirical and molecular formulas more. The difference between empirical and molecular formulas is empirical formulas are based on data and molecular formulas are what the compound's formula truly is. Molecular formulas are generally more complicated, and empirical formulas are more simple. To find the empirical formula of a compound you have to start by finding the number of moles of each element in the compound. Then, you find the ratio of the moles to find out how many atoms of each elements are in a molecule of that compound. This will tell you what the empirical formula is. To find the molecular formula you start the same way as you do to find the empirical formula, but then you go on by taking the given mass of the compound and dividing it by the sum of the atomic masses of the elements in the compound. This will give you a number which you will multiply by the number of atoms you found for each element in the empirical formula. By multiplying, you find the more complicated molecular formula of the compound.

     Before the test we were given a review guide to cover everything we went over in Unit 5. From the definition of a mole to finding the percentages of elements that make up a compound, we reviewed for the test. Then on Friday we took it. I was a little nervous coming in, but as we started and I began to answer the questions with confidence, I wasn't nervous anymore. I feel that I have a very good understanding of the concepts in Unit 5, Counting Particles.

Week 4 Blog

Ariana Trossen

SG Chem 2

Dr. Finnan

10/10/15

     This week in class we continued to discover how to count particles. First, we ratio of masses of chicken eggs and quail eggs. Then, we looked at Avogadro's number, 6.022×10²³, which is a mole. After that, we began to calculate the number of moles in something and how one mole compares to other large numbers of things, such as cells in the human body. To further our understanding of counting particles, we did an experiment of reactions with zinc chloride to determine the empirical formula. We learned about Tare Weight next and how it helps us determine the mass of contents in a bottle. For homework this weekend we got a worksheet to help us to continue to practicing calculating the moles, atoms, and molecules in different objects and compounds. Before looking at Avogadro's number, we looked at the masses of chicken and quail egg samples.

     A standard chicken egg has a mass of 37.44g and a standard quail egg has a mass of 2.34g. From this we were able to determine that the ratio of masses between chicken and quail eggs is 16:1. Then, we compared larger samples of chicken and quail eggs, such as 10 eggs, 438, 1 dozen, and 1 million, In each scenario we found that the ratio of their masses is always 16:1. Using this information, we compared the masses of elements instead and found the same thing; the ratio of masses of two different elements, no matter how many atoms are in each sample, as long as they're the same, will always have the same ratio. Knowing this information makes it possible for us to find the the number of atoms in two samples of elements and find the masses of samples of elements. This also helped us begin to understand what a mole is.

     A mole is 6.022×10²³ particles, this is also known as Avogadro' s number. A mole is so big that you can't even count that high. To really understand how large a mole is, we compared it to the number of cells in the body of every human on earth. We found that mole is still greater than that. It is important to compare and understand ratios between samples and elements because it helps us to determine the relationships between them. We experimented on zinc to find out how it reacted with chlorine and their relationship.

     This lab we did was called the Empirical Formula Lab. In this experiment we put zinc and hydrochloric acid together. From this we were supposed to find the empirical formula of zinc chloride, a product of the reaction between zinc and hydrochloric acid. Empirical means "based on experimental evidence." Based on our experimental evidence we found that the empirical formula for zinc chloride is ZnCl2. We determined this by calculating masses and determining the number of moles of zinc and chlorine. By finding the ratio of the moles of chlorine and zinc and it equaled approximately 2, so the ratio of moles of chlorine to zinc is 2:1. After looking at how zinc and chlorine compare in zinc chloride, we moved on to look at sample containers of different elements and compounds to calculate the number of atoms, molecules, and moles.

