Lab 21: Titration

Procedure: We began the lab by rinsing out the burette(the large pipe/dropper thing in the picture) with a small amount of NaOH, then filling the burette up to the 0 mL line with 0.25 M NaOH with a funnel. Next, we added 7.5 mL of vinegar to an Erlenmeyer flask using a graduated cylinder. We repeated the same measure-and-add with 20 mL of distilled water, so as to end up with enough mixture to submerge the magnetic stirrer. We added 4 drops of phenolphthalein into the flask. We then placed the flask onto the hot plate and turned up the stir setting to 360 rpm, which caused the magnetic stirrer to spin and mix the solution. We added NaOH slowly into the flask from the burette by turning the valve, which allowed water to drip into the flask. When it reached about 18 mL, we slowed the NaOH down to very, very slow drops so we could decide when the faint pink color had appeared. Our first trial finished badly, but our second-- and last-- trial went very well. 

Percent Ionization of Vinegar: .474%

Why is this such a low number?The number is low because vinegar is a weak acid, meaning that there aren't many hydromium atoms in the solution.

Chicago River Field Trip

River Test: Nina, Christian and I took charge of the Dissolved Oxygen test. Dissolved oxygen is oxygen gas dissolved in water. This test checked the amount of oxygen in the river available for aquatic animals such as fish and other invertabrates. Without this dissolved oxygen, aquatic life would not exist.

Procedure: For maximum efficiency, we took half a bucket sample of the river water and used it for all of the tests. First, we took the temperature of the water, and before we began the actual testing, we rinsed all the bottles and tubes we would be using. We then filled the oxygen bottle with river water and added Dissolved Oxygen 1 Reagent Powder (manganous sulfate) and Dissolved Oxygen 2 Reagent Powder (alkaline iodine-azide reagent), and stoppered and left it for about 5 minutes so as to let the precipitate settle. Afterwards, we added Dissolved Oxygen 3 Reagent Powder Packet (Sulfamic Acid) and put in the glass stopper again. After shaking the bottle vigorously, the mixture turned a bright yellow, signifying the presence of bright yellow. We then filled a small plastic tube with the solution, and then poured that amount of solution into a square mixing bottle, where we added Sodium Thiosulfate Standard Solution one drop at a time until it became colorless. The total drops used was the total ppm of Dissolved Oxygen.

Calculations: We used our ppm to find the percent saturation, which we then used to find our q value. The Q value was about 93 out of 100. This was very good. It was also pretty close to the overall Water Quality Score, which was 92.1. This was found as a total of all the test results. (Mostly neutral pH value, good phosphate value, no nitrate, good turbidity, etc.)

What We Can Do: Things we can do vary from simply minimizing water usage by turning off the tap while brushing your teeth and washing the dishes, or the washing machine and dishwasher only when they are fully loaded, etc. to helping protect and preserve wetlands and prevent invasive species from entering the lake.

Lab 17: Calories in Food

Summary: In this lab, we burned three different types of foods: cheese puffs, brazil nuts, and cashew. The purpose of this was to find the number of the Calories per gram of different food items. We did this by lighting the snack on fire, placing a soup can calorimeter over the piece, and then placing a flask of water on top of the can. The calorimeter would keep the heat of the burning food contained, and the water in the flask above would absorb the heat. We kept track of how much its temperature increased from the absorbed heat using a thermometer. Afterwards, we calculated the amount of energy in the food by measuring the heat absorbed by the water, which was equal to the heat lost by the food.

Data:

Food	Initial Mass of Food (g)	Final Mass of Food (g)	Mass of Water (g)	Tinitial (°C)	Tfinal (°C)
Cheese Puff	0.960	0.150	25.0	22.1	42.3
Brazil Nut	3.86	1.70	25.0	24.4	60.2
Cashew	1.47	0.130	25.0	28.9	60.3

Food	Mass of food burned (g)	Change in Temperature (°C)	Q (calories)	Q (Calories)	Q / Mass of food burned (Cal/g)
Cheese Puff	0.810	20.2	505	0.505	.623
Brazil Nut	2.16	35.8	895	0.895	0.414
Cashew	1.34	31.4	785	0.785	0.506

Questions:

1. Did you measure a temperature change in the food sample or the water?

The food sample was impossible to measure; even if we had tried it most likely would have been inaccurate as it may have been exposed to the room temperature. In general, it would have been very hard to "submerge" the thermometer in the solid. We instead measured the temperature change in the water by placing a soup can calorimeter over the burning food and placing a flask of 25 mL water on top of it in order to indirectly heat the water. We placed a thermometer inside the beaker as well to let us track the starting and final temperatures.

