Bacteria Eats Plastic

The Development of Bacteria, Capable to Degrade Plastics on Prezi

BEER'S LAW

In this lab, we demonstrated Beer's Law, which states that the more concentrated a solution is, the darker it is, since it has more of the solute dissolved in the solvent.

Purpose: To determine the concentration of unknown nickel (II) sulfate solutions as determined by the absorbance measured by the colorimeter hooked up to the mac book through logger pro.

Materials:

-Nickel nitrate solution (0.4 M)

-Test tubes
-Test Tube Rack
-2 pipets
-Cuvette
-Vernier Colorimeter
-LoggerPro program
-Distilled Water



Procedure:

1.) Goggle time! Put those goggles on people, Caution when using NiSO4 solutions.

2.) Add distilled water to a cuvette and insert it in the colorimeter, then set the colorimeter to no light and calibrate the unit.

 ( Mr. Ludwig-Why dont we want it to be on green if the solution is green?
Kandace-We wont see anything!)<-- See i was here for that conversation!!! :)
2) Set the colorimeter to red light not green!

3) To begin with add 2 mL of  nickel nitrate solution to a test tube. To reduce add 8 mL of distilled water to the test tube. Stir thoroughly to make the solution concentrated. Pour solution into a cuvette. Insert cuvette into the colorimeter and record the absorbance.

4) Next, add 4 mL of the nickel nitrate solution to a test tube, this time add 6 mL of distilled water to the test tube and of course stir thorough it properly. Pour this concentrated solution into a cuvette as well and insert the cuvette into the colorimeter and record the absorbance.

5) Continuing on, add yet another 6 mL of nickel nitrate solution to a test tube. After that add 4 mL of distilled water to the test tube and stir thoroughly. Pour the solution into a cuvette. Insert the cuvette into the colorimeter and record the absorbance.

6) Add 8 mL of the nickel nitrate solution to a test tube. Add 2 mL of distilled water to the test tube. Stir thoroughly. Pour solution into a cuvette. Insert cuvette into the colorimeter and record the absorbance.

7) Add the pure 0.4 M nickel nitrate solution to a cuvette. Insert cuvette into the colorimeter and record the absorbance.
8) Set a line of best fit that runs through the middle of your data to have an average of the data plotted.
9) Add the 1st solution with one of the unknown concentrations to a cuvette. Insert cuvette into the colorimeter and use the line of best fit to determine its concentration.

10) Add the 2nd solution with one of the unknown concentrations to a cuvette. Insert this cuvette into the colorimeter and use the line of best fit to determine its concentration.

11) Add the 3rd solution with an unknown concentration to a cuvette. Insert the cuvette into the colorimeter as well and use the line of best fit to determine its average concentration level.



Steven's Groups Data! Thank you Steven!!
Absorbance of the solutions are:

Unknown 1: .186

Unknown 2: .551

Unknown 3: .367

Concentrations of the solutions are:

Unknown 1: .155 M

Unknown 2: .365 M

Unknown 3: .26 M

Conclusion:

Although I did not actually do this lab I did have to research a little to actually understand what was going on within this experiment. And thanks to Steven and his group I was able to get the data they had collected while conducting this experiment. In conducting this experiment it became aparent that in order to determine the concentrations of the given solutions, each of the group's had to use Beer's Law acurately. By using Beer's Law each group was able to compare the absorbancies of each of the solutions.

Silver and Copper Lab

Silver

Copper

 Replacement

  Lab

- Soooooo I was gone for majority of this lab, but hey! I'm going to try my best thanks to some help my friend Serena!

Materials:

-Silver Nitrate

-Copper

-Distilled Water

-Test Tube

-Filter Paper

-Funnel

-Beaker

Procedure for Day #1:

Step 1-Obtain 30cm of bare copper wire. Clean the wire, then coil it around a pencil to form a loose spring. Make sure that the other end of the wire reaches the top of your test tube and is uncoiled.

Step 2-Weight the coil with a balance, then place the copper wire in the test tube to check that is is the correct length. Remove the coil and set it aside for later.


Step 3-Weight the weighing dish of silver nitrate and record it's number.


Step 4-Transfer the contents of silver nitrate to your test tube.


Step 5-Pour distilled water into your test tube until the water is about 2 cm from the top.


Step 6-Cover the top of the test tube with parafilm.


