Chemistry Entry 8

Our week started with the stressful news that the majority of our class performed poorly on the chapter 6 test, primarily on the "naming ionic compounds" portion. As a a result, we were assigned a series of quizzes focused on each section of chapter six, with the goal to enhance our skills in the subject. We began these quizzes last Tuesday in class, and I have finished all of them except for the last one. The first two quizzes didn't give me that much of a challenge, as they were focused on ionic compounds with both fixed and variable cations. At first I was answering all the questions incorrectly and I kept getting frustrated, but then I realized that the answers I had to type in are case-sensitive, and I was needlessly capitalizing them. The third quiz was slightly harder than the first two, as it focused on polyatomic ions. I always get intimidated when I see a large string of compounds and I overthink what I have to do, which results in mistakes. As long as I stay focused and know that its the same as any other compound, just composed of a larger amount of elements, then I'll be fine.

Phosphate has a charge of 1, not 4! There are 4 oxygens and
1 phosphorus present in phosphate. Having sodium there
threw me off at first.

Quiz 5 was a breeze for me as I'm confident with covalent bond questions, but Quiz 4 was my biggest challenge. Quiz 4 focused on counting atoms, ions, and particles. For the first mini-quiz, I had to redo it twice as I kept misreading the question. I continuously mixed up on counting ions and the # of atoms. I tried to include them together in terms of having the same number for the answer, but  eventually a lightbulb went off and I realized my dingbat mistake. My main issue was forgetting that polyatomic compounds as a whole, have the same charge. There aren't two different charges present in the compound as if that were the case, then the equation wouldn't be balanced. I kept wanting to count the subscript for both categories, which only applies to the # of atoms of the element it's representing. Once I realized this, my percentage didn't drop below a 90% and I felt much more confident on the matter.

Ammonium!! 

On a side note, in general I need to review more on equations involving ammonium, the unique polyatomic ion with a positive charge. Whenever there is a question involving ammonium I always seem to incorrectly answer it, and only correctly answer it if I have previously memorized the answer. Something about it always throws me off, so I'm planning on consistently redoing quiz 4 so as to strengthen my confidence One more thing I'm planning on studying for our upcoming retake, which is also present in quiz 4, are formula units. These mess me up simply because I occasionally incorrectly  balance the charges for the equation, so the then the diagram is incorrect as well. I also need to work on identifying an equation through a given formula unit, as this type of question consistently messed me up on quiz 4. I hope we can go over some of these problems in class this week, as class discussions improve my understanding 99% of the time.

Our first lab we completed in Unit 7 was called the Nail Lab. The purpose of of this was to determine the rate of copper produced to iron consumed in a replacement reaction. Our first step mass our beaker so as our measurements weren't off later on, and then we combined copper (II) chloride with distilled water in our beaker, and began to swirl the contents. I noticed the contents turned an electric blue upon stirring it, but wondered if the color would change as we continued the lab. During this process,

Setting up for the nail lab.

Day 2 of the nail lab.         

my other table members massed out 2 nails together, and then we placed them into the solution. Somewhat true to my prediction, the nail itself reacted by turning a reddish brownish color and it began to crumble. I was surprised to see, in the next day of our observations, that the nails seems to be combined together (as shown below) and the color was a dark green. I also noticed that the reaction between the nails and the copper (II) chloride ceased. My thought process on why this occurred is that all of the iron on the nail must have been used up, so the reaction couldn't continue. This makes sense to me because, when we massed out the beaker and compared it to the previous days, the data showed that all of the iron (Fe) was used up overnight. I love doing these types of labs, especially because I find it so interesting how seemingly normal substances can combine to create an astonishing outcome. To conclude this lab, all table groups compiled their data and recorded them in a table drawn up on the whiteboard, so as to compare results and see how other groups differed from each other.

Combination (synthesis) reaction.

Decomposition reaction.

