Yezi Cho
1. Purpose
The following reaction between iodine ions and iron (III) ions in an aqueous solution was conducted experimentally to determine its rate expression.
As the reaction progresses, the solution becomes a yellow-orange color, which can be measured quantitatively using a colorimeter. By systematically changing the initial concentration of iodine ions and iron (III) ions, the rate order for each of the reactants can be determined. The effect of the initial concentrations on the rate of reaction was recorded for 120 seconds through LoggerPro.
2. Materials and Methods
5cm3 pipette
10cm3 pipette
Colorimeter and laptop with Logger Pro
Cuvettes
0.020 mol dm–3 KI
0.020 mol dm–3 FeCl3
Distilled water
Safety goggles
1) Set the wavelength of the colorimeter to 470 nm and calibrate the colorimeter using 1.5 cm3 of distilled water and 1.5 cm3 of KI solution.
2) Create the various concentration manipulations of iodine ion and iron (III) ion solutions by dilution.
3) Add specified amounts of iodine and iron (III) solutions and distilled water into the cuvette.
4) Place the cuvette in the colorimeter and start collecting data every 15 seconds for a total of 120 seconds.
5) Repeat the process above for the four remaining concentrations.
Two trials were collected for each type of solution.
3. Results
Figure1. Absorbances over time of varying concentrations of KI and FeCl3Â to determine the rate expression of the reaction between potassium iodide and iron (III) chloride (%)
Absorbance (%)Â | |||||||||
Time (s) → | 15 | 30 | 45 | 60 | 75 | 90 | 105 | 120 | |
1: 1.60 mL KI, 0.80 mL H2O, 0.80 mL FeCl3 | Trial 1 | 0.163 | 0.185 | 0.201 | 0.220 | 0.228 | 0.240 | 0.250 | 0.253 |
     | Trial 2 | 0.121 | 0.149 | 0.172 | 0.194 | 0.211 | 0.226 | 0.241 | 0.255 |
Ave. absorbance | 0.142 | 0.167 | 0.187 | 0.207 | 0.220 | 0.233 | 0.246 | 0.254 | |
2:Â 1.60 mLÂ KI, 1.20 mL H2O, 0.40 mL FeCl3 | Trial 1 | 0.001 | 0.001 | 0.026 | 0.034 | 0.038 | 0.047 | 0.053 | 0.059 |
          | Trial 2 | 0.001 | 0.001 | 0.029 | 0.038 | 0.049 | 0.057 | 0.066 | 0.072 |
Ave. absorbance | 0.001 | 0.001 | 0.028 | 0.036 | 0.044 | 0.052 | 0.060 | 0.066 | |
3:Â 0.80 mL KI, 0.80 mL H2O, 1.60 mL FeCl3 | Trial 1 | 0.092 | 0.120 | 0.143 | 0.165 | 0.184 | 0.203 | 0.219 | 0.235 |
     | Trial 2 | 0.109 | 0.132 | 0.154 | 0.170 | 0.189 | 0.204 | 0.219 | 0.232 |
Ave. absorbance | 0.100 | 0.126 | 0.149 | 0.168 | 0.187 | 0.204 | 0.219 | 0.234 | |
4:Â 0.80 mL KI, 1.60 mL H2O, 0.80 mL FeCl3 | Trial 1 | 0.032 | 0.048 | 0.063 | 0.076 | 0.087 | 0.100 | 0.110 | 0.119 |
     | Trial 2 | 0.027 | 0.042 | 0.054 | 0.064 | 0.075 | 0.085 | 0.093 | 0.101 |
Ave. absorbance | 0.030 | 0.045 | 0.059 | 0.070 | 0.081 | 0.093 | 0.097 | 0.110 | |
5:Â 1.60 mL KI, 0.00 mL H2O, 1.60 mL FeCl3 | Trial 1 | 0.252 | 0.303 | 0.337 | 0.374 | 0.406 | 0.440 | 0.468 | 0.495 |
     | Trial 2 | 0.265 | 0.324 | 0.370 | 0.409 | 0.443 | 0.473 | 0.498 | 0.523 |
Ave. absorbance | 0.259 | 0.314 | 0.354 | 0.392 | 0.425 | 0.457 | 0.483 | 0.509 |
The average absorbances over time were calculated using the two trials.
Figure 2. Average absorbance (%) for 120 seconds for solution 1, solution 2, solution 3, solution 4, solution 5
Average Absorbance (%) | |||||
Time | Solution 1 | Solution 2 | Solution 3 | Solution 4 | Solution 5 |
15 | 0.142 | 0.001 | 0.1 | 0.03 | 0.259 |
30 | 0.167 | 0.001 | 0.126 | 0.045 | 0.314 |
45 | 0.187 | 0.028 | 0.149 | 0.059 | 0.354 |
60 | 0.207 | 0.036 | 0.168 | 0.07 | 0.392 |
75 | 0.22 | 0.044 | 0.187 | 0.081 | 0.425 |
90 | 0.233 | 0.052 | 0.204 | 0.093 | 0.457 |
105 | 0.246 | 0.06 | 0.219 | 0.097 | 0.483 |
120 | 0.254 | 0.066 | 0.234 | 0.11 | 0.509 |
Figure 3. Average absorbance for solution 1, solution 2, solution 3, solution 4, solution 5 over time
Figure 4. Slope of graph of absorbance over time
1 | 2 | 3 | 4 | 5 | |
[KI] | 0.01 | 0.01 | 0.005 | 0.005 | 0.01 |
[FeCl3] | 0.005 | 0.0025 | 0.01 | 0.005 | 0.01 |
Slope | 0.00106 | 0.000659 | 0.00126 | 0.000740 | 0.00233 |
The concentration of KI and FeCl3 was calculated using the total volume including distilled water. The reaction rate over 120 seconds was determined using the slope of each graph of absorbance (%) over time.
4. Conclusion
Using the slopes as an indication of rate of reaction, the manipulations above show that the orders with respect to KI and FeCl3 are both closest to 1. Thus the rate expression would be Rate = k[KI][FeCl3].
It should be noted that the specific rate constant k cannot be calculated from the data above, as the data recorded was not the concentration of ions but the absorbance.
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