     In class we were given multiple different bottles that contained different elements and compounds. Our worksheet asked us questions about how many atoms, or how many molecules, or how many moles were in each bottle. However, in order to find these things we needed to subtract the mass of the bottle from the mass of the bottle plus the element or compound inside. The mass of the bottle itself is known as the Tare Weight (T.W.).  Once we found the individual element or compound mass we could continue to calculate the moles, atoms, or molecules. To do this you have to know that the atomic number of each elements is equivalent to one mole of that element. If you are given a sample of an element, you can determine the portion of a mole that the sample equals. Each group in our class was assigned a specific problem to white board out, we were assigned problem #7. In this problem we were asked to find the number of atoms in a bottle of iron nails. In order to do this, we found the mass of the bottle plus the iron nails. Then, we subtracted the T.W. from the total mass, which gave of the mass of the iron nails. After that, we divided the mass of the iron nails by the atomic mass of iron. We then took that number and multiplied it by Avogadro's number, 6.022×10²³, and this gave us the answer of how many atoms were in the sample of iron nails.

     To end the week we were given a worksheet for extra practice on calculating the moles, atoms, and molecules. I am glad we got this worksheet. I feel the more practice I get with these problems, the better I am at solving them. However, I still don't feel 100% sure about all of them yet. I still struggle sometimes with knowing what numbers to divide or multiply by what. Ultimately, I feel I am getting better.

Week 3 Blog

Ariana Trossen
Dr. Finnan
SG Chem 2
3/10/15

     This week we took our Unit 4 test and started Unit 5. On Monday and Tuesday we spent time reviewing and preparing for the test. We started class on Monday by white boarding out the answers to the last two problems from Unit 4 Worksheet 4. Our group was assigned problem 6. By comparing the ratio of masses in two compounds, we can determine the formula for the  second compound based on the first. In this case we knew that compound 1 was water, with the formula H2O, and had a mass ratio of 1/8 (hydrogen/oxygen). The mass ratio of compound 2 was 1/16, from this we found that its formula would be H2O2, hydrogen peroxide. This is because we know the number of hydrogen in each formula stays the same based on the formula for water having two hydrogens and the numerators in each mas ratio, which represents hydrogen in each compound, are the same. The number of oxygens in compound 2 is twice the amount in compound 1 because the denominator, which represents oxygen in each compound, in compound 2 is twice that in compound one. We were able to determine the formula for one compound based on the mass ratios for each and knowing what one of the compounds was.

     On Tuesday we worked on the Unit 4 Review packet. We went back to the beginning of this unit and worked our way to the end testing our knowledge, before the actual test. First, we studied the differences between elements, compounds, mixtures, and pure substances. Elements cannot be broken down by chemical or physical means. Mixtures are variable; their physical properties depend on the composition, and can be separated by physical means. Compounds can only be separated by chemical means, and consist of two or more elements in a fixed mass ratio. We then identified different compounds, mixtures, and elements as particle drawings. 

    We continued our review by sketching out how different gases react to form compounds. An important thing we learned was about diatomic atoms. Hydrogen, oxygen, fluorine, chlorine, bromine, iodine, and astatine are all diatomic. This means that they don't occur as single atoms, but two of the same element bonded together.

      We reviewed further with mass ratio comparisons of different compounds, similar to problem 6 on Unit 4 Worksheet 4. With the ratios we were able to sketch particle diagrams of the two different compounds, and then write out formulas for each.

     Different elements have different properties, such as melting and boiling points. We showed this by using time and temperature graphs lines showing the relationships between different elements and their boiling points. If you were to have a mixture of two different elements, you could heat the mixture to the temperature of the lowest boiling point of the two elements and separate them in this way.

     On Wednesday we took the Unit 4 test, which covered all of these concepts. Then we went  on to start Unit 5, Counting Particles.

     To begin to understand the concept of counting particles, we looked at relative mass of different kinds of hardware. We discovered that by weighing things you can count them, by knowing the mass of something you can determine how many of that something there is.

     To end the week we began to dive into comparing the relative masses of different elements in various oxides. By this we were able to see how Dalton's assumption of relative mass compares to the adjusted relative mass. Since we did not have enough time to finish this worksheet in class, I am still a little confused on how they compare. Altogether, this week I feel very good about my knowledge on Unit 4 after scoring a 33/34, and am excited to continue learning how to count particles in Unit 5.