To finish up the week, our class learned how to balance equations and learned how to describe difference and similarities between various chemical reactions. Funnily enough, the information we covered in the end of the week is the stuff I have to study for the most. In terms of equation balancing, I have a bad habit of overthinking the problem and then coming up with a random answer that doesn't make sense. I figured out that I need to balance off the reaction with the monatomics last, as well as learn to leave diatomics towards the end of the balancing. Remembering to use fractional coefficients first and then multiplying everything so as to make the fraction a whole number is also a section that needs work, which I need to keep in mind as I balance out any future equation. Resuming on Monday, I hope we discuss more as a class on describing chemical reactions and their specific names, as I am completely unfamiliar with this and how to correctly accomplish this. Specifically differences between reactions such decomposition reactions or double replacement ones, as well as the difference between single and double replacement reactions.

I found it really cool that I could tell the reaction was taking
place by the increase in temperature in the beaker!!

Chemistry Entry 7

Metals that form multiple ions.

On Monday our class began a worksheet based on naming ionic compounds. An ionic compound is a chemical compound in which ions are held together in a structure, but contains a neutral charge. The worksheet was specifically focused on expanding our knowledge of the structural units that make up ionic compounds, and using that information to name them. The key to understanding this, for me at least, was keeping in mind that the chemical formulas for any compound must have a net charge of zero, resulting in varying amounts of atoms that need to be added to an equation, so as the formula equals zero. I found this rule confusing at first as I kept failing to make the connection that each element has its own unique charge, and that occasionally ionic compounds (specifically metals), can form multiple ions. To represent this, the roman numeral of the ion used has to be written after the element it is representing, so as the viewer isn't confused if the charges were balanced incorrectly. When I first learned this I thought that the roman numerals were another way to represent the charge of the entire compound, and thought it was an exception to the net charge of zero rule, so I was thoroughly confused on this topic for a day. I finally caught on and ceased to over think the formulas, however it took me a fair amount of practice problems to do so. As I initially struggled with naming ionic compounds, I found that continually practicing these problems, as well as becoming more familiar with a periodic table, was incredibly useful and helped me to master the material by the time the test rolled around. An observation I made while doing additional practice was that the metal is always written first in the formula. I'm wondering if whether or not this a universal rule that was specifically established by a scientist?

Types of ions.

The next day, our class began discussing polyatomic ions, with the lingering question of, can a group of atoms have a charge? A polyatomic ion is a charged ion composed of two or more atoms bonded together, acting as a single unit. These ions differ from monatomic ions, as a monatomic ion is an ion consisting of a single atom. My hypothesis for this question was yes, simply because it didn't make sense to me that an ion wouldn't have a charge, regardless of how many atoms it consists of. On the worksheet accompanying this topic, there was a diagram I found extremely helpful, as it gave example of multiple types of ions. They were then categorized depending on the number of atoms in the ion. Depending on the number of atoms in the ion, the name of the compound changes to fit the amount of atoms. For example, a monatomic ion labelled nitride, can either be called nitrite or nitrate, depending on the number of oxygen atoms present. I had to study this section a fair amount as I was unaware that a change in atoms caused a change in the ionic compounds name. I found that the list of ions in the green booklets we have in class helped me to pinpoint why the names were changed, specifically by looking at the charge and the number of atoms present.

Representing ions and formula units.

The next worksheet our class completed was additional practice on representing ions and formula units. We were given a two different ions, and using that information we had to draw two separate pictures depicting the ions, and then a diagram showing them combined. We also had to write out the actual name of the compound, including the prefixes we learned. Besides balancing the charges, I found that the prefixes gave me the most trouble. The prefixes range from 1-10 (mono-deca), and are used to write out the name of a compound, but only if specific conditions apply. Prefixes can only be used when the ions from covalent bonds, or a bond between two nonmetals. For instance, when combining sulfur and oxygen, the correct name of this compound would be "sulfur trioxide." Tri, or 3, is used as the amount of atoms needed for the net charge to equal zero between the sulfur and oxygen ions is 3 oxygen ions. If the sulfur was replaced with a metal, then "trioxide", or any other prefix that matches with the number of atoms needed to balance the equation, would be incorrect as a bond between a metal and nonmetal is NOT a covalent bond. I initially had trouble with this, as I kept switching around the rule for the prefixes, due to me overthinking it too much. Like before, I resolved this issue through many practice problems.

Additional practice problem.

