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The Iodide - Persulphate Reaction: Determining the Effect of Concentration on Reaction Rate

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The Iodide - Persulphate Reaction: Determining the Effect of Concentration on Reaction Rate Name: April Yue Date: February 10, 2004 Student ID: 20131652 Section: 006 T.A. : Mary Diep April Yue (20131652) CHEM 123L Feb. 10, 2004 Experiment 3: The Iodide - Persulphate Reaction: The Effect of Concentration on Reaction Rate Introduction: In this experiment, we utilized the ability for the iodide ion to become oxidized by the persulphate ion. Our general reaction can be described as: (NH4)2S2O8 + 2KI --> I2 + (NH4)2SO4 + K2SO4 (1a) However, we know that in an aqueous solution, all of these compounds except iodine will dissociate into their ionic components. Thus we can rewrite the equation in a more convenient manner: S2O82- + 2I- --> I2 + 2SO42- (1b) It is important however to note that the NH4 and K ions are still in the solution, they are just unreactive. In order to measure the rate of the reaction, the conventional method would be to measure the species in question at certain times. However, this would be inconvenient, especially for a three hour laboratory period. Since the iodide ion can be oxidized by the persulphate ion, we can use sodium thiosulphate to be an indicator of the presence of iodine in the solution. For this experiment, we can simply calculate the rate of the reaction by timing the amount of iodine being produced in several runs. ...read more.


log [S2O8-2] m= rise = -2.678 - (-2.055) = -0.623 = 1 run -2.041 - (-1.439) -0.602 slope 2: -log ?t vs. log [I-] n= rise = -2.635 - (-2.055) = -0.58 = 1 run -1.74 - (-1.138) -0.602 3. Sample Calculations: Rate of Reaction for run #1: Rate = -?S2O8-2 / ?t = -9.09 x 10-4 M / 113.5 s = 8.01 x 10-6 M s-1 Rate constant (k) for run #1: k = Rate / ( [S2O8-2]m [I-]n ) k = 8.01 x 10-6 M s-1 / {(3.64 x 10-2 M)1(7.27 x 10-2 M)1} k = 3.03 x 10-3 s-1 Ionic Strength (�) for run #1: � = 0.5 ? CiZi2 � = 0.5{([NH4]x(+1)2) + ([S2O8]x(-2)2) + ([K]x(+1)2) + ([I]x(-1)2) + ([Na]x(+1)2) + ([S2O3]x(-2)2)} � = 0.187 mol L-1 Table 2 - Calculations Summary Table Run # [S2O8-2] (M) [I-] (M) [S2O3-2] (M) -?S2O8-2 (M) ?t (s) Rate (M s-1) Rate Constant, k (s-1) Ionic Strength (M) 1 3.64 x 10-2 7.27 x 10-2 1.82 x 10-3 -9.09 x 10-4 114 -8.01 x 10-6 3.03 x 10-3 0.187 2 1.82 x 10-2 7.27 x 10-2 1.82 x 10-3 -9.09 x 10-4 218 -4.16 x 10-6 3.15 x 10-3 0.187 3 9.09 x 10-3 7.27 x 10-2 1.82 x 10-3 -9.09 x 10-4 476 -1.91 x 10-6 2.89 x 10-3 0.187 4 3.64 x 10-2 3.64 x 10-2 1.82 x 10-3 -9.09 x 10-4 228 -3.98 x 10-6 3.01 x 10-3 0.187 ...read more.


Since the ionic strength has also decreased, it has some effect on the resulting rate constant and therefore skews the results a bit. The rest of the results seem to agree with the logical way the experiment should have occurred. For example, the runs with the longer elapsed times had the slower reaction rates and vice versa with the runs with the shorter elapsed times. This makes sense due to the linear relationship between reaction rate and time. Some sources of error in this experiment may have been a mistake in mixing certain reactants, or inaccuracy with measuring volume of the solutions. It was more likely that there was inaccurate measuring of the solutions because it was quite difficult to always use the Mohr and transfer pipettes precisely. Conclusions: The purpose of this experiment was to determine how concentration of a certain reactant in a reaction can affect the rate of the entire reaction. The experiment was overall a success because we could see that when we varied the concentrations of certain compounds, the reaction rate was affected accordingly. Overall we know that the rate of the reaction is linearly proportional to the concentration of your reactant. However, if your reaction can exist in equilibrium and you increase the concentration of a product, the reaction will favour in the left direction, and if you are measuring rate of product formation, this will result in a decrease in reaction rate. Reference(s): Chemistry Department, First Year Chemistry: Chem 123L Laboratory Manual. University of Waterloo: 2004 ...read more.

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