Fecl3 And Ki Essay

Trial | [FeCl3] (M) | [KI] (M) | Initial rate (sec -1) | 1 | .01 | .01 | .002246 | 2 | .01 | .005 | .001348 | 3 | .005 | .01 | .001627 | 4 | .0075 | .005 | .001126 | 5 | .005 | .0075 | .001267 |

Order of Reaction: n=1;m=1
R1/R2=.002246/.001348=k[.01][.01]/k[.01][.005]
1.666172=2n n=1

Rate law Expression:
Rate=k[FeCl3][KI]

Rate Law constant:
Trial 1: .002246=k[.01][.01] k=22.46
Trial 2: k=26.96
Trial 3: k=32.54
Trial 4: k=30.027
Trail 5: k=33.787
Average Rate Constant: Trial1+2+3+4+5=145.774/5=29.1548

Procedure
Obtain and wear goggles. Connect a Colorimeter to Channel 1 of the Vernier computer interface. Connect the interface to the computer with the proper cable. Start the Logger Pro program
on your computer. Open the file “25 Rate and Order” from the Advanced Chemistry with Vernier folder.
Set up and calibrate the Colorimeter. Prepare a blank by filling an empty cuvette ¾ full with distilled water. Place the blank in the cuvette slot of the Colorimeter and close the lid. If your Colorimeter has a CAL button, set the wavelength on the Colorimeter to 430 nm, press the CAL button.
Obtain the materials you will need to conduct this experiment.
• Three 25 mL graduated cylinders.
• Approximately 100 mL of 0.020 M KI solution in a 100 mL beaker.
• Approximately 100 mL of 0.020 M FeCl3 solution in a separate 100 mL beaker.
• Approximately 60 mL of distilled water in a third 100 mL beaker.
During this experiment you will conduct 5 trials. This step describes the process for conducting the trials using the Trial 1 volumes. When you repeat this process, use the correct volume for each trial based on the table below.
Trial FeCl3 (mL) KI (mL) H2O (mL)
1 20.0 20.0 0.0
2 20.0 10.0 10.0
3 10.0 20.0 10.0
4 15.0 10.0 15.0
5 10.0 15.0 15.0
Measure 20.0 mL of FeCl3 solution into a 100 mL beaker.
Measure 20.0 mL of KI solution into a second 100 mL beaker. Remove the cuvette from the Colorimeter and pour out the distilled water. Add the 20.0 mL of FeCl3 solution to the beaker of KI solution. Swirl the beaker to mix. Rinse the cuvette twice with ~1-mL amounts and then fill it ¾ full. Wipe the outside of the cuvette with a tissue, place it in the Colorimeter, and close the lid.
Click to begin collecting absorbance data. Data will be gathered for 2 minutes. Observe the progress of the reaction in the beaker. When the data collection is complete, carefully remove the cuvette from the Colorimeter. Dispose of the contents of the beaker and cuvette as directed. Rinse and clean the beakers and the cuvette for the next trial. Examine the graph of the first trial. Select a linear region of the graph that covers about 30-40 seconds of the reaction. Click the Linear Regression button, . Record the slope, as the initial rate of the Trial 1 reaction, in your data table and then close the Linear Regression box.
10. Repeat Steps 6-9 to conduct Trials 2-5. Note: You will not perform Step 6c in Trials
2-5.

Propose
The purpose of this experiment is to determine the rate and order of a reaction.
To do this one uses Beers Law which states that the absorbance of a molecule is linearly proportional to the thickness of the sample, the concentration of the absorbing medium, and the absorption, which is a measure of a given molecule's ability of absorb light. To perform this experiment one will use a Colorimeter. A colorimeter is a light-sensitive instrument that measures how much color is absorbed by an object or substance. It determines color based on the red, blue, and green components of light absorbed by the object or sample. When light passes through a medium, part of the light is absorbed, and as a result, there is a decrease in how much of the light reflected by the medium. A colorimeter measures that change so one can analyze the concentration of a particular substance in that medium. A colorimeter works on the basis of Beers law, which says that the absorption of light transmitted through a medium is directly proportional to the concentration of the
medium.

In this experiment the reaction was conducted between solutions of potassium iodide and iron(III) chloride. The equation of the reaction is below, in ionic form. 2 I– (aq) + 2 Fe3+ (aq) → I2 (aq) + 2 Fe2+ (aq)

Conclusion
In conclusion the group ended up with the reaction being a 2nd order reaction. After doing the experiment the rate law for our experiment came out to be Rate=k[FeCl3][KI]. The lab asked if it was possible to calculate the rate constant k, and it was possible. The average of our rate constant k’s came out to be 29.1548. I am not sure if that is an exceptable rate constant since there is nothing to compare it to. If our results are not an exceptable rate constant, there may have been a few errors during the process of our experiment. Finger prints could have been on the clear side of the cuvette preventing the light from passing through. Another reason for error could have been that we let the solutions sit to long before pouring them into the cuvette for testing.

Chelsea Roy
AP Chemistry
p.6
Abstract

In this experiment one will determine the rate and the order of the reaction:
2 I– (aq) + 2 Fe3+ (aq) → I2 (aq) + 2 Fe2+ (aq)
To determine the rate and the order of the reaction, one must use a colorimeter and Beers Law. A colorimeter is a piece of equipment that lets you determine the amount of light absorbed through a solution. Beers Law states that absorbance is directly proportional to the concentration of the solution.After the experiment our group found the overall order of the reaction to be 2nd order. After some calculations we found the rate law to be Rate=k[FeCl3][KI].