Finishing up our unit, we were given our review packets to study for the test. The topics I felt the most worried for was the balancing of the charges, and then giving the chemical formula its correct name. To help with my test anxiety, I completed an additional practice worksheet solely based on what I wanted to focus on the most. I found that as I continued to practice these types of problems, it took me less and less time to complete them. The one thing I had to watch out for was accidentally naming an element to be either a metal or nonmetal, when its actually the opposite. I was worried about this because confusing these up could alter your entire answer, as the elements in the compound determine whether or not prefixes are necessary in the equation. Overall I thought that the test went decently for me, although I am slightly worried that I missed a few of the "naming ionic compounds" portion.

My eyes still hurt from this!

Chemistry Entry 6

Thomsons "Plum Pudding" Model

Kicking off the week, our class dove into our new unit titled Unit 6 - Particles with Internal Structure. As with every new chapter, our class received and objectives sheet filled with goals that should be met by the time the test rolls around. We also got a worksheet concerning a famous scientist J. J. Thomas, and his experiments involving cathode rays. Cathode rays are beams of electricity emitted from the cathode of a high vacuum tube, while a cathode itself, in this case, is a negatively charged electrode by which electrons enter an electric device. A cathode can also be positively charged electrode, that supplies current for an electronic device. Thomson conducted several experiments in 1897 involving these rays, intent on trying to understand electricity. In class we were given iPads to go on the website A Look Inside the Atom, to answer questions on Thomson's three experiments, and the conclusions he drew from them. Out of the three, only two were accepted by other physicists, while the other one was proven to be incorrect. Out of the two hypothesis that were deemed correct, Thomson developed a model of an atom, titled the "plum pudding" model. It shows how an atom is made up of negative charged rays zooming around in a positively charged location, who sums up Thomson's observations during his two experiments.

Set-up for the sticky tape activity.

Building off of what we learned on Monday, on both Tuesday and Wednesday we concentrated our efforts on a lab with the driving question of: Do particles have charge interactions? We went about answering this question through fairly simple means, involving tape, tinfoil, and paper. Hanging from a stand stretched equidistant from each other, we placed the mentioned materials, with an added piece of tape. Once this was set-up at our table respective workspaces, we took two more strips of tape, stuck them together, and then peeled them off of each other quickly so as to fully charge both sides of the tape. This process was also done before taping the other two pieces of tape to the handle. To test if there were any charge interactions, we moved the tapes close to the materials, and found that both pieces of tape attracted the tin foil and paper, while the tape with the same charge repelled each other, and attracted when the opposite charges were placed near the other. We also discovered that as the tape moved closer and closer to whatever we were testing it with, their attractions would become stronger which was represented by them moving either closer or farther apart from one another, depending on the charge.

On Thursday we did an activity titled "Lets Conduct Maestros". Fifteen unique compounds were spaced throughout the room, and all we had to do was measure them to see if they were either conductive or not conductive. While most of them ended up not being conductive, about 1/4 of them didn't disappoint. Our table concluded that the substances which were conductive, were all similar in

One of the compounds we tested.

Explanation on why substances
have varying amounts of conductivity.

 that they were all positively charged metal ions, found in an aqueous solution. From this activity, I also learned of a classification called "salts", which are merely positive ions from a metal combined with negative ions from a nonmetal. An example of this is sodium chloride (NaCl).

Electrolysis of CuCl(II) diagram.

Finishing up the week, on Friday (MOLE DAY) each table group recorded their own observations of a lab set-up concerning the electrolysis of copper chloride, or CuCl(II), which had been set up the previous day. Set up in a U-shaped test tube, on the right side a negative electrode gave off a current, with a positive electrode feeding a current on the other side. What I immediately noticed was that on the positive side bubbles were forming that gave off the scent of chlorine, while on the opposite side a weird pink substance was forming on the electrode surface. As I eventually learned, the "pink stuff" was actually copper forming, due to the positively charged copper sulfate and hydrogen ions attracted to its opposite charge. While this was occurring, the negatively charged chlorine ions were attracted to the positive electrode, resulting in the chlorine smelling bubbles forming on the surface. I found that collaborating as a class with our separate data results, concerning each activity we completed throughout the week, helped me to understand the material a lot more, compared to if we had continued moving on instead of sharing our results. I would like to work a little more on understanding the relationship between positive and negatively charged ions, and how the combination of different elements effects their conductivity.

