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Transcript
10. Factors Affecting the Rate of a Chemical Reaction
What you will accomplish in this experiment
You’ll investigate two of the four factors that affect the rate of a chemical reaction.
Specifically, you’ll study the impact of:
•
Increasing the concentration of each reactant (increasing the number of moles of reactant per liter of solution).
•
Increasing the temperature of the reaction by heating.
You’ll determine the rate of a reaction by measuring the decrease in reactant concentration over a period of time
(as it is being converted into product). You’ll then need to make a conclusion as to whether reaction rate increases
or decreases when the concentration of each reactant is increased, and when the overall temperature of the reaction
is increased.
Concepts you need to know to be prepared
The study of the rate of a chemical reaction is called chemical kinetics.
The rate of a chemical reaction can be studied by measuring either:
• How quickly a reactant is used up, or
• How quickly a product is formed.
More specifically, the change in concentration of a reactant (or of a product) is measured over a certain period of
time as the reaction is occurring.
For example, the figure to the left below illustrates how one might measure the rate of decomposition of hydrogen
peroxide, H2O2. The hydrogen peroxide molecules decompose to produce water and oxygen gas.
2 H2O2 (aq)
2 H2O (l) + O2 (g)
The change in mass as the bubbles of oxygen gas diffuse out of the reaction flask provides a way to measure how
quickly the product is being formed. The graph to the right shows the increase in oxygen concentration (as well as
the decrease in hydroxide peroxide concentration) over a period of time.
1L
C. Graham Brittain
Page 1 of 12
11/14/2010
Chemical reactions can only occur when the reactant molecules or ions collide with one another. This idea is
referred to as the Collision Theory of Chemical Kinetics:
•
In order for reactant “particles” (atoms, ions, or molecules) to react, they must
collide.
•
At the instant of collision, the kinetic energy of particle motion is converted to
collision energy.
•
The collision energy must be greater than or equal to the Activation Energy (Ea)
in order to have an effective (successful) collision – one that has sufficient energy
to break key bonds and rearrange the atoms.
•
The Activation Energy (Ea) is the minimum energy that must be supplied by the
collision in order for the reaction to occur by a particular pathway. It is the
difference between the energy level of the reactants and the energy level of the
“transition state” (the intermediate that exists at the moment of collision, when
the kinetic energy, converted into collision energy, is used to break bonds and
rearrange the atoms).
•
A certain fraction of all molecules in a sample will have the necessary Activation
Energy to react, and that fraction will increase with increasing temperature.
•
The proper spatial orientation of the colliding molecules is also essential to an
effective collision. (See the figure to the right.)
Not every molecular collision can be “effective” or “successful,” in that it leads to the breaking of bonds in the
reactant molecules and the formation of new bonds in the product molecules. In fact, the result of many collisions
is that the molecules simply bounce apart without reacting.
Only molecules that collide with sufficient energy (to break the bonds) and with the correct orientation (to
exchange atoms) will have an effective (successful) collision that leads to the formation of products.
There are four factors that can influence the rate of a chemical reaction:
1. The Ionic or Molecular Nature of the Reactants
In general, chemical reactions that occur between
ions in aqueous solution are extremely rapid: they
happen almost instantaneously when the solutions
containing the reacting ions are mixed. This is
because the attractive forces between the ions (the
ionic bonds) are broken as the compound dissolves
and the ions become hydrated by the water molecules
(as you learned in last week’s experiment).
Additionally, most ions exert attractive forces
equally in all directions. There is no preferred
orientation for collision, and usually no covalent
bonds need to be broken (unless the ions are
polyatomic).
In contrast, reactions between molecules (that require
covalent bonds to be broken or formed) tend to
proceed much more slowly. Certain structural
characteristics of the reactant molecules, such as
bond polarity, geometry, overall size, and orientation
will influence the reaction rate as well.
