A student studies the kinetics of reaction 1 below using a clock reaction technique, reaction 2, and the procedures given in this experiment. For reaction mixture 1, he measures a time to color change of 121 seconds and for reaction mixture 2, he measures a time of 58 seconds. The temperatures of the reaction mixtures were the same for both experiments.
6Iâ (aq)+ BrO3â (aq) + 6H+(aq) --> 3I2 (aq) + Brâ (aq) + 3H2O(l) (Reaction1,slow)
I2 (aq) + 2 S2O32â (aq) --> 2 Iâ (aq) + S4O62â (aq) (Reaction 2, fast)
a. Determine the change in concentration of S2O32â that occurred during the time period measured. This should be the same for both the reaction mixtures. First, use M1V1=M2V2 to determine the initial concentration of the S2O32â after mixing the contents of the two flasks but before the reaction begins. Next, find the final concentration of S2O32â by remembering that all the S2O32â is consumed when the color change occurs. You can then find the change in the concentration of S2O32â.
[S2O32â]0 =
â[S2O32â] =
b. Use the stoichiometry of Reactions 1 and 2 to determine the change in concentration of BrO3â during
the same time period. â[BrO3â] =
c. Calculate the initial rate for each reaction mixture. Recall that Rate = â â[BrOâ ] / ât
Rate for mixture 1:
Rate for mixture 2:
d. Method of Initial Rates
Complete the table below using M1V1=M2V2 to find the initial concentration of each reactant after mixing the contents of flasks I and II for both reaction mixtures. Enter the initial rates computed above.
Rxn Mixture [Iâ]0(M) [BrO3â]0(M) [H+]0 (M) Rate (M/s)
1
2
Use the method of initial rates and the relevant data above to determine the order of reaction 1 with respect to the concentration of iodide ions.
Order with respect to [Iâ]:_________________________
e. It is critical for the success of this experiment that the concentrations of the reactants Iâ, BrO3â, and H+ do not change significantly over the time period measured.
Why is this important?
f. By what percent did the bromate ion concentration change during the time period considered for mixture 1 in this experiment?
PercentChange = (|â[BrO3â ]| / [BrO3â ]0 ) x 100%
A student studies the kinetics of reaction 1 below using a clock reaction technique, reaction 2, and the procedures given in this experiment. For reaction mixture 1, he measures a time to color change of 121 seconds and for reaction mixture 2, he measures a time of 58 seconds. The temperatures of the reaction mixtures were the same for both experiments.
6Iâ (aq)+ BrO3â (aq) + 6H+(aq) --> 3I2 (aq) + Brâ (aq) + 3H2O(l) (Reaction1,slow)
I2 (aq) + 2 S2O32â (aq) --> 2 Iâ (aq) + S4O62â (aq) (Reaction 2, fast)
a. Determine the change in concentration of S2O32â that occurred during the time period measured. This should be the same for both the reaction mixtures. First, use M1V1=M2V2 to determine the initial concentration of the S2O32â after mixing the contents of the two flasks but before the reaction begins. Next, find the final concentration of S2O32â by remembering that all the S2O32â is consumed when the color change occurs. You can then find the change in the concentration of S2O32â.
[S2O32â]0 =
â[S2O32â] =
b. Use the stoichiometry of Reactions 1 and 2 to determine the change in concentration of BrO3â during
the same time period. â[BrO3â] =
c. Calculate the initial rate for each reaction mixture. Recall that Rate = â â[BrOâ ] / ât
Rate for mixture 1:
Rate for mixture 2:
d. Method of Initial Rates
Complete the table below using M1V1=M2V2 to find the initial concentration of each reactant after mixing the contents of flasks I and II for both reaction mixtures. Enter the initial rates computed above.
Rxn Mixture [Iâ]0(M) [BrO3â]0(M) [H+]0 (M) Rate (M/s)
1
2
Use the method of initial rates and the relevant data above to determine the order of reaction 1 with respect to the concentration of iodide ions.
Order with respect to [Iâ]:_________________________
e. It is critical for the success of this experiment that the concentrations of the reactants Iâ, BrO3â, and H+ do not change significantly over the time period measured.
Why is this important?
f. By what percent did the bromate ion concentration change during the time period considered for mixture 1 in this experiment?
PercentChange = (|â[BrO3â ]| / [BrO3â ]0 ) x 100%
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Related questions
Everytime I make the graph based off my calculations in the chart it looks like this and I have no idea if it is even close? Please help ! For the questions I just need to understand how to start them and what info is nessiary to do so. very confused. 
Data Analysis
Fill in the following table. Look below the table for instructions on how to calculate the values for each row of the table.
| a | total volume (L) | 0.065L |
| b | mol S2O32-consumed | 0.010 mol |
| c | mol I3-produced | 0.015M |
| d | [S2O32-] consumed (M) | 0.154M |
| e | [I3-] produced (M) | 0.231M |
Calculate the final volume of the reaction mixture after the contents of beaker B are added to beaker A. Report your answer in liters.
All S2O32- is consumed at the end of the reaction. Therefore, the moles of S2O32- consumed can be calculated using the equation below.
(moles S2O32-)consumed = Mstock x (Vstock)
The I3- produced in reaction 1 reacts with S2O32- in reaction 2 as shown.
I3- (aq) 2S2O32- (aq) â 3I- (aq) S4O62- (aq)
Therefore, for every mole of I3- produced, 2 moles of S2O32- have reacted.
(moles I3-)produced = (moles S2O32-)consumed / 2
[S2O32-]consumed = (moles S2O32-)consumed / (Vtotal)
Calculation is the same as d, but using I3- instead of S2O32-.
Fill in the following table. Look below the table for instructions on how to calculate the values for each column of the table.
| a | b | c | d | e | f | g | h | i | |
|---|---|---|---|---|---|---|---|---|---|
| Trial | Ît | Î[I3-] / Ît (M/s) | log(Î[I3-] / Ît) | V 0.2 M KI added (mL) | [Iâ]0(M) | log[Iâ]0 | V of 0.2 M (NH4)S2O8 (mL) | [S2O8]0 (M) | log[S2O8]0 |
| 1 | 19 | .011 | -1.94 | 25 | .077 | -1.35 | 25 | .077 | -1.11 |
| 2 | 40 | .0055 | -2.24 | 25 | .077 | -1.11 | 12.5 | .038 | -1.42 |
| 3 | 85 | .0030 | -1.84 | 25 | .077 | -.1.11 | 6.25 | .019 | -1.72 |
| 4 | 42 | .0055 | -2.26 | 12.5 | .038 | -1.42 | 25 | .077 | -1.11 |
Conclusions
If a large amount of heat was released at the start of the reaction, what effect would this have on the rate measurements?
| The rate of reaction would be much slower due to its exothermic reaction properties. |
Click the box underneath the graph to show the trendline. It will automatically calculate your slope and intercept. Record them below.
| Determine the value of q from the graph. Explain your answer. |
What is the rate law for the reaction?
For each trial, calculate the rate constant. What is average value of the rate constant?
In this experiment, you assumed that [S2O82â] >> [S2O32â]. To find out if this assumption is correct, calculate the ratio of [S2O82â] / [S2O32â] for Trial 3. (In Trial 3, the [S2O82â] was lowest, and therefore the ratio [S2O82â] / [S2O32â] is the smallest).
In this experiment, you assumed that only a small amount of S2O82â was used during the time trial so that the concentration of this reactant did not change appreciably during the course of the reaction. Calculate how much S2O82â was used in Trial 3 and the percent remaining at the end of the reaction. Was this assumption valid?
