Rate Law Lab
Maya Parks
Partners: Ben Seufert, Brandon Busick
3/27/15 and 3/30/15
Abstract:
This lab was performed to determine the reaction rate equation of the creation of iodine. My
group found the equation to be
. From there we determined the value
of k at different temperatures and in the presence of a catalyst. We found larger k values at
higher temperatures and with the addition of a catalyst in the reaction. We found the correct
order of the reactants, however, there was a large deviation for the k values at 22.0°C. On several
occasions, we might not have had enough of the reactant or perhaps too much, and quite a few
times, we stopped the timer a little later than we should have. This caused error in the data, and
can be seen when calculating k. Our Arrhenius graph was not a straight line and had some
deviation. Regardless of the error in this lab however, it served its purpose in demonstrating the
effects of concentration, temperature, and a catalyst on the reaction rate. We were able to witness
a real-world application of concepts taught in class and gained invaluable experience in the
laboratory.
Purpose:
This lab was performed to demonstrate how the different variables affect the reaction rate, such
as temperature, concentration, and a catalyst. We learned the equation for reaction rate in class
and we will be finding this equation using this real world example. (the experiment) We will use
data collected to determine the value of k as well as the order of each reactant.
Materials:
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.010M KI
.001M Na2S2O3
Distilled water
.040M KBrO3
.10M HCl
.5M (NH4)2MoO4
125mL Erlenmeyer Flask
250mL Erlenmeyer Flask
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2 10mL graduated cylinders
2 50mL graduated cylinders
Pipet
Timer
Starch indicator
Thermometer
Hot water bath
Ice bath
Goggles
Procedure:
PART A. DEPENDENCE OF REACTION RATE ON CONCENTRATION
In Table 1 we have summarized the reagent volumes to be used in carrying out the
several reactions whose rates we need to know in order to find the general rate law for (EQ. 6).
First, measure out 100 mL of each of the listed reagents (except H2O) into clean, labeled
beakers. Use these reagents in your reaction mixtures.
Table 1
Reaction Mixtures at Room Temperature
(Reagent Volumes in mL)
1
2
3
4
5
Reaction Flask I (250 Flask)
.010M KI
.001M Na2S2O3
H2O
10
10
10
20
10
0
10
10
0
10
10
0
8
10
12
Reaction Flask II (125 flask)
.040M KBrO3
.10M HCl
10
10
10
10
20
10
10
20
5
15
The actual procedure for each reaction mixture will be much the same, and we will
describe it now for Reaction Mixture 1.
Since there are several reagents to mix, and since we don’t want the reaction to start until
we are ready, we will put some of the reagent into one flask and the rest into another, selecting
them so that no reaction occurs until the contents of the two flasks are mixed. Using the 10 mL
graduated cylinder to measure volumes, measure out 10mL of 0.010M KI, 10 mL of 0.001M
Na2S2O3, and 10 mL of distilled water into 250 mL Erlenmeyer flask (reaction Flask 1). Then
measure out 10 mL of 0.040M KBrO3 and 10 mL 0.10M HCl into a 125 mL Erlenmeyer flask
(Reaction Flask II). To Flask II add 3 mL of starch indicator solution.
Pour the contents of Reaction Flask II into Reaction Flask I and swirl the solution to mix
them thoroughly. Note the time at which the solutions are mixed. Continue swirling the
solution. Note the time in seconds from the time of mixing to the instant that the blue/black
color appears. Record the temperature of the blue solution to 0.5oC.
Repeat the procedure with the other mixtures in Table 1. Don’t forget to add the starch
indicator before mixing the solutions in the two flasks. The reaction flasks should be rinsed with
distilled water between runs. When measuring out reagents, rinse the graduated cylinder with
distilled water between measuring the reagents for Flask I and Flask II. Try to keep the
temperature just about the same in all the runs. Repeat any experiments that did not appear to
proceed properly. Check with your instructor.
PART B. DEPENDENCE OF REACTION RATE ON TEMPERATURE
In this part of the experiment, the reaction will be carried out at several different
temperatures, using Reaction Mixture 1 in all cases. The temperature we will use will be about
0oC , 20oC, 40oC and 60oC.