C. Graham Brittain
Page 2 of 12
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2. The Concentration of the Reactants
For many reactions, the rate will increase as the concentration of the reactants increases. If concentration is
expressed as the number of moles of a reactant in a particular volume of solution, realize that increasing the
reactant concentration means raising the number of reactant molecules that are contained in that same volume.
For many reactions – but by no means all – there is a direct relationship between concentration and reaction rate:
when the concentration of a reactant is doubled, the reaction rate is also doubled.
This result is easily understood when the collision theory of reactions is considered: if there are twice as many
molecules contained in the same volume, there will be twice as many collisions occurring over the same time
period.
3. The Temperature of the Reactants
In nearly all cases, the rate of a reaction increases with increasing temperature.
A rule of thumb for many reactions: Each time the temperature is increased by 10oC, the rate of the reaction
doubles!
This temperature effect can be extremely powerful. Consider a reaction conducted at room temperature (20oC)
versus the same reaction run at 90oC. Since there are seven 10oC increments between 20oC and 90oC, the reaction
will occur 27 = 128 times faster!
Expressed differently, if the reaction takes 20 hours
to convert 100 grams of reactant to product at 20oC,
then the same reaction would take only 10 minutes at
90oC!
Again, the collision theory of reactions explains this
phenomenon: the average kinetic energy of
molecules (energy of motion) is a direct function of
temperature.
When the temperature is increased, molecules move
more rapidly, which means that they collide more
frequently, but more importantly, that they collide
with GREATER ENERGY. Higher energy collisions
are more likely to be effective collisions, and lead to
the rearrangement of atoms and the formation of
products.
4. Presence of a Catalyst
A substance that increases the rate of a reaction (without itself being used up by the reaction) is called a catalyst.
Catalysts work by providing an alternate pathway for the reaction: one in which less energy is required in order to
have an effective collision. Put another way, a catalyst lowers the Activation Energy that is required for reactants
to be converted into products.
Some catalysts can increase the rate of a great number of different reactions, while other catalysts, such as the
enzymes in living systems, are specific to just one reaction, or even to just a single type of reactant molecule.
In this experiment, you’ll investigate the effect of concentration and temperature on the rate of the reaction
between solutions of potassium permanganate (KMnO4, an ionic compound) and oxalic acid (H2C2O4, a molecular
compound).
5 H2C2O4 (aq) + 2 MnO4- (aq) + 6 H+ (aq) → 2 Mn2+ (aq) + 10 CO2 (aq) + 8 H2O (aq)
C. Graham Brittain
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You may remember from last week’s solubility experiment that potassium permanganate dissolves in water to
yield an intensely purple-colored solution. You’ll use this purple color to monitor how quickly the permanganate
ion reacts with oxalic acid.
As the permanganate ion is used up by this reaction, the purple color will fade to red, and then to a yellow-brown.
This color change will signal that the reaction has gone to completion, and that the permanganate ion has been
used up (its concentration has reached 0 moles/liter).
To establish the reaction rate, you’ll measure the amount of time that elapses between the mixing of the reactants
and the disappearance of the permanganate ion’s purple color.
The rate of the reaction can be calculated by the following equation:
ܴܽ‫= ݁ݐ‬
∆ሾ‫ܱ݊ܯܭ‬ସ ሿ
∆ ‫݁݉݅ݐ‬
The symbol “∆” in the equation above means “change in,” and the square brackets around the chemical formula
for potassium permanganate means “concentration of” in Molarity units (of moles/liter).
So the reaction rate is calculated by dividing the change in the permanganate ion concentration by the elapsed
time.
But the change in the permanganate concentration is simply the difference between the initial concentration (at the
start of the reaction) and the final concentration (at the end of the reaction). And as discussed above, the final
concentration is assumed to be 0 moles/liter (when the disappearance of the purple color signals that the reaction
has gone to completion).
Thus the change in concentration will be the same numerical value as the initial concentration.
And the change in time is simply the number of seconds that elapses from the start of the reaction at time zero
(when the reactant solutions are mixed) until the purple color has faded to yellow-brown.
