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