Step 7-Place your thumb on top of the test tube and invert it until all the silver nitrate is disolved.


Step 8-Weigh the empty weighing dish so that you can determine the mas of silver nitrate that was added to the test tube.


Step 9-Add the copper coil to the test tube.


Step 10-Set the tube aside until the next day or class period.


Step 11-Make sure you record all data!

Procedure for Day 2!! (I was here for this part!)

Step 1- Weigh a piece of fiter paper for use in separating the silver.



Step 2- Shake the test tube and copper wire to dislodge the silver.

Step 3- Set up a funnel with your filter paper in it.


Step 4- With a waste beaker beneath the funnel, lift the copper wire out of the test tube and hold it over the filter system.


Step 5- Using a water bottle, let distilled water run down the wire so that any silver will wash into the filter.


Step 6- Lay the copper wire on a labeled piece of paper to allow it to dry.


Step 7- Pour off the solution in the test tube through the filter paper into the filter, trying to keep the silver precipitate in the test tube.


Step 8- Rinse with distilled water and decant several times into the filter to wash the silver.


Step 9- Finally wash the silver from the test tube onto the filter paper.


Step 10- Allow to dry overnight.

Procedure for Day 3!!

Step 1- Weigh the copper coil and record it's mass.



Step 2- Weigh the silver and filter paper - record the mass.

Numbers, Numbers, Numbers!!! Data, Data, Data!!!! Conclusion, Conclusion, Conclusion!!!!!

Silver Nitrate's Mass: 1.369 grams
Mass of the Copper Coil Before: 3.68 grams
Mass of the Copper Coil After: 3.26 grams
Mass of Copper Reaction: .423 grams
Filter Paper Mass: 1.389 grams
Filter Paper & Silver: 1.892 grams
Mass of Silver Produced From Reaction: .503 grams
Silver Production in MOLES!: .0032 moles
Copper Consumtion: .0036 moles

EQUATION TIME AGAIN!
2AgNO3 + Cu --> Cu(No3)2 + 2Ag

Evaporation and Intermolecular Attractions

Lets Begin!!!
Evaporation and Intermolecular Attractions
Yes that's ATTRACTION not Reaction! It's time to get down to the relationship status!

1st Come the Materials:

• Computer (DUH! OBVIOUSLY!)
• Serial Box Interface ULI
• Data Logger
• Two Probes
• 6 Pieces of Filter Paper
• 2 Small Rubber Bands
• Methanol ( Methyl Alcohol)
• Ethanol (Ethyl Alchohol)
• 1-Propanol
• 1-Butanol
• n-Hexane
• n-Pentane

Procedure, Procedure, Procedure! This Is What it Comes Down TO!:

1. Open "Experiment 9" from the Chemistry with Computers from the experiment files of Logger Pro.
2. Wrap filter paper around the tips of probes 1 and 2, and secure them with the rubber bands. By the way this was the hardest part!!!!
3. Place both probes into a container holding a sample of either methanol or ethanol.
4. After the probes have been in the liquids for around 30 seconds, begin data collection. Monitor the temperature for around 15 seconds to establish the initial temperature of each liquid.
5. Simultaneously remove the probes from the liquids and tape them with the tips of the probes extended of the edge of the counter.
6. When both temperatures have reached minimums, stop data collection, it kinda stops itself...
7. Find the maximum and minimum temperatures, then subtract them to find the change of temperature during evaporation.
8. Remove the rubber bands and discard the used filter paper.
9. Predict, using the data just collected, how the next set of substances will react!
10. Repeat these steps for the final substances as well!! :)

Data Analysis:

1.  Most of these substances contain hydrogens bonds which is very apparent for 1-Butanol and n-pentane. Their overall weights were very similar, however the temperature change between the two was vastlly different. Can you say hydrogen bonds were the cause! Definitely! 1-Butanol contains hydrogen bonds, but n-pentane does not. This made it easier for n-pentane to evaporate compared to 1-Butanol.
2. Methanol had the weakest intermolecular forces out of the six tested. However, 1-Butanol had the strongest. We can see that methanol had the weakest bonds because of the change of temperature. Methanol changed 17.38 degrees. This shows that it had weak bonds, allowing it to evaporate more quickly than the other liquids. 1-Butanol had the strongest bonds because its temperature had a low change rate. It's temperature only went down by about 2.41 degrees. This means that it had stronger intermolecular forces which didn't allow it to evaporate as much or as quickly as methanol. So basically the lower the change the better and the stronger intermolecular force the substance has!
3. Alkanes! n-Hexane had the weakest intermolecular forces, and n-pentane had the strongest molecular forces. n-Hexane changed a total of 10.00 degrees, and n-pentane only changed 7.28 degrees.