Thus the rate equation simplifies to:
ܴܽ‫= ݁ݐ‬
ሾ‫ܱ݊ܯܭ‬ସ ሿ௜௡௜௧௜௔௟
݈݁ܽ‫݁݉݅ݐ ݀݁ݏ݌‬
As mentioned above, you’ll study the effects of concentration and temperature on the rate of this reaction.
Effect of Concentration:
•
You’ll conduct a series of experiments in which the concentration of one reactant will be varied, while the
temperature and concentration of the second reactant will be held constant.
Effect of Temperature:
•
You’ll conduct a series of experiments in which only the temperature of the reaction will be varied, while
the concentration of all reactants will be held constant.
For each set of reaction conditions, you must determine the initial permanganate concentration and the elapsed
time in order to determine the rate of the reaction.
Procedure that you will follow
Safety Precautions: Both the potassium permanganate and oxalic acid solutions can be both corrosive and toxic.
Permanganate also stains both skin and clothing. If either of these solutions comes into contact with your skin,
you must wash your skin with water for a full fifteen minutes and notify your lab instructor immediately. Gloves,
safety glasses or goggles, and lab coats are required at all times during this experiment.
C. Graham Brittain
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A. Experimental Set-up
1. Dispense 5.0 ml of the oxalic acid solution (H2C2O4) from the pump bottle into a labeled 4” test tube. Record
the concentration of the H2C2O4 solution in moles/liter (M).
2. Dispense 5.0 ml of the potassium permanganate solution (KMnO4) from the pump bottle into a second labeled
4” test tube. Record the concentration of the KMnO4 solution in moles/liter (M).
3. Add approximately 5.0 mL of distilled water to a third labeled 4” test tube.
4. Place all three test tubes in your rack. Note that it is essential to label the tubes, as both the distilled water and
oxalic acid solution are colorless, and could easily be confused. Place a Beral pipet in each test tube for
measuring drop size quantities of each reagent.
B. Effect of Concentration on Reaction Rate
1. Conduct THREE trials of EACH of the experiments listed in the table below. For each experiment, the
procedure will be as follows:
2. To a clean, dry 4” test tube, add the specified number of drops of distilled water and the H2C2O4 solution.
Swirl the tube to mix the reagents.
3. One lab partner is to add the specified number of drops of KMnO4 solution to the test tube (swirling to mix).
The second lab partner begins timing the reaction at the instant the KMnO4 is added. Then, at the instant that
the purple color has faded (from red, and then) to yellow-brown, the elapsed time in seconds must be recorded.
Experiment 1 (three trials)
Experiment 2 (three trials)
Experiment 3 (three trials)
Distilled water
6 drops
1 drop
5 drops
H2C2O4 solution
5 drops
10 drops
5 drops
KMnO4 solution
1 drop
1 drop
2 drops
12 drops
12 drops
12 drops
Total Volume
4. Dispose of the used solution in a waste beaker at your lab bench. Rinse and dry the reaction test tube.
5. Record the temperature of the remaining distilled water in the stock supply test tube. You’ll assume that this is
the temperature at which all three trials of all three experiments were conducted. The three trails of
Experiment 1 will also serve as your room temperature data for Part C.
C. Effect of Temperature on Reaction Rate
1. Fill a large beaker half-full with tap water. Warm the beaker on a hot plate until the water temperature is 3032oC above room temperature. Hold the temperature constant to + 1oC by adding ice or cold water or by
carefully adjusting the heat.
2. Place the test tube containing the stock solution of KMnO4 into the warm water bath.
3. To a clean, dry 4” test tube, add 6 drops of distilled water and 5 drops of the H2C2O4 solution. Swirl the tube to
mix the reagents. Place this test tube in the warm water bath, and allow it two minutes to equilibrate to the
temperature of the bath.