2. Did you measure the energy released by the food sample or the energy gained by the water?

We measured the energy gained by the water, as it was easier to track than the energy the food sample released. As heat is related to energy, and temperature is related to heat, the fact that measuring the temperature of the water was done much more easily than had we tried to measure the food sample's temperature shows that measuring the energy released by the food sample would've been just as hard.

3. What happens to the small amount of energy that is not absorbed by the water?

We did the lab based on the assumption that "all energy was absorbed into the water," even though it was obviously not possible. The small amount was negligible, however, which is why the assumption was made. As energy cannot be destroyed, the small amount of energy not absorbed by the water was instead released into the air. 

4. Were you surprised by any of the results? Explain.

I was surprised the nuts had more calories than the cheese puffs; nuts have always seemed to be more "healthy" than snacks such as cheese puffs, but the latter had more than 200 less calories than either the cashew or the brazil nut.

Lab 15: Evaporation and Intermolecular Attractions

Calculations and Results:

2. Explain the difference in temperature (ΔT) of these substances as they evaporated. Explain your results in terms of intermolecular force.

The differences in temperature varied for the five substances. Methanol's temperature went down 13.4, while glycerin actually went up 2.5. This was caused by the difference in intermolecular forces. Methanol, whose temperature went down the most, had the weakest bonds; it had just one hydrogen bond and three London Dispersion bonds. (Hydrocarbons form only dispersion bonds, and any N-H, F-H, or H-O compounds have hydrogen bonds). On the other hand, glycerin, the only substance whose temperature increased, had three hydrogen bonds and five London Dispersion bonds. (Further explained in the next two questions).

3. Explain the difference in evaporation of any two compounds that have similar molar masses. Explain your results in terms of interolecular forces.

It is possible for compounds of similar molar masses to have different changes of temperature. What we were testing was the change of temperature as a result of evaporation from the room temperature; evaporation is liquid becoming gas through energy overcoming intermolecular forces to break apart the bonds holding molecules together. Stronger intermolecular forces means slower evaporating rates because the energy used must be greater in order to break the bonds. Methanol and Ethanol had the closest molar masses, but ethanol's temperature decreased less than methanol's, and had stronger intermolecular forces.

4.  Explain how the number of -OH groups in the substances affects the ability of the tested compounds to evaporate. Explain your results in terms of interolecular forces.

OH forms hydrogen bonds. These are the strongest type of intermolecular forces, excluding ionic bonds. Hydrogen bonds are stronger than london dispersion bonds, which are what hydrocarbons (HC) forms. As a result the number of OH groups results in more evaporation and greater decrease in temperature. In fact, glycerin, which had three OH groups, actually went up in terms of its temperature. Methanol, on the other hand, had just one OH group, and evaporated the most, causing its temperature to decrease. As stated before, stronger intermolecular forces means slower evaporating rates because the energy used to break the bonds must be larger. 

Lab 14: Create a Smell Lab: Ester Synthesis

Summary: In this lab, we combined different acids with alcohol in order to synthesize, or produce, three esters with various smells. Our goal was to produce a sweet-smelling molecule from a putrid, foul smell. What we added and what we ended up with are shown in the table below:

Microcentri- fuge Tube	Alcohol	Acid	Odor of Ester
I	Isopentyl Alcohol	Sulfuric Acid	Artificial Banana Flavor
E	Ethyl Alcohol	Glacial Acetic Acid and Sulfuric Acid	Nail Polish
M	Methyl Alcohol	Salicylic Acid and Sulfuric Acid	Mint / Peppermint Gum

 Analysis Questions:

1. Compare the odors of the three esters produced to the odors of the starting materials. How are they different?

The starting materials either smelled like nothing or were very pungent-- in a bad way. The alcohols smelled like alcohol or like medicine, and the acids were somewhat strong and putrid. On the other hand, the produced esters all had a kind of sweet smell, other than the nail polish. The one produced in tube M smelled like mint gum while the one in tube I was similar to the banana flavor in, for example, laffy taffy. Even the nail polish smell in tube E was better than other scents from the starting materials.