Week 2 Blog

Ariana Trossen
5th hour SG Chem 2
Dr. Finnan

     This week in class we dove deeper into chemistry by looking at the history of it. We discovered how chemistry began by learning about the different scientists who had a hand in its foundation. From Empedocles to Democritus to Dalton, we learned about what people believed about the things around us are made of. We also studied Joseph Priestley's discovery of oxygen and Antione Lavoisier's theft by taking the credit. Something I didn't know before was that Napoleon Bonaparte loved science more than war and because of this it led to a competition between Humphry Davy and Joseph-Louis Gay-Lussac. After looking the history of chemistry, we studied the ratios of the masses of elements in a compound and their equations. Then we applied the ratio of the masses of elements in known compounds and made inferences on their compositions. Last week we learned all about atoms, elements, particles, and compounds, but people didn't always know about them.

     Empedocles was a Greek philosopher and scientist, he proposed the first theory about what the things around us are made of. He said that all matter was composed of four different elements: fire, air, water, and earth. However, there were many flaws in this because no matter how many times you break something it never ends up as one of those elements. That led to other scientists drawing different conclusions.

     Democritus was also Greek, though unlike Empedocles, his new theory for what things are made of was based on reasoning rather than science. He believed that if something, a stone for example, was cut in half again and again until it couldn't be anymore, then you would end up with these microscopic pieces, which he called "atomos", "Atomos" means invisible, the pieces that make up the stone are so small they can't be seen with the naked eye. His idea was dismissed at first, but later became the foundation of what we know today.

      Joseph Priestley began to experiment with red mercury calx. He discovered that when it is heated it breaks down into two different substances: liquid mercury and a strange gas. Priestley went further by collecting the gas in jars and studying it. He ran some tests with it, discovering that it made flames burn stronger and allowed a mouse to live longer with it than "normal air." Dephlogisticated air is what Priestley decided to call it, but it didn't stick.

   After hearing about Priestley's discovery, Antoine Lavoisier had to perform the experiments himself. He did and decided to rename the strange gas oxygen. Lavoisier went on to discover hydrogen and many other elements as he continued to experiment. He was very technical and weighed everything. This led to Lavoisier establishing the Law of Conservation of Mass: mass cannot be created or destroyed. More and more discoveries were being made on what matter is.

     John Dalton put the pieces together and formed a modern atomic theory with four main concepts. These concepts were: 1) All matter is composed of indivisible particles called atoms. 2) All atoms of a given element are identical; atoms of different elements have different properties. 3) Chemical reactions involve the combination of atoms, not the destruction of atoms. 4) When elements react to form compounds, they react in defined, whole-number ratios. Dalton gave structure to the ideas people and scientist were experimenting on.

    Humphry Davy and Joseph-Louis Gay-Lussac were rivals, and they were quite opposites. Davy was flamboyant, charismatic, loved leisure, and spent time on many different hobbies. Gay-Lussac was patient, careful, fully devoted to science, and spent time with his family when he wasn't working. Davy discovered sodium and potassium, and after he did so Gay-Lussac and himself raced to discover as much as they could about the new elements. When a strange substance that formed black crystals and could produce a purple vapor was found in seaweed, Gay-Lussac was asked to review and perform the scientists who found it's experiments. He decided to call it "iode," which means purple in Greek, due to its purple vapor. He had no idea that Davy had been told about it and started experimenting as well. Another competition began between them. Gay-Lussac beat Davy to the press by one day with the result that what is called iodine today is an element itself and not composed of chlorine as they suspected at first. There are many different elements, about 116 known today, and they all have their own unique properties.