We will take the time at about 20oC to be that for Reaction Mixture 1 as determined in
Part A at room temperature. To determine the time at 40oC proceed as follows. Make up
Reaction Mixture 1 as you did in Part 1, including the indicator. However, instead of mixing the
solutions in the two flasks at room temperature, put the flasks into the 40oC hot water bath set up
in the lab. Check to see that the water is indeed at about a constant 40oC (record the actual
temperature), and leave the flasks in the water for several minutes to bring them to the proper
temperature. Then mix the two solutions, nothing the time of mixing. Continue swirling the
reaction flask in the warm water. When the color change occurs note the time and temperature
of the solution in the flask.
Repeat the experiment at about 60oC, heating all the reactant in a water bath at that
temperature (record the temperature) before starting the reaction. Record the time required for
the color to change and the final temperature of the reaction mixture. Repeat once again at about
0oC, this time using an ice bath.
PART C. DEPENDENCE OF REACTION RATE ON THE PRESENCE OF A CATALYST
Some ions have a pronounced catalytic effect on the rats of many reaction in water
solution. Observe the effect on this reaction by once again making up Reaction Mixture 1.
Before mixing, add 1 drop of 0.5M ammonium molybdate, (NH4)2MoO4, and a few drops of
starch indicator to Reaction Flask II. Swirl the flask to mix the catalyst thoroughly. Then mix
the solution, nothing the time required for the color to change.
Data:
Mixture
Time
1
4min 36.31s
2
2min 26.83s
3
2min 27.03s
4
1min 14.65s
5
5min 56.16s
*refer to table in procedure
Temperature
1.0°C
22.0°C
45.3°C
66.2°C
Time
9min 23.75s
4min 36.31s
40.73s
14.39s
Mixture 1 with catalyst
Time
1 (+2 drops (NH4)2MoO4)
10.78s
*compare with mixture 1 time in first table
Calculations:
PART A CALCULATIONS
Reaction
mixture
Time t (sec)
for color to
change
Relative rate
of reaction,
1000/t
Reactant concentrations in
reacting mixture [M]
[I-]
[BrO3-]
[H+]
Temperature
o
C
K
1
276.31
3.6191
.0020
.0080
.020
22.0
295.15
2
3
4
146.83
147.03
74.65
6.8106
6.8013
13.40
.0040
.0020
.0020
.0080
.0160
.0080
.020
.020
.040
22.0
22.0
22.0
295.15
295.15
295.15
5
1 (cat)
356.16
10.78
2.8077
92.76
.0016
.0020
.0040
.0080
.030
.020
22.0
22.0
295.15
295.15
(1)
Relative rate of reaction (1000/t)
1.
2.
3.
4.
5.
6.
6.8106
6.8013
13.40
2.8077
92.76
(2)
Initial Concentrations of Reactants (one sample per reactant)
.020M I-
1.
2. .0080M BrO33. .020M H+
(3)
Orders of Reaction
(Show all work or explain in detail how you came to this conclusion)
(1) Iodide (I-)
I- Concentration 1= .020M
relative reaction rate= 3.6191
I Concentration 2= .040M
relative reaction rate= 6.8106
2(M/s)/2(M/s)=1
(2)
Bromate (BrO3-)
m=1
BrO3 concentration 1= .0080M
relative rate of reaction= 3.6191
BrO3 concentration 3= .0160M
relative rate of reaction=6.8013
2(M/s)/2(M/s)=1
(3)
n=1
Hydrogen ion (H+)
H+ concentration 1= .020M
relative rate of reaction= 3.6191
H+ concentration 3= .040M
relative rate of reaction=13.40
4(M/s)/2(M/s)=2
p=2
Using the orders you have calculated and the relative rates, calculate the rate constant, k for
mixtures 1 to 5. Show work for one example in the space below. Calculate the average k.
Mixture
K
1
565484375
2
532078125
3
531351563
4
523437500
5
487447917
Average k = 527959896
Add the values for k and divide by 5
Sample Calculation
Divide rate by the concentrations to their respective orders (iodide and bromate to the first and
hydrogen ion to the second) for mixture 1
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565484375
PART B. DEPENDENCE OF REACTION RATE ON TEMPERATURE
Calculations Table
Reaction
Mixture
Time t
(sec) for
color to
change
Relative Specific rate
rate of
constant, k
reaction,
1000/t
ln k
Temperature
o
C
K
K-1
6
563.75
1.7738
277161863
19.4401122 1.0
274.2
.003647
7
276.31
3.6191
565484375
20.1531932 22.0
294.2
.003399
8
40.73
24.55
3835937500
22.0676797 45.3
318.5
.003140
9
14.39
69.49
10857812500
23.1081507 66.2
339.4
.002946
(1) Using the Rate law you determined in Part A, calculate the specific rate constant, k,
at the temperatures studied and record them in the table above. (show one example
calculation here)
a.