2. Based on the smell of the mixtures after heating, what functional group must be present in the final molecules that were produced? Draw it.

The functional group that was definitely in the final molecules was the esters. The prelab states that "in order to transform a putrid smelling molecule into a sweet-smelling molecule, it is necessary to transform a carboxylic acid into an ester molecule..." As the end result was sweet, we can safely say that there was ester in the final product.

3. Were the esters easily identified as a specific fragrance, i.e., apple or banana? In the case where a specific fragrance was detected, how does the odor compare to the natural fragrance?

The smells detected were definitely things smelled before, though they also resembled things I had never smelled before (my partner Daniel stated that the ester produced in tube M smelled like root beer, which I'd never had before, while I immediately stated that it was similar to peppermint gum). At some points it took some thinking to recall what the fragrance was, but in th end we were able to identify the smells. The smells all resembled the natural fragrances very closely. Had I closed my eyes, I could've imagined being at home with my sister painting her nails as I opened tube E. However, tube M was not exactly the same smell as mint gum, and the banana artificial flavor was not very similar, especially as I had to look for it (aka someone had to say that it resembed the fragrance) to find it. Admittedly, it was pretty similar once I had found it. All esters definitely smelled much better than the starting materials, however.

Lab 12: Electron Configuration Battleship

Biggest Challenge: The biggest challenges I faced included finding Nina's ships. In fact, this took so long we din't actually get to finish. It was also difficult trying to find and say the electron configurations quickly, another thing that made us unable to finish.

What I Learned: I learned how to find electron configurations quicker and more accurately.

Lab 11: Flame Test Lab

Pre-lab Questions:

1. What is the difference between ground state and an excited state?
An atom or ion in its ground state is in the lowest energy levels possible, and is stable. An atom or ion in its excited state, however, has "jumped" to a higher energy level by absorbing energy, and is unstable.

2. What does the word "emit" mean?
Oxford dictionaries defines "emit as "[To] produce and discharge (something, especially gas or radiation)." In other words, to emit means to give off, or to release. When the electrons "fall back" to their normal positions of lower energy, they emit electromagnetic radiation, which includes visible light.

3. In this experiment, where are the atoms getting their excess energy from?

In this lab, the atoms are getting their excess energy from the Bunsen Burner's flame. The flame heats the atoms to high temperatures, which allows them to absorb energy and therefore reach higher energy levels.

4. Why do different atoms emit different colors of light?
Atoms of different elements have different number of "shells" of electrons, and each shell contains a different amount of energy. When they are heated, some electrons absorb energy and "jump" to shells with higher energy. Because this is an unstable state, however, the electrons eventually settle back to their lower state of energy, emitting the absorbed energy as photons, which are little "packets" of light. The amount of energy in the photon determines the wavelength of the light. Different elements have shells of different energies, so the photons released give different wavelengths. As colors are visible light with different wavelengths, the atoms give off different colors of photons depending on the energies of the different shells.

5. Why is it necessary for each station to have separate wooden splints for each inividual flame test?

We used different wooden splints for each individual flame test so that the different compounds wouldn't mix and cause unreliable and possibly invalid results.

Purpose: To identify two unknown chlorine compounds using flame tests.

Summary: To fulfill the purpose of this lab, we took (separate) wooden splints soaked in 8 different chlorine compound solutions and held them over a Bunsen Burner flame, and then noted the flame appearance with the naked eye. We compared these flame colors to the flame colors of the two unknowns to determine their identities. 