4. One lab partner is to add 1 drop of KMnO4 solution to the test tube (swirling to mix). The second lab partner
begins timing the reaction at the instant the KMnO4 is added. The reaction tube may be lifted from the water
bath briefly to add the reagent or to monitor the color, but it should be returned to the bath as quickly as
C. Graham Brittain
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possible to avoid the loss of heat. Then, at the instant that the purple color has faded (from red, and then) to
yellow-brown, the elapsed time in seconds must be recorded.
5. Dispose of the used solution in the waste beaker at your lab bench. Rinse and dry the reaction test tube.
Repeat steps 3 and 4 to conduct two more trials of the experiment at this temperature
6. Add ice to the warm water bath to reduce its temperature to 20-22oC above room temperature. Hold the
temperature constant as you did in step 1 above. Follow the same procedure that you did in steps 2 through 5,
to conduct a total of three trials of the experiment at this temperature.
7. Add ice to the warm water bath to reduce its temperature to 10-12oC above room temperature. Hold the
temperature constant as you did in step 1 above. Follow the same procedure that you did in steps 2 through 5,
to conduct a total of three trials of the experiment at this temperature.
D. Calculation of Results
1. Realize that when the KMnO4 solution was added to the distilled water and H2C2O4 mixture, it was diluted
from its original stock concentration by the volumes of water and oxalic acid solution. That means you’ll need
to calculate the concentration of KMnO4 that was present in the reaction test tube at the instant the reaction
began. This requires use of the dilution formula you learned in your lecture course:
‫ܥ‬ଵ ܸଵ = ‫ܥ‬ଶ ܸଶ
In this formula:
• C1 is the stock concentration of KMnO4
• V1 is the volume of KMnO4 added to the distilled water and H2C2O4 mixture.
• V2 is the total volume of solution in your reaction test tube (the volumes of distilled water, H2C2O4, and
KMnO4).
When these substitutions are made, and the dilution formula is rearranged to solve for “C2”, the result is:
‫ܥ‬ଶ = ሾ‫ܱ݊ܯܭ‬ସ ሿ௜௡௜௧௜௔௟ =
ሺ‫ܱ݊ܯܭ ݂݋ ݊݋݅ݐܽݎݐ݊݁ܿ݊݋ܿ ݇ܿ݋ݐݏ‬ସ ሻ × ሺ# ݀‫ܱ݊ܯܭ ݂݋ ݏ݌݋ݎ‬ସ ݅݊ ‫ܾ݁ݑݐ ݊݋݅ݐܿܽ݁ݎ‬ሻ
‫ ݈ܽݐ݋ݐ‬#݀‫ܾ݁ݑݐ ݊݋݅ݐܿܽ݁ݎ ݊݅ ݊݋݅ݐݑ݈݋ݏ ݂݋ ݏ݌݋ݎ‬
2. Finally, you’ll calculate the Average Rate of Reaction for EACH Experiment:
‫= ݊݋݅ݐܴܿܽ݁ ݂݋ ݁ݐܴܽ ݁݃ܽݎ݁ݒܣ‬
ሾ‫ܱ݊ܯܭ‬ସ ሿ௜௡௜௧௜௔௟
‫݁݉݅ܶ ݀݁ݏ݌݈ܽܧ ܧܩܣܴܧܸܣ‬
Note that for the Effect of Concentration study (part B of the procedure), you’ll calculate the Average Rate of
Reaction for EACH of the three Experiments, using
•
The [KMnO4]initial used in EACH Experiment
•
The average elapsed time (the average time of reaction from the three TRIALS you ran for each
EXPERIMENT).
And note that for the Effect of Temperature study (part C of the procedure), you’ll calculate the Average Rate
of Reaction at EACH of the three Temperatures, using
•
The [KMnO4]initial (which was held constant for all trials at all temperatures)
•
The average elapsed time (the average time of reaction from the three TRIALS you ran at each
TEMPERATURE).
