The rate law of a chemical reaction is a mathematical equation that describes how the reaction rate depends upon the concentration of each reactant. Two methods are commonly used in the experimental determination of the rate law: the method of initial rates and the graphical method. In this experiment, we shall use the method of initial rates to determine the rate law of a reaction. You should review the sections on determining rate laws via this method in the chapter on chemical kinetics in your textbook before proceeding with this experiment. The reaction to be studied in this experiment is represented by the following balanced chemical equation:6I−(aq)+BrO−3(aq)+6H+(aq)⟶ 3I2(aq)+Br−(aq)+3H2O(l)(1.1)This reaction proceeds relatively slowly. The rate law for this reaction is of the form:Rate=k[I−]x[BrO−3]y[H+]z(1.2)w here the value of the rate constant, k , is dependent upon the temperature at which the reaction is run. The values of x , y , z and k must be found for this reaction in order to specify the rate law completely. The values of the reaction orders, x , y , and z , are usually, though not always, small integers. The method of initial rates allows the values of these reaction orders to be found by running the reaction multiple times under controlled conditions and measuring the rate of the reaction in each case. All variables are held constant from one run to the next, except for the concentration of one reactant. The order of that reactant concentration in the rate law can be determined by observing how the reaction rate varies as the concentration of that one reactant is varied. This method is repeated for each reactant until all the orders are determined. At that point, the rate law can be used to find the value of k for each trial. If the temperatures are the same for each trial, then the values of k should be the same too. The rate of the reaction can be defined as the rate of decrease of the concentration of BrO−3 :Rate of Reaction=−Δ[BrO−3]Δt(1.3)Note that this is actually the average rate of the reaction over the time interval Δt since the concentration of BrO−3 and therefore the rate of the reaction is continuously decreasing during Δt . In this experiment we can assume that Δ[BrO−3] is negligible compared to [BrO−3] thus allowing us to approximate our measured average rate as being equal to the initial, instantaneous rate. The rate of the reaction will be measured indirectly by running a second reaction, known as a clock reaction, simultaneously along with the reaction of interest. The clock reaction must be inherently fast relative to the reaction of interest and it must consume at least one of the products of the reaction of interest. In this case the parallel reactions are the following.6I−(aq)+BrO−3(aq)+6H+(aq) ⟶3I2(aq)+Br−(aq)+3H2O(l)slow(1.4)(c lock reaction)3I2(aq)+6S2O2−3(aq)⟶6I−(aq )+3S4O−2(aq)fast(1.5)Note that since the clock reaction is relatively fast, the I2(aq) will be consumed by the clock reaction as quickly as it can be produced by the reaction of interest thus holding the I2 concentration at a very low value close to zero. Only after the S2O2−3 has been entirely consumed will the concentration of I2 start to increase. After the S2O2−3 has been consumed the concentration of I2 will increase rapidly allowing the I2 to react with a starch indicator resulting in the solution turning to a deep blue color. In this experiment you will measure the time required for the solution to turn blue. This is essentially the amount of time required for all of the S2O2−3 to be reacted. Inspection of Table 1 shows that the number of moles of S2O2−3 is the same for all reaction mixtures and is relatively small compared to the number of moles of each reactant. Furthermore a stoichiometric calculation can be used to determine the number of moles of BrO−3 reacted at the point in time when all of the S2O2−3 has been consumed. By dividing the number of moles of BrO−3 reacted by the total volume of the reaction mixture we can obtain the change in the BrO−3 concentration. We can obtain the reaction rate by dividing this change in concentration by the amount of time required for the blue color to appear ( Δt ).