the rate law for a reaction allows the speed of a reaction to be related to the concentrations of reacting species, and also provides important information about the mechanism and activation energy. In this case, the reaction to be studied is the iodination of acetone in the presence of an acid catalyst. I2 (aq) + CH3COCH3 (aq) HI (aq) + ICH2COCH3 (aq) Many methods are possible to track the progress of a reaction, often relying on some form of spectroscopy. Conveniently for us, iodine in solution has a yellow-brown color, whereas the acetone and both products are colorless. This means that we can run the reaction using iodine as the limiting reagent, and when we see that the yellow-brown color of the iodine has disappeared, we know that the reaction has stopped and all of the iodine has been consumed. So how does this help us? Recall that to calculate a rate of reaction, we would like to know the change in concentration of a reactant or product during a fixed amount of time. We will choose the reactant iodine in this case, because its brown color allows us to determine when its concentration has fallen to zero. Rate = –Δ[I2] / Δt = – ([I2]final – [I2]initial) / ([tfinal – tinitial) = – (0 – [I2]initial) / (tfinal – 0) = [I2]initial / tfinal In this lab, you will run several trials at different concentrations of reactants and use the method of initial rates to obtain values for k, x, y, and z in the rate law. Rate = k [I2]x [acetone]y [H+]z You will also run the reaction at fixed concentrations but at different temperatures, so that the activation energy and frequency factor can be determined from the Arrhenius equation. This equation can be rearranged into the form below, which mirrors that of a straight line: ln k = (– Ea/R) (1/T) + ln A y = m x + b Thus once the value for k is known at several different temperatures, a graph of ln k vs (1/T) should approximate a straight line, allowing the slope (which equals –Ea/R) and y-intercept (which equals ln A) to be determined. PROCEDURE: Part I: Changing Concentrations In preparation for Part II of the lab, begin heating a 400 mL beaker of water on a hot plate. Your target temperature is approximately 40oC. Place the hot plate out of the way, and watch that cords do not get in the way. The four experiments outlined in the chart below will allow you to determine the order of reaction with respect to each reactant. Use volumetric pipets to measure all of the required chemicals. You will combine everything except the iodine into a 125 mL Erlenmeyer flask. The iodine should be measured into a separate small beaker. Once you are ready to start the reaction, begin timing and at the same time add the iodine into the Erlenmeyer flask – as quickly as safety allows. Swirl the flask above a white piece of paper, and when you judge that the yellow color has disappeared, stop the timer. Repeat each experiment until two trials are within 20 seconds of each other, and then calculate the average time for each experiment. Experiment #: Water (mL) 1.0 M HCl (mL) 4.0 M Acetone (mL) 0.0050 M I2 (mL) 1 10 5 5 5 2 5 10 5 5 3 5 5 10 5 4 5 5 5 10 Part II: Changing Temperature To determine the activation energy and frequency factor, the reaction will be compared at three different temperatures. Your average value for Experiment 1 above provides the reaction time at room temperature; now the same concentrations will be reacted at a colder temperature and a hotter temperature. You will only need to do one trial at each of these temperatures. Using the same volumes as Experiment 1 from Part 1, combine everything except the iodine in a 125 mL Erlenmeyer flask and stopper the flask. Place the flask in the hot water bath and wait five minutes for the temperature to equilibrate. While you are waiting, measure the iodine into a separate small beaker. Once you are ready to start the reaction, record the exact temperature. Begin timing and at the same time add the iodine into the Erlenmeyer flask – as quickly as safety allows. Swirl the flask, and when you judge that the yellow color has disappeared, stop the timer. Repeat the above procedure for a low temperature bath, which you will make by adding ice to a 400 mL beaker of water. Your target temperature is slightly above 15oC. You should not go significantly below this temperature, or the reaction will take too long. The visual cue will happen when ____ (what number of moles?) of _____ (what substance?) has been _____ (produced or consumed?). Experiment #1 water: 10 mL 1.0 M HCl: 5 mL 4.0 M acetone: 5 mL 0.0050 M I2: 5 mL