C. Graham Brittain
Page 6 of 12
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Report Sheet 10: Factors Affecting the Rate of a Chemical
Reaction
Student ______________________________ Lab Partner__________________________ Date Lab Performed__________
Section #_________ Lab Instructor__________________________________________ Date Report Received ___________
Lab Notebook: Data and Observations
Stock Concentrations of Reagents
Oxalic Acid Solution (moles/liter, M)
________________________________
Permanganate Solution (moles/liter, M)
________________________________
Effect of Concentration
Experiment 1:
Drops of Distilled Water
Drops of H2C2O4
Drops of KMnO4
Trial 1
6
5
1
Trial 2
6
5
1
Trial 3
6
5
1
Elapsed Time, sec
Average Elapsed Time
for Experiment 1:
Experiment 2:
Drops of Distilled Water
Drops of H2C2O4
Drops of KMnO4
Trial 1
1
10
1
Trial 2
1
10
1
Trial 3
1
10
1
Elapsed Time, sec
Average Elapsed Time
for Experiment 2:
Experiment 3:
Drops of Distilled Water
Drops of H2C2O4
Drops of KMnO4
Trial 1
5
5
2
Trial 2
5
5
2
Trial 3
5
5
2
Elapsed Time, sec
Average Elapsed Time
for Experiment 3:
C. Graham Brittain
Page 7 of 12
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Effect of Temperature
Temperature 4 (30-32 oC > room temperature): _________________________
Drops of Distilled Water
Drops of H2C2O4
Drops of KMnO4
Trial 1
6
5
1
Trial 2
6
5
1
Trial 3
6
5
1
Elapsed Time, sec
Average Elapsed Time
for Temperature 1:
Temperature 3 (20-22 oC > room temperature): _________________________
Drops of Distilled Water
Drops of H2C2O4
Drops of KMnO4
Trial 1
6
5
1
Trial 2
6
5
1
Trial 3
6
5
1
Elapsed Time, sec
Average Elapsed Time
for Temperature 2:
Temperature 2 (10-12 oC > room temperature): _________________________
Drops of Distilled Water
Drops of H2C2O4
Drops of KMnO4
Trial 1
6
5
1
Trial 2
6
5
1
Trial 3
6
5
1
Elapsed Time, sec
Average Elapsed Time
for Temperature 3:
Temperature 1 (room temperature): _________________________ (Same as Effect of Concentration Experiment 1)
Drops of Distilled Water
Drops of H2C2O4
Drops of KMnO4
Trial 1
6
5
1
Trial 2
6
5
1
Trial 3
6
5
1
Elapsed Time, sec
Average Elapsed Time
for Temperature 4:
C. Graham Brittain
Page 8 of 12
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Formal Report: Results and Conclusions for Effect of Concentration
Calculation of Initial Concentration of KMnO4, in moles/liter:
ሾ‫ܱ݊ܯܭ‬ସ ሿ௜௡௜௧௜௔௟ =
ሺ‫ܱ݊ܯܭ ݂݋ ݊݋݅ݐܽݎݐ݊݁ܿ݊݋ܿ ݇ܿ݋ݐݏ‬ସ ሻ × ሺ# ݀‫ܱ݊ܯܭ ݂݋ ݏ݌݋ݎ‬ସ ݅݊ ‫ܾ݁ݑݐ ݊݋݅ݐܿܽ݁ݎ‬ሻ
‫ ݈ܽݐ݋ݐ‬#݀‫ܾ݁ݑݐ ݊݋݅ݐܿܽ݁ݎ ݊݅ ݊݋݅ݐݑ݈݋ݏ ݂݋ ݏ݌݋ݎ‬
Experiment 1:
Experiment 2:
Experiment 3:
Calculation of Average Rate of Reaction, in moles/liter-sec:
‫= ݊݋݅ݐܴܿܽ݁ ݂݋ ݁ݐܴܽ ݁݃ܽݎ݁ݒܣ‬
ሾ‫ܱ݊ܯܭ‬ସ ሿ௜௡௜௧௜௔௟
‫݁݉݅ܶ ݀݁ݏ݌݈ܽܧ ܧܩܣܴܧܸܣ‬
Experiment 1:
Experiment 2:
Experiment 3:
C. Graham Brittain
Page 9 of 12
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Formal Report: Results and Conclusions for Effect of Temperature
Calculation of Initial Concentration of KMnO4, in moles/liter (held constant for all temperatures):
ሾ‫ܱ݊ܯܭ‬ସ ሿ௜௡௜௧௜௔௟ =
ሺ‫ܱ݊ܯܭ ݂݋ ݊݋݅ݐܽݎݐ݊݁ܿ݊݋ܿ ݇ܿ݋ݐݏ‬ସ ሻ × ሺ# ݀‫ܱ݊ܯܭ ݂݋ ݏ݌݋ݎ‬ସ ݅݊ ‫ܾ݁ݑݐ ݊݋݅ݐܿܽ݁ݎ‬ሻ
‫ ݈ܽݐ݋ݐ‬#݀‫ܾ݁ݑݐ ݊݋݅ݐܿܽ݁ݎ ݊݅ ݊݋݅ݐݑ݈݋ݏ ݂݋ ݏ݌݋ݎ‬
