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Objectives: To determine reaction orders with respect to reactants and formulate a rate law.
Background
The rate at which a chemical reaction occurs depends on the following factors: the nature of the reactants, the concentrations of the reactants, the temperature, and the presence of possible catalysts. For a given reaction, the rate typically increases with an increase in the concentration of any reactant. The rate law gives the relation between rate and concentration of reactants. For a reaction of the following form,
aA + bB ® cC (1)
The rate is expressed by Eqn. 2,
Rate = k [A]m [B]n (2)
where m and n are generally, but not always, integers 0, 1,2, or possibly 3; [A] and [B] are the concentrations of A and B, respectively. The constant k is called the rate constant of the reaction. The rate constant does not depend on concentrations of reactants, but depend on temperature.
In this experiment the kinetics of the reaction in Eqn. 3 will be studied. This reaction is called a “clock” reaction because of the means of monitoring of the reaction.
6 I–(aq) + BrO3–(aq) + 6 H+(aq) 3 I2(aq) + Br–(aq) + 3 H2O(l) (3)
This reaction is somewhat slow at room temperature. Its rate depends on the concentration of the reactants and on the temperature. We can express the rate of the reaction in terms of the rate of disappearance of the iodide ion, as in Eqn. 4.
The exponents m, n, and p are called the “orders” of the reaction with respect to the indicated reactants, and show how a change in the concentration of each reactant affects the rate of reaction. The values of m, n, and p will be determined in this experiment.
To determine the rate of a reaction we need some way of monitoring the reaction progress, i.e. measuring change in a reactant or product concentration as a function of time. For the reaction in Eqn. 3, the rate at which iodine forms can be measured indirectly by reacting it with S2O32-. The I2/S2O32- reaction, Eqn. 5, is much faster than the reaction in Eqn. 3, which is a precondition for this method reaction monitoring.
I2(aq) + 2 S2O32-(aq) 2 I–(aq) + S4O62- (aq) (5)
The clock reaction in Eqn. 5 uses much less than stoichiometric amount of thiosulfate ions. When all S2O32- ions are used up, iodine begins to build up in the solution. In the presence of starch, iodine forms a deep blue‑colored complex that is readily apparent. Since this is a very sensitive test, the blue color appears almost immediately after the thiosulfate ions are all consumed.
The same amount of thiosulfate ion will be used in all reaction mixtures (kinetics runs). One can see from the clock reaction, Eqn. 5, that two moles of S2O32- react with one mole of I2. Knowing the initial concentration of thiosulfate employed allows the determination of [I2] consumed in each kinetic run, which in turn allows the determination of the decrease in [I–], D[I–], at the instant of time the blue color appears. Measurement of the period of time, Dt, allows the calculation of the reaction rate according to Eqn. 6.
The experiment is designed in such a way that D[I–] is the same in all kinetics runs (D[I–] = [S2O32-]o). However, the time period (Dt) will be different since different initial concentrations of reactants will be used.
APPARATUS AND CHEMICALS
Beakers (100-mL), pipets (10-mL), stop-watch, 50-mL burets, 0.01 M KI, 0.0010 M Na2S2O3, 0.040 M KBrO3, 0.10 M HCl, 1% starch solution.
PROCEDURE
Table 1 shows the volumes of the different reagents that need to be mixed together in a given kinetic run.
Table 1. Reaction mixtures at room temperature
| Rxn. No. | Reaction Beaker I (100 mL) | Reaction beaker II (100 mL) | |||
| 0.01 M KI (mL) | 0.0010 M Na2S2O3 (mL) | H2O (mL) | 0.04 M KBrO3 (mL) | 0.10 M HCl (mL) | |
| 1 | 10 | 10 | 10 | 10 | 10 |
| 2 | 20 | 10 | 0 | 10 | 10 |
| 3 | 10 | 10 | 0 | 20 | 10 |
| 4 | 10 | 10 | 0 | 10 | 20 |
| 5 | 8 | 10 | 12 | 5 | 15 |
The procedure for carrying out Reaction #1 is described below. Reactions 2-4 can be carried out following a similar procedure.
- Using 10-mL pipets, measure out 10 mL 0.010 M KI, 10 mL 0.0010 M Na2S2O3, and 10 mL distilled water into a 100-mL beaker (beaker I).
- Measure out 10 mL 0.040 M KBrO3 and 10 mL 0.10 M HCl into another 100-mL beaker (beaker II).
- To beaker II, add three or four drops of starch indicator solution.
- While a student is noting time of mixing, pour the contents of beaker II into beaker I and pour the mixture back and forth a couple of times. Stir the mixture carefully with a clean glass rod. The solution should turn blue in less than 2 minutes.
- Note the time at the instant the blue color appears. Also measure the temperature of the solution.
- Repeat steps 1-5 with the remaining reaction mixtures.
- Repeat any experiments that did not appear to proceed properly.
DISPOSAL
The spent reaction mixtures can be poured into the sink, since very dilute solutions were used.
After you have read the procedure, view a video of:
Bromate – Iodide Clock Reaction Lab
North Carolina School of Science and Mathematics, •Dec 15, 2011;
(Please note that in this video, temperature is varied instead of concentration, so do not use the results as data. I have supplied the time data instead).
Copy the following video link into your web browser and play the video:
Experiment 7: Kinetics: Determination of Overall Reaction Order
Name: ________________________ Section: ______________
Partner(s): _______________________________ Date: ________________
I. DATA AND CALCULATIONS
Table 1. Initial concentrations of reactants, time (s), and temperature (oC) data
| Rxn. No. | Time(s) (Dt) | Initial concentrations of reactants (mol/L) | Temp. (oC) | Rate (M/s) | ||
| [I–]o | [BrO3–]o | [H+]o | ||||
| 1 | 62 | 0.0020 | ||||
| 2 | 15 | |||||
| 3 | 17 | |||||
| 4 | 32 | |||||
| 5 | 258 |
Calculate the initial concentrations of the reactants in each reaction mixture. The volume of the reaction mixtures is 50 mL in all cases. The initial concentration of I– in Reaction (1) is calculated as follows as an example:
[I–]stock . Vstock = [I–]mixture . Vmixture
=> [I–]mixture = (0.010 M x 10 mL) / 50 mL = 0.0020 M
II. Reaction Order with respect to each reactant
The individual reaction orders can be determined from the ratios of rates of two runs. For instance, according to Eqn. 6 the rates of Reactions (1) and (2) are:
Rate1 = k [I–]m [BrO3–]n [H+]p and Rate2 = k [2I–]m [BrO3–]n [H+]p
=>
The value of “m’ is determined using the periods of time, Dt, measured for the two runs.
a) Calculate the value of m in the space given below. m = _______ (nearest integer)
b) Use Rate1 and Rate3 to determine the value of n. n = __________ (nearest integer)
c) Use Rate1 and Rate4 to determine the value of p. p = _________ (nearest integer)
d) Calculate the value of k for each kinetic run using the following equation. Use the values of m, n, and p you determined experimentally.
e) Using the average value of k, predict the initial rate of reaction mixture 5. Use the initial concentrations from Table 1.
f) Calculate the initial rate of disappearance of iodide ions in reaction mixture 5.