Result: We found the identity of Unknown 1 to be SrCl2, or Strontium Chloride, and Unknown 2 to be KCl, or Potassium Chloride. Unknown 1 had a red-pinkish color similar to bright red or strawberry. We matched this to both Lithium Chloride, which had a rich ruby-like, hot pink to magenta-like color, and Strontium Chloride, which was more red than pink, but after redoing all three flame tests, we decided that the color was closer to the latter's. Unknown 2 was relatively easier to determine compared to Unknown 1. The flame test caused a purple, lavender-like color that only one other compound had: Potassium Chloride.

Lithium Chloride

Sodium Chloride

Strontium Chloride

Copper Chloride

Calcium Chloride

Potassium Chloride

Unknown #2

Unknown #1

Lab 10: Mole-Mass Relationships Lab

Summary: The purpose of this lab was to better understand the mole-mass relationship in a chemical reaction and the interpretation of a balanced chemical equation through a chemical reaction involving sodium hydrogen carbonate and hydrochloric acid. We used about 2 grams of NaHCO₃ and a large source of HCl and slowly added the latter into the former until the reaction had stopped. We then heated the mixture on a hot plate until all that was left was salt NaCl from the three products NaCl, CO₂, and H₂O. Next, we compared this yield to a theoretical yield calculated through stoichiometry. Finally, we took both yields to calculate the percent yield.

Reaction between NaHCO₃ and HCl:

Possible Problems: Errors most likely resulted from the time we left the dish on the hot plate and how much acid we used.

Calculations:

Lab 9: Composition of a Copper Sulfate Hydrate Lab

Data: In moles, the amount of water that evaporated was 0.017 moles. We got this by dividing the grams of water evaporated by the molar mass of water. Similarly, the moles of Copper Sulfate, the anhydrate, that was left in the evaporating dish was the grams of Copper Sulfate left divided by the molar mass of Copper Sulfate, or 0.0035 moles. To find the empirical formula, we needed the ratio of moles of CuSO₄ to H₂O. We did this by dividing by the smallest amount, which was the moles of CuSO₄. We got a ratio of 5 water to 1 Copper Sulfate. As such, the empirical formula was CuSO₄ * 5H₂O. Our percentage of water was 36.9%, less than a tenth of a percent away from the accepted value for the percentage of water. Our percent error was, as such, 2.85%. This means our coefficient is very close to the actual. It is probably a little high, as our experimental value was as well.

Hydrates before and after heating (from left to right):

 Calculations:

Lab 8: Mole Baggie Lab

Summary: The purpose of this lab was to determine the substances in two Ziploc bags with only the mass of the empty bag, the substance's representative particles or the number of moles of the substance, a balance, and a calculator. The first bag given to us was bag B3. The first thing we did was find the molar mass of each of the possible compounds so we could compare them to our answers from our calculations. We then converted the representative particles to moles; this was done by dividing the R.P. by Avogadro's number. Next, weweighed the bag to get the mass of both the bag and the substance, and then subtracted the mass of the empty bag to get the mass of just the substance. Afterwards, we divided the result by the moles of the substance in the bag. We had to do this process twice because of a faulty balance that, on the first try, led us to an answer much less than the actual. Bag A4 was much easier; we used a correct balance this time, and again subtracted the mass of the empty bag from that. We first divided the result by the number of moles of the substance, the latter of which was given, and then compared it to our possible compounds.

Result: Bag B3 was Potassium Sulfate, and bag A4 was Calcium Carbonate.

Lab 6: Double Replacement Reaction Lab

Well Plate:

Balanced Equations and Net Ionic Equations:

Biggest Challenge: The most challenging part of this lab was balancing the equations, specifically figuring out which reactant or product needed a coefficient or parenthesis or subscripts.

Lab 5: Nomenclature Puzzle

Summary: This "lab" was to use our knowledge of the nomenclature of ions, acids, and covalently-bonded molecules to match their names to their formulas. This was done as a puzzle; we had to match a total of 180 sides-- 90 names, 90 formulas-- into the shape of a large square. Our finished puzzle is shown below.