the rate law for a reaction allows the speed of a reaction to be related to the concentrations of reacting species, and also provides important information about the mechanism and activation energy. In this case, the reaction to be studied is the iodination of acetone in the presence of an acid catalyst. I2 (aq) + CH3COCH3 (aq) HI (aq) + ICH2COCH3 (aq) Many methods are possible to track the progress of a reaction, often relying on some form of spectroscopy. Conveniently for us, iodine in solution has a yellow-brown color, whereas the acetone and both products are colorless. This means that we can run the reaction using iodine as the limiting reagent, and when we see that the yellow-brown color of the iodine has disappeared, we know that the reaction has stopped and all of the iodine has been consumed. So how does this help us? Recall that to calculate a rate of reaction, we would like to know the change in concentration of a reactant or product during a fixed amount of time. We will choose the reactant iodine in this case, because its brown color allows us to determine when its concentration has fallen to zero. Rate = –Δ[I2] / Δt = – ([I2]final – [I2]initial) / ([tfinal – tinitial) = – (0 – [I2]initial) / (tfinal – 0) = [I2]initial / tfinal In this lab, you will run several trials at different concentrations of reactants and use the method of initial rates to obtain values for k, x, y, and z in the rate law. Rate = k [I2]x [acetone]y [H+]z You will also run the reaction at fixed concentrations but at different temperatures, so that the activation energy and frequency factor can be determined from the Arrhenius equation. This equation can be rearranged into the form below, which mirrors that of a straight line: ln k = (– Ea/R) (1/T) + ln A y = m x + b Thus once the value for k is known at several different temperatures, a graph of ln k vs (1/T) should approximate a straight line, allowing the slope (which equals –Ea/R) and y-intercept (which equals ln A) to be determined. PROCEDURE: Part I: Changing Concentrations In preparation for Part II of the lab, begin heating a 400 mL beaker of water on a hot plate. Your target temperature is approximately 40oC. Place the hot plate out of the way, and watch that cords do not get in the way. The four experiments outlined in the chart below will allow you to determine the order of reaction with respect to each reactant. Use volumetric pipets to measure all of the required chemicals. You will combine everything except the iodine into a 125 mL Erlenmeyer flask. The iodine should be measured into a separate small beaker. Once you are ready to start the reaction, begin timing and at the same time add the iodine into the Erlenmeyer flask – as quickly as safety allows. Swirl the flask above a white piece of paper, and when you judge that the yellow color has disappeared, stop the timer. Repeat each experiment until two trials are within 20 seconds of each other, and then calculate the average time for each experiment. Experiment #: Water (mL) 1.0 M HCl (mL) 4.0 M Acetone (mL) 0.0050 M I2 (mL) 1 10 5 5 5 2 5 10 5 5 3 5 5 10 5 4 5 5 5 10 Part II: Changing Temperature To determine the activation energy and frequency factor, the reaction will be compared at three different temperatures. Your average value for Experiment 1 above provides the reaction time at room temperature; now the same concentrations will be reacted at a colder temperature and a hotter temperature. You will only need to do one trial at each of these temperatures. Using the same volumes as Experiment 1 from Part 1, combine everything except the iodine in a 125 mL Erlenmeyer flask and stopper the flask. Place the flask in the hot water bath and wait five minutes for the temperature to equilibrate. While you are waiting, measure the iodine into a separate small beaker. Once you are ready to start the reaction, record the exact temperature. Begin timing and at the same time add the iodine into the Erlenmeyer flask – as quickly as safety allows. Swirl the flask, and when you judge that the yellow color has disappeared, stop the timer. Repeat the above procedure for a low temperature bath, which you will make by adding ice to a 400 mL beaker of water. Your target temperature is slightly above 15oC. You should not go significantly below this temperature, or the reaction will take too long. The visual cue will happen when ____ (what number of moles?) of _____ (what substance?) has been _____ (produced or consumed?). Experiment #1 water: 10 mL 1.0 M HCl: 5 mL 4.0 M acetone: 5 mL 0.0050 M I2: 5 mL
Chemistry & Chemical Reactivity
9th Edition
ISBN:9781133949640
Author:John C. Kotz, Paul M. Treichel, John Townsend, David Treichel
Publisher:John C. Kotz, Paul M. Treichel, John Townsend, David Treichel
Chapter14: Chemical Kinetics: The Rates Of Chemical Reactions
Section: Chapter Questions
Problem 86IL: The acid-catalyzed iodination of acetone CH3COCH3(aq) + I2(aq) CH3COCH2I(aq) + HI(aq) is a common...
