Rate Equation

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Objective
1. To determine how the concentration of a species can affect reaction rate in the determination of rate law and rate constant. 2. To determine how temperature affects reaction rate.
Introduction
Chemical kinetics deals with the speed, or rate, of a reaction and the mechanism by which the reaction occurs. We can think of the rate as the number of events per unit time. The rate at which you drive (your speed) is the number of miles you drive in an hour (mi/hr). For a chemical reaction the rate is the number of moles that react in a second. In practice, we usually monitor how much the concentration (the number of moles in a liter) changes in a second. Reaction rates are usually expressed in units of moles per liter per second, or molarity per second (M/s). In this experiment you will investigate how reactant concentration and temperature influence the rate of the following oxidation–reduction reaction between peroxydisulfate ion and iodide ion: S2O82 + I– —> SO4–2 + I2

The dependence of rate on concentration is given by the rate law: r = k[S2O8–2]m[I–]n
The rate of a reaction is the change in concentration of a reactant with time. In this experiment we will monitor the peroxydisulfate concentration, so the rate will be expressed as: r = –∆ [S2O8–2]/∆ t (the minus sign insures that the rate is a positive quantity)

The rate constant k is a constant for a given reaction and varies only with temperature. The Arrhenius equation: k = Ae–Ea/RT
This will be accomplished by using a second reaction, in which thiosulfate ions consume I2 molecules as soon as they are produced: I2 + S2O3–2 —> I– + S4O6–2
A starch indicator is also added to the solution. The starch indicator turns blue–black in the presence of I2. When the reagents are mixed, I2 is produced in the first reaction, which is then immediately consumed in the second reaction. As soon as all of the thiosulfate (in reaction 2 ) has been consumed, I2 molecules accumulate, react with the starch indicator, and the solution turns blue–black. By measuring the time it takes for the solution to turn color as a function of reactant concentration, the rate law for the reaction can be determined. Apparatus

Stopwatch, 250 mL conical flask, burette, pipette.
Chemicals
0.20 M potassium iodide (KI solution), 0.0050 M sodium thiosulfate (Na2S2O3 containing 0.4% starch indicator), 0.10 M potassium peroxydisulfate (K2S2O8) . Procedure
Part 1 : Finding the Rate Law and Rate Constant
1. The most accurate piece of glassware that allows for variable volumes to be dispensed for all volume additions was used. 2. The solution for each trial based on the volume as stated in table 3.1 was prepared. 3. Timing was beginning as soon as the peroxydisulfate is added and was immediately swirl the reaction to ensure good mixing. The timing was stopped when the solution changes to blue. 4. Each trial was performed at least two trial or until reproducible reaction times are obtained. The concentration of KI and K2S2O8 are such that a stoichiometric mixture for reaction was result from equal volumes. 5. To ensure uniform conditions for all runs, a constant total volume of 50.0 mL in a 250 mL conical flask was used and was maintained initial [S2O32-]0 = 0.0010 M for each run.

Volume used for the preparation each of trial
Trial | Volume I-| Volume S2O32- (mL)| Volume H2O (mL)| Volume S2O8(mL)| Total volume (mL)| 1| 20| 10| 0| 20| 50|
2| 15| 10| 5| 20| 50|
3| 10| 10| 10| 20| 50|
4| 5| 10| 15| 20| 50|
5| 20| 10| 5| 15| 50|
6| 20| 10| 10| 10| 50|
7| 20| 10| 15| 5| 50|

Part 2 : Temperature dependence of the rate rate constant
1. The mixture containing I—S2O32- was warmed or heated using hot plate or cool it using an ice bath. 2. The timing was beginning as soon as peroxydisulphate (S2O82-) is added and the reaction was immediately swirled to ensure good mixing. The temperature was...
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