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(2) Using a calculator, find the natural logarithm of each k value and record them on the
table above
a. ln(277161863)= 19.4401122
(3) Determine the Kelvin Temperatures for each of the temperatures above. Record in
the table above and show work for one calculation below.
a. 1.0°C+273.15=274.2K
(4) Determine the inverse of the Kelvin Temperatures for each of the trials above.
Record in the table above and show work for one calculation below.
a.
.003647
(5) Use Excel to Graph the ln k vs 1/T. Attach graph below. Be sure to have your
linear regression data on your graph.
(6) Use your Graph to calculate the Activation Energy, Ea, (see EQ. 10).
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Error Analysis:
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Two errors that might have impacted our results were the amount of substance involved in the
reaction and with a few trials, merely paying attention. With this lab in particular, it is important
to have very close to the exact amount. Concentrations can greatly affect the rate of reaction, and
this lead it to be slightly off double or quadruple. It rounded to the correct number, however if it
had been much more, it would have rounded to 3 or perhaps 5. Additionally, a few times, no one
in my group was paying attention to the solution as it was taking quite a while to change color,
and was a few seconds late to stop the timer. This caused our time to be longer, leading to a
smaller relative reaction rate. This caused it to be closer to the rate tripling instead of
quadrupling in some instances. In our case, fortunately, it rounded to 4 rather than 3 and 2 rather
than 1, but if we stopped the timer or put just a little less solution in, it very well could have
rounded down.
2.
Give at least one way you could improve your lab results. Offer one lab improvement
with explanation on how it could improve results or learning.
It would be much easier to see the color change if we held it over a white paper, like in the
titration lab. Our times would then be better and hopefully the results more accurate.
Conclusions:
In this experiment, the relative rate of a reaction was determined by measuring the time required
to produce a constant amount of I2 each time. What chemical was added to the reaction
determined this specific quantity? (Hint: it was not starch.) Explain how you came to this
conclusion.
The Na2S2O3 determined what quantity of I2 was necessary. The starch turned black, but
the Na2S2O3 was the reaction that delayed it turning black for length of time, until all the
Na2S2O3 was used up in the reaction.
What was the minimum number of moles of I2 that had to be produced in each reaction (the same
amount for each) before you could see the color change?
.000005moles of I2
Compare the times for Mixture 1 catalyzed and uncatalyzed. Would you expect the activation
energy for the catalyzed reaction be greater than, less than, or equal to the activation energy for
the uncatalyzed reaction? Why?
The activation energy for the catalyzed reaction would be less than the activation energy
of the uncatalyzed reaction, because the catalyzed reacted faster. The temperature was the
same, the concentrations were the same, so the activation energy must have been lowered
if the reaction rate was to be greater. Also, that is how a catalyst works; it lowers the
activation energy for a specific reaction so it goes faster.
In a study of the kinetics of the reaction represented below, the following data were obtained at
298K.
5Br-(aq) + BrO3-(aq) + 6H+(aq)  3Br2(l) + 3H2O(l)
Initial [Br-]
Initial [BrO3-]
Initial [H+]
(mol L-1)
(mol L-1)
(mol L-1)
Rate of
disappearance of
BrO3- (mol L-1s-1)
1
0.00100
0.00500
0.100
2.50 x 10-4
2
0.00200
0.00500
0.100
5.00 x 10-4
3
0.00100
0.00750
0.100
3.75 x 10-4
4
0.00100
0.01500
0.200
3.00 x 10-3
Experiment
a.) From the data given above, determine the order of the reaction for each reactant. Show
your reasoning.
Br-: m=1
The reaction rate doubles whenever the concentration of bromine ions doubles
BrO3-: n=1
The reaction rate increases by a factor of 1.5 whenever the concentration of bromate ions
increases by a factor of 1.5
H+: p=2
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b.) Write the rate law for the overall reaction.
Rate=k[Br]1[BrO3]1[H]2
c.) Determine the value of the specific rate constant for the reaction at 298K. include the
correct units.
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d.) If initially [Br-] = 0.0025M, [BrO3-] = 0.0100M and [H+] = 0.300M, what would be the
rate of appearance of Br2 in mol L-1 s-1.
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