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Pre-Lab Questions 1

Pre-Lab Questions

  1. How many grams of IKI would it take to obtain a 100 mL solution of 0.300 M IKI?  How many grams of IKI would it take to create a 100 mL solution of 0.600 M IKI?

8.79 grams

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17.57 grams

  • What is the molarity of 100 mL of a 3.0% H2O2 (mass/volume) solution?  100 mL of a 2.25% H2O2 solution?



Experiment 1: Calculating Rate of Reaction

In this experiment you will calculate the rate of reaction of potassium iodide and hydrogen peroxide. The order of the reaction will also be determined.

Materials: 20 mL 3% Hydrogen peroxide, H2O2
15 mL Iodine-Potassium Iodide Solution, IKI (1% Iodine, 2% Potassium Iodide)
100 mL Beaker
10 mL Graduated cylinder
100 mL Graduated cylinder
250 mL Beaker
(2) 250 mL Erlenmeyer flasks
1-Hole Rubber stopper
2-Hole Rubber stopper
(2) 3 in. Rigid tubing pieces
(1) 6 in. Rigid tubing
6 Pipettes
  Waste beaker (any volume)*Stopwatch
*Permanent marker
*Access to graphing software
*Access to a computer
*10 mL Distilled water *24 in. Flexible tubing * You must provide *You must cut this piece of tubing into two, 12 inch pieces. (if you have not already done so).


Preparation of Apparatus

  1. Set up apparatus as shown in Figure 2. To do this, begin by labeling the Erlenmeyer Flasks as 1 and 2. The reaction will take place in Flask 1.
  2. Fill Flask 2 approximately three quarters of the way full with water.
  3. Press the 2-hole rubber stopper into the top of Flask 2. Place one three in. piece and one six in. piece of rigid tubing into each hole of the rubber stopper. This should create an airtight system.
  4. Place the one-hole stopper on Flask 1, and fit the remaining 3 in. piece of rigid tubing in the stopper hole.
  5. Connect Flask 1 and Flask 2 with the two, 12 in flexible tubing pieces. One piece should connect Flask 1 to Flask 2, and the second piece should connect Flask 2 to the graduated cylinder. The tubing which connects Flask 2 to the graduated cylinder should be positioned low enough to be immersed in the water in Flask 2.
Figure 2: Apparatus set-up. Note this is a sample set-up and is not drawn to scale. Your specific equipment may vary slightly.

Part A: Preparation of Reactants

  1. Pour five mL of the IKI solution into a 10 mL graduated cylinder.
  2. Add five mL of distilled water to the graduated cylinder to bring the total volume to 10 mL. This is the 0.5% – 1.0% (diluted) IKI solution.
  3. Pour 15 mL of 3% H2O2 solution into a 100 mL beaker.
  4. Add five mL of distilled water to this beaker and mix with a stir rod. This is the 2.25% (diluted) H2O2 solution.