Chemistry Entry 5

Simple mass & mole problems.

Two problems I initially had difficulty with.

Our class began the week with reviewing our homework we had been assigned over the weekend. These problems consisted of calculations involving moles, various elements,  and the masses of assorted objects. Each problem varied in that you were given a specific amount of information, but each time the given information was divergent, so the unknowns had to be solved differently than usual. Sometimes more than enough information was given, so the reader had to be able to understand exactly what was being asked, and had to know exactly was was needed and what wasn't to solve the problem. Personally, I had a lot of difficulty in solving problems such as these, mainly because whenever I see something given that is unneeded, I fail to grasp the fact that the information is irrelevant to the question. This is partly due to me still learning about the material and that I am collecting more knowledge about how to solve these types of problems, but I also need to get over the mental roadblock I have in terms of problems like these. What really helped me begin to figure out how to tell which information was needed or not, was white boarding the homework as a class. I heard all of the input from my fellow peers and the various ways they went about solving the homework, and was happy to overhear that I wasn't the only student struggling with this. Solving out difficult problems such as these always seems to cause a light bulb to go off in my brain, and everything seems to click all it once. I would love to continue white boarding out problems as a class, as this learning style really seems to stimulate my learning abilities.

Solving for the empirical formula.

Finding the empirical AND molecular formula's.

Our next task was a worksheet involving Empirical and Molecular Formulas. I learned that the main difference between these two is that empirical means "based on experimental data", and the molecular formula is simply a formula giving the number of atoms of each of the elements present in one molecule of a specific compound. To solve one of these types of problems, you have to be given the grams of the specific elements, and have a periodic table present so as to find the atomic number of the given elements. Once located, for each element, you multiply the given mass of the element by 1 mole, which is then divided by the elements atomic number. Once you have solved that, you compare the numbers between the elements, and whichever is smallest, you divide all them by that amount. The approximated answer of this is simply the amount of the elements present in the empirical formula equation. If you're then instructed to find the molecular formula, all you have to do is add up the atomic numbers of the given elements, and then divide that number by the molar mass of the compound.

I'm a dingbat!!

The rest of the week consisted of a quick quiz and then studying for our test on Friday. The quiz seemed easy to me when I took it, but as it turns out, I made a really stupid which earned me the title of a "Ding Bat." I accidentally had multiplied an O, mistaking it for a 0 while it actually stood for oxygen. I somehow over looked this incident and had an interesting surprise when we were handed back the quizzes. In terms of the test, I felt a lot more comfortable and confident than I felt on the quiz, with nothing surprising me or striking me as being too difficult.

Chemistry Entry 4

Mole fact.

Continuing into what we started the previous week, this week our class delved further into Unit 5. The first thing we covered was the topic of Relative Mass and the Mole, and how we could use them to count atoms by using just a balance. I was familiar with relative mass, but a mole was an entirely new concept for me, that I am still trying to understand as we progress further into using this measurement. A mole is the mass of a substance (in the case of the worksheet; oxygen and sulfur) containing the same number of fundamental units as there are atoms, in the specific number of 6.022 x 10e23. The units of a mole can range anywhere from a molecule to an atom, but they always have the same number. This number is called Avogadro's number, with Avogadro being a famous scientist we discussed about in Unit 4. Back to the worksheet, we were given a table of oxygen and sulfur atoms, and needed to find the mass of the sample, through knowing the number of atoms in that sample. This could be found through multiplying the mass of the first sample (contains one atom), by the number of atoms in each column. Once this was completed, we made a ratio between the oxygen and sulfur atoms, comparing the number and masses of the atoms. Our table came to the conclusion that, the ratio of the sample masses will be equal to the ratio of the atom's masses. We also realized that, by knowing the relative masses of the above elements, you can use them to "count" atoms, with the quantity being the mole. As the unit Mole can be confusing and  add complexity to an "average" question, the next day we burrowed deeper into this enormous number. To give somebody an idea of how massive a mole is, if we had a mole of rice grains, all the land area of the earth would be covered with rice to a depth of about 75 meters.