Challenge: The biggest challenge my partner Nina and I faced while doing this lab was finding each corresponding piece. Sodium, Barium, Iron, and Potassium showed up often on pieces of the puzzle, and to find the correct piece took a lot of time and patience, especially as sometimes we would look for one and find it 10 minutes later, and realize as we look for the first piece that we had somehow lost it. This happened 5-6 times for me. Sometimes, a name would come up that we just could not find the formula for in the reference packet. This problem also was an obstacle when we had formed two large parts of the puzzle and could not figure out where to put them together; it was only with the assistance of Saloni that we could figure this out, in the last two minutes of class.

Contribution: One of my biggest contributions was finding the corresponding names and formulas. I was able to find the names and chemical formulas in the reference packet pretty often, and was able to use this to finally complete the puzzle by putting our two separate big solved parts of the puzzle together. This was greatly helped by my ability to "strategize," meaning choose priorities; I first went after the ions with uncommon atoms, such as the Palladium (II) Iodide; Palladium showed up only once, and Iodide around 2-3 times; this meant it would stand out. I ended up finding this piece almost immediately after I looked for it. This worked a few times, but definitely helped in finishing the puzzle in the 45 minutes left in class.

Lab 4: Atomic Mass of Candium

Purpose: The purpose of this lab is to find the average atomic mass of the element candium using an arbitrary sample of the three candium isotopes: regular M&Ms, peanut M&Ms, and pretzel M&Ms.

Average Atomic Mass: We found the average atomic mass of candium to be 1.76456 amu.

Lab Questions:

1. Ask a group nearby what their average atomic mass was. Why would your average atomic mass be different than theirs?

The group next to us had an average atomic mass of 1.7 amu. Our average atomic masses were different because we had different samples; it was as if we had different "planets," and thus had different "abundances" of the element in them. For example, they could've had more regular M&Ms, which have a smaller mass. They could also have had more small peanut or pretzel M&Ms, as each does not have a constant mass.

2. If larger samples of candium were used, for example if I gave you a whole backpack filled with candium, would the differences between your average atomic mass and others' average atomic masses be bigger or smaller? Defend your answer.

The differences between our average atomic mass and others' would most likely get smaller with a larger sample, as larger sample sizes usually means less variation in the data.

3. If you took any piece of candium from your sample and placed it on the balance, would it have the exact average atomic mass that you calculated? Why or why not?

There is a very tiny chance that I would have the exact average atomic mass I had calculated. A data point is almost never the actual average; more likely, it would be somewhere very close to my average. The average is, after all, the total atomic masses-- outliers and all-- summed up and then divided by the total number.

4. You have been chosen to design the square that will be placed on the periodic table for candium. Draw a rough sketch of what your square would look like. Include an atomic symbol for candium as well as the average atomic mass for candium.

Lab 3: Chromatography Lab

1. Why is it important that only the wick and not the filter paper be in contact with the cup?

Should water come in contact with the filter paper, then the ink and water may not spread out from the center to the edges, and therefore not create the radial pattern that did result from this lab.

2. What are some of the variables that will affect the pattern of colors produced on the filter paper?

Different inks resulted in different colors; for example, LiquidFlair resulted in only black, but the thin Expo marker created bright shades of blue. Other variables included the pattern drawn in the center. The longer it was vertically, the farther the water and ink ran outwards. Of course, this also includes the amount of ink on the paper. When I used a thick Mr. Sketch marker and drew bold lines, the ink spread very far, versus when I put a circle of tiny dots, which just spread a little. The distance from the center also made an impact, because the ink spread outwards only, and not inwards. 

3. Why does each ink separate into different pigment bands? 

Chromatography is defined as "the separation of a mixture by passing it in solution... through a medium in which the components move at different rates." As the focus of our lab today was on this, we separated a mixture through a medium. The mixture was the ink, and the medium was, of course, the water. When the water reached the ink and passed through it, it separated the ink into the different pigments that made it up; in other words, it divided the homogenous mixure that was the ink into all of its different components, or pigments. This resulted in bands of these pigments. Each color pigment is unique in its physical properties; these differences show in their ability to travel further or not as far on the filter paper. One property would be its water solubility; if the pigment is highly water soluble, it will "float," or remain in the water for longer, and therefore travel more. If its water solubility is not high, it will not stay in the water for long, and will create a pattern closer to the center. 