Calculation of Average Rate of Reaction, in moles/liter-sec:
‫= ݊݋݅ݐܴܿܽ݁ ݂݋ ݁ݐܴܽ ݁݃ܽݎ݁ݒܣ‬
ሾ‫ܱ݊ܯܭ‬ସ ሿ௜௡௜௧௜௔௟
‫݁݉݅ܶ ݀݁ݏ݌݈ܽܧ ܧܩܣܴܧܸܣ‬
Temperature 4 (30-32 oC > room temperature): _________________________
Temperature 3 (20-22 oC > room temperature): _________________________
Temperature 2 (10-12 oC > room temperature): _________________________
Temperature 1 (room temperature): _________________________
C. Graham Brittain
Page 10 of 12
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Discussion Questions
1. In the Effect of Concentration study, compare your procedure and results for Experiments 1 and 2.
a) Which reactant concentration was increased (KMnO4 or H2C2O4)? ________________________________________
How much was it increased (doubled, tripled, quadrupled)? ________________________________________
Which reactant concentration was held constant (KMnO4 or H2C2O4)? ___________________________________
b) How did this increase in reactant concentration affect the rate of reaction (no effect, doubled, tripled, quadrupled)?
_______________________________________
c) Is this result consistent with your expectation? Explain.
(Recall that for many reactions there is a direct relationship between concentration and rate: when the
concentration of a reactant is doubled, the rate is also doubled.)
2. In the Effect of Concentration study, compare your procedure and results for Experiments 1 and 3.
a) Which reactant concentration was increased (KMnO4 or H2C2O4)? ________________________________________
How much was it increased (doubled, tripled, quadrupled)? ________________________________________
Which reactant concentration was held constant (KMnO4 or H2C2O4)? ___________________________________
b) How did this increase in reactant concentration affect the rate of reaction (no effect, doubled, tripled, quadrupled)?
_______________________________________
c) Is this result consistent with your expectation? Explain.
(Recall that for many reactions there is a direct relationship between concentration and rate: when the
concentration of a reactant is doubled, the rate is also doubled.)
C. Graham Brittain
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3. What would be your estimate of the reaction TIME (number of seconds required for the purple color to
disappear) if you reacted 1.5 drops of KMnO4 with 5 drops of H2C2O4? Explain your reasoning.
(Realize that the # of drops of water would be adjusted so that the total volume of solution would still be 12 drops.)
4. What effect did increasing the temperature have on the RATE of this reaction?
Did the change in rate match the “rule of thumb” (each time the temperature is increased by 10oC, the rate of the
reaction doubles)?
Explain your results.
5. What would be your estimate of the reaction RATE if you decreased the temperature by 10oC by allowing the
reaction test tube to cool in an ice-water bath?
(Think through this – Would the time required for the reaction be longer or shorter? And how would that affect
the rate?) Explain your reasoning.
C. Graham Brittain
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