Heating up the substances.

Post heating.

In addition to the problems we solved regarding the unit mole, last week we did a lab called the "Empirical Formula Lab", empirical meaning based on previous experimental evidence.  In this experiment, a specific amount of zinc was mixed with hydrochloric acid. With our table groups, we had to collect data that would enable us to determine the empirical formula of zinc chloride (Zn?Cl?).  This lab took two days to complete, as the reactants needed a specified amount of time to combine. On day 1, all we had to do was find the mass of a regular, empty beaker, and then add zinc pieces to find the mass of both the zinc and bottle. Once this was complete, we had to add 3 moles of hydrochloric acid, so we could place the beaker in a safe location so as to finish the rest of the lab the next day. Picking up where we had left off, day 2 was slightly more complex, for me at least. recording the observations in day 1 just involved us to examine and interpret what was occurring, but day 2 required a little more calculation. Our workplace was set up to enable us to heat the beaker, so as to remove the unneeded water from the zinc chloride. Our only restriction was to remove the heated beaker immediately off the bunsen burner, if the contents began smoking. Once the zinc chloride was finished heating, it was interesting to see how it solidifies from its molten state. Once this process was complete, the mass of the beaker and zinc chloride had to be found twice. The idea of measuring it twice was to verify that the second mass is lighter than the first (by approximately .02 grams).

White boarding of question 8.

Finishing up the week, as a class we white boarded a worksheet involving problems with sealed sample containers. We had to find the mass of said containers as well as of its contents. A new vocabulary word I learned from this was Tare Weight. Tare weight is the containers empty mass, and is helpful as the number is needed for various calculations throughout the sheet. The questions varied, but the main idea of this was to learn how to find the number of moles or atoms in a container, given a formula. The formula of the contents is incredibly useful as you need to multiply the number of atoms in the formula, by their respected average atomic numbers, so you can plug it into the main equation once all the values are found. I found this week slightly confusing due to the mole, but also intriguing as you can apply the mole multiple chemistry concepts, which helps me to understand them more.

Naturally, I enjoyed the below demonstration you gave for our class.

I took a video and then screen shotted this!

Chemistry Entry 3

Starting off our week in Chemistry, Monday was our designated review day for the Unit 4 test, our first one this trimester. Previously on last Friday, our class acquired an objective sheet illustrating main concepts that we needed to know for the test on Tuesday. The listed concepts on the sheet were everything we had covered in the weeks leading up to the test day. Within our table groups, we filled out the boxes on the sheet to the best of our abilities without using our notes, referred back to them once everything we remembered was written down, and then white boarded the worksheet to share our answers with the class, as shown below.   

I really enjoy this type of learning as it is an easy way to place yourself in terms of understanding the material and noticing where you are in comparison to your peers. Also, through white boarding review packets like this, if you noticed you made the same mistake as someone else, than you can work together to solve it or bring it up as a class discussion. Similarly, you can scan another groups board to see if they solved any problems differently and possibly in a simpler fashion, or query on how they got a specific answer that your group might not be sure on. Either way, I recognize that this type of review session is beneficial for a quality test score, although I am still curious on what points I lost as they are yet to be handed back.

The Monday review session for our Unit 4 test was composed of filling out a review packet, similar to what was completed on Friday. This one differed from Fridays because it contained actual questions that could appear on the test but in a different format/new scenario. Using the same whiteboard form of study, we concluded our worksheet and updated our whiteboard. A concept I found interesting was that atoms can be monatomic, diatomic, triatomic, and so on. I was aware that atoms differ depending on the element, but unaware of the specific classifications and how to demonstrate the differences in a particle diagram. I had a bit of difficulty diagramming how atoms of different classifications (say monatomic and diatomic) compounded together to create a new element, but by white boarding the problem out, I had it down in no time. The only other problem I faced was was not mixing up the definitions and properties of mixtures, compounds, elements, atoms, and molecules. I continuously quizzed myself on this topic both in class and out, but I'm worried I haven't mastered it as I believe those were the scarce problems I incorrectly completed on my test.