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Question
the rate law for a reaction allows the speed of a reaction to be
related to the concentrations of reacting species, and also provides important information
about the mechanism and activation energy. In this case, the reaction to be studied is the
iodination of acetone in the presence of an acid catalyst.
I2 (aq) + CH3COCH3 (aq) HI (aq) + ICH2COCH3 (aq)
Many methods are possible to track the progress of a reaction, often relying on some
form of spectroscopy. Conveniently for us, iodine in solution has a yellow-brown color,
whereas the acetone and both products are colorless. This means that we can run the reaction
using iodine as the limiting reagent, and when we see that the yellow-brown color of the iodine
has disappeared, we know that the reaction has stopped and all of the iodine has been
consumed.
So how does this help us? Recall that to calculate arate of reaction , we would like to
know the change in concentration of a reactant or product during a fixed amount of time. We
will choose the reactant iodine in this case, because its brown color allows us to determine
when its concentration has fallen to zero.
Rate = –Δ[I2] / Δt = – ([I2]final – [I2]initial) / ([tfinal – tinitial) = – (0 – [I2]initial) / (tfinal – 0) = [I2]initial / tfinal
In this lab, you will run several trials at different concentrations of reactants and use the
method of initial rates to obtain values for k, x, y, and z in the rate law.
Rate = k [I2]x [acetone]y [H+]z
You will also run the reaction at fixed concentrations but at different temperatures, so
that the activation energy and frequency factor can be determined from the Arrhenius
equation. This equation can be rearranged into the form below, which mirrors that of a straight
line:
ln k = (– Ea/R) (1/T) + ln A
y = m x + b
Thus once the value for k is known at several different temperatures, a graph of ln k vs (1/T)
should approximate a straight line, allowing the slope (which equals –Ea/R) and y-intercept
(which equals ln A) to be determined.
related to the concentrations of reacting species, and also provides important information
about the mechanism and activation energy. In this case, the reaction to be studied is the
iodination of acetone in the presence of an acid catalyst.
I2 (aq) + CH3COCH3 (aq) HI (aq) + ICH2COCH3 (aq)
Many methods are possible to track the progress of a reaction, often relying on some
form of spectroscopy. Conveniently for us, iodine in solution has a yellow-brown color,
whereas the acetone and both products are colorless. This means that we can run the reaction
using iodine as the limiting reagent, and when we see that the yellow-brown color of the iodine
has disappeared, we know that the reaction has stopped and all of the iodine has been
consumed.
So how does this help us? Recall that to calculate a
know the change in concentration of a reactant or product during a fixed amount of time. We
will choose the reactant iodine in this case, because its brown color allows us to determine
when its concentration has fallen to zero.
Rate = –Δ[I2] / Δt = – ([I2]final – [I2]initial) / ([tfinal – tinitial) = – (0 – [I2]initial) / (tfinal – 0) = [I2]initial / tfinal
In this lab, you will run several trials at different concentrations of reactants and use the
method of initial rates to obtain values for k, x, y, and z in the rate law.
Rate = k [I2]x [acetone]y [H+]z
You will also run the reaction at fixed concentrations but at different temperatures, so
that the activation energy and frequency factor can be determined from the Arrhenius
equation. This equation can be rearranged into the form below, which mirrors that of a straight
line:
ln k = (– Ea/R) (1/T) + ln A
y = m x + b
Thus once the value for k is known at several different temperatures, a graph of ln k vs (1/T)
should approximate a straight line, allowing the slope (which equals –Ea/R) and y-intercept
(which equals ln A) to be determined.
PROCEDURE:
Part I: Changing Concentrations
In preparation for Part II of the lab, begin heating a 400 mL beaker of water on a hot
plate. Your target temperature is approximately 40oC. Place the hot plate out of the way, and
watch that cords do not get in the way.