     By looking at the ratio of the masses of a compound, we drew out different particle drawings and equations to represent the compound. One hypothesis was to draw out the atoms with them having the same mass. The second hypothesis was to draw out the atoms with one being heavier than the other based on the ratio. As a class we determined that the second hypothesis was the better option because it made more sense based on our learning from Dalotn and others that every element has different properties and therefore different masses. After this we applied the ratio of masses of compounds to compare unknown compounds with a known compound. These images below show how we calculated the ratios and percentages, and how we determined if the unknown compounds matched up with the known compound, in this case sucrose.

  

     As we learned new things each day in class this week, it all built on each other. Starting with Empedocles and ending with what we know today by studying compounds and their chemical equations. Each scientist provided a stepping stone for the next as they made their discoveries. By looking at what the previous scientists found and experimenting on new ideas and making inferences, we were able to apply what we've learned from these scientists, who gave chemistry its foundation, to learn about different elements and compounds and how they relate.

     One question I still have is: what is the Law of Definite Proportions? We didn't have quite enough time to cover it in class, and I'm interested to know how it relates to the comparing of known compounds to unknown compounds.

Week 1 Blog

Ariana Trossen
5th hour SG Chem 2
Dr. Finnan

This week in class we reviewed classification of matter and techniques for separating compounds and mixtures, and we learned some new information, such as, Avogadro's hypothesis. While reviewing classification of matter, we looked at the difference between atoms, particles, and molecules. Atoms are the individual building blocks of particles and molecules. A particle is any single atom or multiple atoms chemically bonded together. A molecule is two or more atoms chemically bonded together. We went further after this to compare pure substances, mixtures, elements, and compounds. A pure substance is any combination of a single type of atom or molecule. A mixture is a combination of different types of atoms and/or molecules. Elements are combinations of a single type of atom, they can be individual or chemically bonded together. A compound is a combination of different types of atoms chemically bonded together. All elements and compounds are pure substances. Compounds and mixtures can be separated, but not in the same ways. Mixtures can be separated physically, such as with a magnet or by filtering, but compounds have to be separated by chemical means because there are chemical bonds to hold the atoms together, ways to do this include electrolysis and decomposition. We went on to learn about Avogadro's hypothesis, which states that with two given samples of an ideal gas, of the same volume and at the same temperature and pressure, contain the same number of molecules. This makes it possible to find the formulas of compounds formed when gases react.226 In some cases we found that in order to keep the mass the same the molecules of some gaseous elements have to have two atoms. An example of this is when hydrogen and chlorine react to form hydrogen chloride.

An experiment we did in class was with sugar cubes to see how they would react when placed in water, ethanol, and a 50/50 mix of water and ethanol.

https://www.evernote.com/shard/s83/share/eee0-s83/res/d06e68b0-e06d-48a6-955a-9dd572f7a0e9/dissolvingSugarClass.mp4 (video link of experiment)

In this experiment we discovered that the sugar cube dissolves completely in water, not at all in ethanol, and not quite completely in the 50/50 mix and also at a slower rate than in pure water. After doing the experiment we white boarded out particle drawings of the before and after for the sugar cube placed in water and the sugar cube placed in ethanol.

All these things we learned this week are connected. We started from the beginning with the basics of how matter is classified. Atoms are the building blocks of molecules and both an atom and a molecule are considered particles. Then we have atoms and molecules that make up pure substances, mixtures, elements, or compounds. A mixture is different molecules and/or atoms mixed together. A pure substance is any single type of atom or molecule. An element is any atom or molecule made up of a single type of atom. A compound is any type of molecule made up of different atoms chemically bonded together. Elements and compounds are both pure substances. Then we expanded by talking about how compounds and mixtures are separated. Compounds have to be separated by chemical means, which we were showed in the experiment of separating water molecules. This also proved that water is made up of two hydrogens and one oxygen. Mixtures, however can be separated by physical means, which we were showed when a magnet was used to separate iron from a mixture with sulfur. Avogadro's hypothesis ties into this as we learned to draw out the formulas for how elements react and form compounds with particle drawings. I feel we learned a lot this week and am excited to further my knowledge and understanding of chemistry.