Part B: Performing the Reaction

  1. Remove the stopper from Flask 1 and place five mL of the 3% (undiluted) H2O2 solution and 10 mL of the undiluted IKI solution provided into the flask. Immediately replace the stopper on the flask.
    Note: At this point, you should select an extra beaker (any volume) from your lab kit to use as a supplemental collection container beaker for Step 6. You do not need to use the beaker yet, but keep it in close proximity.
  2. Swirl Flask 1 until you observe a steady dripping of water going into the 10 mL graduated cylinder. This could take 3 – 5 minutes. Check for leaks in the tubing or system if water does not start rising up the plastic tubing coming from Flask 2 and traveling towards the graduated cylinder within one minute.  
  3. Stop swirling Flask 1 when you notice the steady flow of water droplets. When you stop, the water drop -rate will significantly decrease (to around one drop every 5 – 20 seconds) and could take a few minutes to stabilize. If a steady flow of drops of water does not occur within a few minutes, swirl Flask 1 for 1 more minute and check again. Repeat this process until there is a steady flow of drops of water after you have stopped swirling Flask 1.
  4. Quickly empty liquid that has collected in the 10 mL graduated cylinder and replace the empty cylinder back under the flexible tubing.
  5. Allow the flow of drops to become steady again. This could take 1 – 3 mL of water.
  6. Start timing once the drop rate is steady and the volume of water collected is at a whole number (such as three mL). Record the time in Table 1 each time 2 mL of is water displaced. Continue taking data until you have at least 10 data points (20 mL displaced).
    Note: Use the extra beaker (located in Part B: Step 1) to collect additional fluid when the volume of displaced water exceeds 10 mL.
  7. Return the collected water from your 10 mL graduated cylinder to Flask 2. Ensure the seal is air tight.
  8. Empty, clean and dry Flask 1 and the graduated cylinder.
  9. Repeat Steps 1 – 8 for the following trial conditions: 5 mL 3% (undiluted) H2O2 mixed with 10 mL of 0.5%-1.0% IKI solution (placed in Flask 1); and, 5 mL of 2.25% H2O2 mixed with 10 mL of 1.0%-2.0% IKI solution (placed in Flask 1). Record the data in Table 2 and Table 3, respectively.
    Note: Clean the graduated cylinder and extra collection beaker before it is used to measure any additional reagents for Trial 2 or Trial 3; and, before it is used for collecting the water from the reaction in the apparatus.
  10. Use a graphing software program to make a graph of each trial. The graph should demonstrate the relationship formed between time vs. mL of water displaced.
  11. Find and record the slope and the inverse slope for each trial.
Table 1: 10 mL Undiluted (1.0 -2.0%) IKI and 5 mL 3% H2O2
mL water displaced Time (seconds)
2 72 
4 115 
6  164
8  206
10  244
12  283
14  323
16  368
18 408 
20  448
22  490
24  530
Slope:  0.0484
Inverse Slope: 20.66
Table 2: 10 mL Diluted (0.5-1.0% IKI) and 5 mL 3% H2O2
mL water displaced Time (seconds)
2 112 
4  178
6  262
8  335
10  398
12  465
14  536
16  612
18  687
20  768
22  855
24  930
Slope:  0.0271
Inverse Slope: 36.90
Table 3: 10 mL Undiluted (1.0 -2.0%) IKI and 5 mL 2.25% H2O2
mL water displaced Time (seconds)
2 51 
4  93
6  127
8  156
10  184
12  214
14  241
16  271
18  296
20  322
22  347
24  374
Slope:  0.0696
Inverse Slope:  14.37


Post-Lab Questions

  1. Determine the order of the KI in this reaction.


  • Determine the order of the H2O2 in this reaction.


  • Calculate the rate law constant.
  • What is the overall rate law?


  • When finding the order of H2O2, why was Trial 1 and Trial 3 used?

Because they represent different concentration levels for H2O2

  • When finding the order of KI, why was Trial 1 and Trial 2 used?

Because those trials represent different concentration levels for the KI

  • Research and identify at least two catalysts that could be used to accelerate the     decomposition of hydrogen peroxide. Evaluate these catalysts and determine which option is “greener”. Provide approximately ½ page summary of your research.

     From what I was able to come up with, it seems like the best choices were the enzyme catalase, and yeast, with yeast being the superior choice.  The reason being is that yeast is a naturally-occurring, and faster-acting decomposer of the hydrogen peroxide.  Below, I have provided an example of a decomposition reaction for hydrogen peroxide that should highlight how the introduction of a catalyst can speed up the decomposition process, and the role it plays in the reaction overall:  “This demonstration is based on the decomposition of hydrogen peroxide into water and oxygen gas.  Reactions like these that are both oxidations and reductions are known as disproportionation reactions:

     2 H2O2(aq) -> 2 H2O(l) + O2(g)

     Left on its own at room temperature, this reaction happens at a rate so slow that, for practical purposes, it may as well not even exist.  A catalyst is added to speed things along.   The KI added dissociates into K+ and I, at which point the I begins its work.  The reaction pathway represented below has a lower activation energy than the straight decomposition represented above:

     H2O2(aq) + I(aq) -> OI(aq) + H2O(l)

     H2O2(aq) + OI(aq) -> H2O(l) + O2(g) + I(aq)

     Note that the iodide ion is conserved in this reaction (it is not consumed in the sum of the reactions, and the same iodide ion could potentially go through many such cycles).  This is the definition of a catalyst: a substance that increases the rate of a chemical reaction without being consumed in the reaction.”

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