Part I: Changing Concentrations
In preparation for Part II of the lab, begin heating a 400 mL beaker of water on a hot
plate. Your target temperature is approximately 40oC. Place the hot plate out of the way, and
watch that cords do not get in the way.
The four experiments outlined in the chart below will allow you to determine the order
of reaction with respect to each reactant. Use volumetric pipets to measure all of the required
chemicals. You will combine everything except the iodine into a 125 mL Erlenmeyer flask. The
iodine should be measured into a separate small beaker. Once you are ready to start the
reaction, begin timing and at the same time add the iodine into the Erlenmeyer flask – as
quickly as safety allows. Swirl the flask above a white piece of paper, and when you judge that
the yellow color has disappeared, stop the timer. Repeat each experiment until two trials are
within 20 seconds of each other, and then calculate the average time for each experiment.
Experiment #: Water (mL) 1.0 M HCl (mL) 4.0 M Acetone (mL) 0.0050 M I2 (mL)
1 10 5 5 5
2 5 10 5 5
3 5 5 10 5
4 5 5 5 10
of reaction with respect to each reactant. Use volumetric pipets to measure all of the required
chemicals. You will combine everything except the iodine into a 125 mL Erlenmeyer flask. The
iodine should be measured into a separate small beaker. Once you are ready to start the
reaction, begin timing and at the same time add the iodine into the Erlenmeyer flask – as
quickly as safety allows. Swirl the flask above a white piece of paper, and when you judge that
the yellow color has disappeared, stop the timer. Repeat each experiment until two trials are
within 20 seconds of each other, and then calculate the average time for each experiment.
Experiment #: Water (mL) 1.0 M HCl (mL) 4.0 M Acetone (mL) 0.0050 M I2 (mL)
1 10 5 5 5
2 5 10 5 5
3 5 5 10 5
4 5 5 5 10
Part II: Changing Temperature
To determine the activation energy and frequency factor, the reaction will be compared
at three different temperatures. Your average value for Experiment 1 above provides the
reaction time at room temperature; now the same concentrations will be reacted at a colder
temperature and a hotter temperature. You will only need to do one trial at each of these
temperatures.
To determine the activation energy and frequency factor, the reaction will be compared
at three different temperatures. Your average value for Experiment 1 above provides the
reaction time at room temperature; now the same concentrations will be reacted at a colder
temperature and a hotter temperature. You will only need to do one trial at each of these
temperatures.
Using the same volumes as Experiment 1 from Part 1, combine everything except the
iodine in a 125 mL Erlenmeyer flask and stopper the flask. Place the flask in the hot water bath
and wait five minutes for the temperature to equilibrate. While you are waiting, measure the
iodine into a separate small beaker. Once you are ready to start the reaction, record the exact
temperature. Begin timing and at the same time add the iodine into the Erlenmeyer flask – as
quickly as safety allows. Swirl the flask, and when you judge that the yellow color has
disappeared, stop the timer.
Repeat the above procedure for a low temperature bath, which you will make by adding
ice to a 400 mL beaker of water. Your target temperature is slightly above 15oC. You should
not go significantly below this temperature, or the reaction will take too long.
iodine in a 125 mL Erlenmeyer flask and stopper the flask. Place the flask in the hot water bath
and wait five minutes for the temperature to equilibrate. While you are waiting, measure the
iodine into a separate small beaker. Once you are ready to start the reaction, record the exact
temperature. Begin timing and at the same time add the iodine into the Erlenmeyer flask – as
quickly as safety allows. Swirl the flask, and when you judge that the yellow color has
disappeared, stop the timer.
Repeat the above procedure for a low temperature bath, which you will make by adding
ice to a 400 mL beaker of water. Your target temperature is slightly above 15oC. You should
not go significantly below this temperature, or the reaction will take too long.
The visual cue will happen when ____ (what number of moles?) of _____ (what substance?) has been _____ (produced or consumed?).
Experiment #1
water: 10 mL
1.0 M HCl: 5 mL
4.0 M acetone: 5 mL
0.0050 M I2: 5 mL
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