Green-Ethanol & Red-Methanol

Red- 1-Propanol & Green- 1-Butanol

Red- n-Pentane & Green- n-Hexane

Dry Lab

A dry lab?  Doesn't that just make you think the ingredients or materials are dry? well thats what i thought until Mr. Ludwig explained that this lab was actually very serious and dangerous and we would not actuallly be performing the procedure. In this lab we would simply be following what others have already done, and simply analyzing the conclusion and results.

PROCEDURE TIME!!!!

1. Clean and dry two beakers
2. Record the number of you chemical- Lab #5 ( I had either 4 or 5 I cant really remember which! Sorrys!)
3. Find the mass of lead(II) nitrate- 1.03 grams
4. Find the mass of potassium iodide- 1.33 grams
5. Add about 50mL of DI water to each beaker
6. Stir the beakers until both chemicals have dissolved
7. Combine the contents of both beakers into one beaker
8. Wash any remaining solution into the one beaker
9. Stir the solution
10. Measure the mass of filter paper- 1.39 grams
11. Filter out the precipitate into the filter paper
12. Wash any remaining precipitate into the filter paper 
13. Allow the precipitate and filter paper to dry overnight
14. Find the mass of the precipitate- .95 grams

QUESTION TIME!!!

Question: What is the balanced equation for the chemical reaction that occured in the experiment?

Answer: The balanced equation for this particular chemical reaction is: 2KI + Pb(NO3)2 --> 2KNO3 + PbI2 (Thank goodness for notes! I'm so slow at balancing equations!)

Question: What was the limiting reagent for your experiment?
Answer: The limiting reagent for this experiment was: Lead (II) Nitrate

Question: How much lead (II) Iodide will theoretically be formed?
Answer: Theoretically speaking there should be: 1.43 grams of lead(II) iodide formed from the experiment

Question: How much lead (II) iodide did you get from your experiment?
Answer: However when i performed the experiment there was actually: .96 grams of lead(II) iodide formed from the experiment

Question: What was the Percent yeild for your experiment?
Answer: The percent yield of lead(II) from the experiment was: 67%

Empirical and Molecular Formula Quick Overview!

Empirical and Molecular Formulas on Prezi

Book Problems!!

Yes that is me with all my beautifully designed problem cards I did for Mr. Ludwig! Mr. Ludwig told me it would be a lot of extra work to put all these problems on individual little cards, but it actually kept my work organized and help me keep the problem clean and accurate. These problems came from chapter 11 in our Glenco Chemistry book.  The problems are 42-57 and 60 to 64 I do believe. These problems  have covered subjects that included the empirical formula, moles, and the molecular formula. 

Element Identification

Element Identification

So in this lab we had a procedure on google docs, provided by none other then Mr. Ludwig and his handy dandy technology equipment! We were to calculate the atomic mass of different elements in order to define which was which!

Here is how we did it…

Element A

                  .163mol / 6.52g × 6.02 × 10²³ / 1 mole = 6.52/.163 = 40

Element = Calcium

Element B

                  .91mol / 24.51g ×  6.02 × 10²³ / 1 mole = 24.51/.91 = 26.93

Element = Aluminum

Element C

                  .160mol / 9.45g × 6.02 × 10²³ / 1 mole = 9.45/.160 = 59.06

Element = Cobalt

Element D

                  .268mol / 53.87g × 6.02 × 10²³ / 1 mole = 53.87/.258 = 208.79

Element = Bismuth

Element E

                  .220mol / 14.25g × 6.02 × 10²³ / 1 mole = 14.25/.220 = 64.77

Element = Zinc

Element F

                  .756mol / 49.25g × 6.02 × 10²³ / 1 mole = 49.25/.756 = 65.145

Element = Zinc

Element G

                  .492mol / 28.65g × 6.02 × 10²³ / 1 mole = 28.65/.492 = 58.23

Element = Nickel

Element H

                  .430mol / 89.11g × 6.02 × 10²³ / 1 mole = 89.11/.430 = 207.2

Element = Lead

Element I

                  .381mol / 21.28g × 6.02 × 10²³ / 1 mole = 21.28/.381 = 55.85

Element = Iron

Element J

                  .259mol / 30.66g × 6.02 × 10²³ / 1 mole = 30.66/.259 = 118.37

Element = Tin

Popcorn Lab!