Kicking off our new unit, like last time, our class received a list of objectives that will be taught through the course of the next unit, which is Unit 5 - Counting Particles. Our first activity had to do with a worksheet focused on Relative Mass. The purpose behind this was to determine the relative mass of different kinds of hardware and to learn to count by massing. We were given a semi-filled out data table, that contained multiple different types of hardware (bolts, washers, hex-nuts) and an empty container. We had to find the general mass, mass of 25 pieces, and the mass of 1 box. This was calculated through a scale, and guided with knowing that 1 kilogram is equal to 1,000 grams. Throughout the rest of the sheet we had to use the given information to solve problems, given various scenarios and new bits of information to add to the calculations. I definitely still need to work on problems like these as I am always either at a loss on how to approach the problem, or am on the right track but still don't understand parts of it.

As Friday was pep rally day, I had to leave early for the marching band. However, I am aware that the worksheet that was covered in the short amount of time had to do with Molar Masses of the Elements. I would like to cover this on Monday as a refresher to the kids who learned it, and as new material for us marching band nerds.

Chemistry Entry 2

Kicking off this week, our class dove right into some new material, which I am still getting the hang of understanding. One of the first worksheets we started off with involved masses of elements in pairs of compounds (compound A & B), and we had to use the given information to suggest formulas that account for the written ratios. The first problem was simple enough, and it involved a certain mass of oxygen and carbon written in a ratio, in two different compounds. Using the information provided, we had to determine the value of the ratio in compounds A & B. To find this, I had to put the mass of oxygen over the mass of carbon, divide it out, and then round the answer to the nearest tenth or hundredth, depending on the answer. As we found for this particular problem, the ratio to compound B was twice as much as the ratio to compound A. Below our ratio computing, we were given two boxes side-by-side, labeled "Hypothesis 1" and "Hypothesis 2". In the hypothesis 1 box we were told that the mass of the carbon and oxygen had to have the same mass, and in 2 we were told that the oxygen atoms are heavier than the carbon atoms by the ration in compound A. To demonstrate this, we had boxes in which to draw particle drawings to represent both compounds. Once they were drawn, we had to write out an equation describing the drawing, which could be found by looking at the amount of the atoms for each element, and then writing out the element name as represented on the periodic table, and a subscript of the number just afterwards.



Originally this part confused me, but as I learned these drawings could be reasoned out by changing the ratio we previously found into an improper fraction. By doing this, the numerator represents the amount of oxygen atoms and the denominator represents the number of carbon. Throughout the worksheet, each single problem had a layout similar to this, yet they progressively got more challenging in terms of the given elements in a compound, and the varying ratios.  I would like to continue solving problems such as these as I am not 100% confident on them, and I would like to perfect my understanding by the time the test rolls around. In continuation of the worksheet described above, as a class we began to slowly understand the main idea behind solving these problems. The problem that eventually "clicked" with our class was comparing the compounds of iron and chlorine. Together we found the ratios and compared them, and reasoned that compound B has 1.5 more grams of chlorine then A does. Once we found this, the question we all asked ourselves was "so what?" I didn't see why this information mattered, or how it was applicable to writing out an equation comparing these two compounds. Slowly but surely, it began to dawn on us (with the help of Toby and Megan), that 1.5 was the amount multiplied by A's chlorine to get B's chlorine. As I'm writing this blog and looking over this worksheet, even still I am learning more and understanding this problem. Figuring out this problem as a class really helped to improve my understanding, so hopefully we will do more of that as the year progresses.

Another piece of information I found interesting and useful, was the history behind matter. It is crazy to imagine that scientists years and years ago forged their own ideas to describe matter, and used what they learned to apply to experiments and the world in general. As years passed however, new scientists with radical, unfamiliar ideas had multiple breakthroughs. These threatened the basis of chemistry and how matter was viewed for the "old school" chemists. In particular, Democritus suggesting the idea of atoms, Bernoulli theorizing that gases consists of small particles that are loosely packed in an empty volume of space, and Priestly's experiments with red mercury calx, leading to Lavoisier's discovery of oxygen. Fortunately these theories were proved and are now the footing of what chemists build off of today. This week started off a little bit overwhelming in terms of the new material and equations, but ended up enlightening me as well as providing me with a useful background of the "chemistry revolution" and those behind the fascinating discoveries.