4. Choose one color that is present in more than one type of ink. Is the pigment that gives this color always the same? Do any of the pens appear to contain common pigments? Explain.

Blue appeared in many different inks, in different forms. For exmple, we had purple on one filter paper, which is just red and blue, and blue by itself in another. Green, a mixture of blue and yellow, was also present in another pen's ink. As such, we can say that blue was present in more than one type of ink. The pigment that gives this color is not always the same; at one point one chromatogram had dark blue on it, but another had a cyan blue color. I did not see any pens that appeared to contain common pigments.

5. Why are only water-soluble markers or pens used in this activity? How could the experiment be modified to separate the pigments in "permanent" markers or pens?

Water-soluble markers and pens were the only ones used in this activity because this property allowed the ink to be separated and carried by the water to spread and create patterns. Had permanent markers been used, the pigments would not have separated and created pigment bands, and as such the experiment would have failed. A good example of this was Nina's chromatogram, in which she'd accidentally used a not water-soluble pen to draw on it. About twenty minutes later, the water-soluble pens she'd used had successfully spread, but the lines drawn by the non-water-soluble had not changed. To separate the pigments in "permanent" markers or pens, it would be best to switch the medium to something that the ink in the "permanent" markers could be separated by; in other words, water should be swapped for something this specific ink is soluble in.

Lab 2: Aluminum Foil Lab

Procedure: This lab tasked us with determining the thickness of a "square" piece of aluminum foil in millimeters. We were given the density and we were given a ruler, but nothing else. Our first step was to use the ruler to measure our piece of aluminum foil. The length and width respectively were 11.30 cm and 10.10 cm. We then weighed the aluminum foil on the balance to find its mass, which was 0.50 g. With this value, we could solve for volume using the formula for density: D = M/V. This equation also meant that volume was mass divided by density. By plugging in the values we already had, we found that the volume was 0.19 cm^3. We were also given the formula for volume, which was V= L x H x W. As the first thing we did was measure the length and width, we could plug in values for every variable except for H, the height, which was what we were looking for. We solved this equation to get a height, or thickness, of 0.002 cm. We converted this number into millimeters, resulting in a final answer of 0.02 mm.

Lab 1: Density Block Lab

Introduction: The objective of the Density Block Lab was to calculate the mass of a block using only its density and a ruler. The density was given to us along with the block; specifically, it was labeled on a sticker on the block. The ruler was used to find the volume, as the equation for volume is l*w*h. Important terms in the lab include mass, which can be defined as the amount of matter in an object; volume, which means the amount of space an object occupies; and density, which is mass per unit of volume. 

Procedure: My partner Nina and I first recieved a block from the instructor. We measured the length, width and height of this block carefully (to three significant figures) and then found the volume by multiplying them together. We then used the density already given to us to find the mass; as density is equal to mass over volume, mass is equal to volume times density. Using the values we'd acquired, we found our calculated mass. Next, our instructor weighed our block, giving us our actual mass. With these two masses we calculated the percentage of error. We repeated the procedure only if the percentage of error was more than 5.

Data: We were able to get the calculated mass within the 5% margin for error on our first try. The calculated mass was 137.5 g, and the actual mass was 144.7 g, allowing us to scrape by with a percent error of 4.98%.

Conclusion: We were able to fulfill the purpose on our first try, although it was very close. There was some definite errors. My partner and I had issues with measuring accuracy; I in particular rounded most things that I thought were too small to eyeball to the nearest millimeter, resulting in a smaller number than had it been done with more care. As such, that is one thing I will do differently. Another, however, would be double-checking those values with my partner, and in general communicating more with them about the process. The main lesson we learned in this lab is how to calculate something's mass with just density and dimensions. We also learned to make measurements as precise as possible, and to communicate more with our partners about what we are doing.