Pop, Pop, POPCORN!

Procedure:

1. Add enough vegetable oil to cover the bottom of the beaker. Make a tight cover for the beaker for the aluminum foil. Weigh the beaker as it is and record the weight

2. Add 25-30 kernels of popcorn to the beaker and replace the lid. Weight the beaker again, with the popcorn in it

3. Poke several holes in the aluminum foil and place it on the ring stand over the Bunsen Burner

4. When the popcorn is finished popping, remove the heat and carefully take off the aluminum cover

5. Let the beaker stand until it is cool. Once it has cooled, weigh the beaker with the popped corn and the cover on

!!!!RECORD THE DATA!!!!

I did this lab with Mr. Steven and our results were…

Beaker and Oil	192 Grams
Beaker, Oil, and Corn	198.5 Grams
Beaker, Oil, and POPcorn	197.5 Grams

Conclusion:

Well overall in this lab I wasn’t here for the first day to get all the yummy popcorn we could actually eat, but hey that’s okay! However I did learn that popcorn loses water mass as the kernels are heated up and reach the point of popping. That’s why when you place a bag of popcorn in the microwave it is always more heavy going in then it does coming out. Water mass is what keeps the kernels from burning when they are being heated to pop. So results showed that our popcorn lost around about 1% give or take of its water content!

Baking Soda Lab!!!

Baking Soda Lab

Purpose: To find the experimental mole ratio of the reaction of baking soda with vinegar

Background: Baking soda is pure sodium hydrogen carbonate (NaHCO₃) Vinegar is an aqueous solution of acetic acid (HC₂H₃O₂)

Balanced Equation: NaC₂H₃O₂ + HC₂H₃O₂ à CO₂ + H₂O + NaC₂H₃O₂

Mole Ratio of Balanced Equation: 1 to 1

Materials: 20ml vinegar in a large pipette

                  Balance

                  100ml beaker

                  1 Gram Baking Soda

Procedure:

1. Get a large plastic pipette filled with vinegar. Measure the mass of the pipette and record the mass

2. Measure the mass of an empty, clean 00 ml beaker

3. Transfer about 1 gram of baking soda to the beaker. Record the exact mass of the beaker with the powder

4. Add vinegar, from the pipette, to the beaker. Swirl the contents and observe the reaction. Continue to add vinegar until no more bubbles are produced. This will take a while so be patient and pay close attention to the reaction

5. Find the mass of the left-over vinegar in the pipette and record. Subtract the original mass of this pipette to find the mass of the vinegar used in the reaction

Initial Mass of Pipette in Grams	1.36 Grams
Final Mass of Pipette at End of Reaction	2.52 Grams
Net Mass of Vinegar Used in Reaction	9.81 Grams
Mass of Empty Beaker	113.73 Grams
Mass of Beaker and Baking Soda	114.73 Grams
Net Mass of Baking Soda	1 Gram

Conclusion Questions:

1.     Calculate the moles of sodium hydrogen carbonate used from the net mass of the baking soda.

1gNaHCO₃ × 1 mole / 81g NaHCO₃ = .0123 moles

2.     There are 5 grams of acetic acid per 100 grams of vinegar. Calculate the grams of acetic acid added to the beaker for the reaction.

5% of 9.81 = .49 grams of acetic acid

3.     Calculate the moles of acetic acid used in the reaction.

.49gHC₂H₃O₂ × 1 mole / 60g HC₂H₃O₂ = .0081 moles

4.     Compare the moles of sodium hydrogen carbonate with the moles of acetic acid used in this experiment. Find the nearest whole number ration between the two values.

Experimental Ratio: 2 to 3

5.     The ratio calculated above is your experiment mole ratio. Is this mole ration the same as the mole ratio in the balance equation?

No, due to rounding errors along with weighing errors resulted in the chance in the results.

Snow!

Chemistry of Snow