Lab 3 Kinetics

7/29/2019 Lab 3 Kinetics

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Marcie Samayoa

Period 1

AP Chemistry

Determining the Rate of a Reaction February 13, 2010

Purpose:

Investigate how the rate of reaction can be measured and how reaction conditions affect reaction rates.

Procedure:

Part 1. Find the Volume of One Drop of Solution

1. Fill the pipet with deionized water.2. Record the mass of a small beaker.3. Put five drops of water in the beaker and find the total mass. Record data.4. Add an additional five drops of water into the beaker, and again determine the mass. Record

data.

5. Put in an additional five more drops and find the mass. Record data.Part 2. Determine the Reaction Rate and Calculate the Rate Law

1. Obtain six pipets and label the pipets KI, H2O, HCl, Starch, Na2S2O3, and KBrO3.2. Using the table given as a guide, fill each numbered well in the first reaction strip. Mix the

solution in each well with a toothpick.

3. To wells 1-9 in the second reaction strip, add 2 drops of 0.040 M KBrO3.4. Record the time when the solution in each cell turns blue.5. When all the cells have turned blue, take the temperature of one of the reaction solutions.

Record this temperature for all the reactions.

6. Repeat steps 4-11 for the combinations that cover experiments 4 and 5 and experiments 6 and 7.Part 3. Determine the Activation Energy

1. Prepare a warm water bath of about 40C.2. Fill the first six wells in the reaction strip with the appropriate number of drops of the reagentlisted. Mix solutions with a toothpick.

3. Place the reaction strip in the water bath with KBrO3 pipet.4. Measure the temperature of the water bath with a thermometer and record the values.5. Add 2 drops of KBrO3 and record time when the first blue color appears.6. Repeat steps 7-9 for the reaction solutions in wells #2 and #3.


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7. Add ice cubes and water to create a cold temperature water bath.8. Place the reaction strip and KBrO3 pipet into the cold temperature water bath.9. Measure the temperature of the water bath with a thermometer and record values.10.Repeat steps 7-9 for wells #4, #5, and #6. Record the time for each reaction.

Part 4. Observe the Effect of a Catalyst on the Rate.

Repeat the procedure in Part 2 for Experiment 1 only and add 1 drop of 0.1 M cupric nitrate solution.

Cu(NO3)2, and 3 drops of distilled water to the mixture. Fill only the first reaction wells. Record the

reaction times.

Data:

Part 1.Data Table. Find the Volume of One Drop of Solution

Mass of empty beaker (a) 28.56

Trial 1 Mass of beaker plus 5 drops of water (b) 28.78

Mass of first 5 drops of water (b) - (a) 0.220

Average mass of 1 drop of water 0.044

Trial 2 Mass of beaker plus 10 drops of water (c) 29

Mass of second 5 drops of water (c) - (b) 0.220


Trial 3 Mass of beaker plus 15 drops of water (d) 29.2

Mass of third 5 drops of water (d) (c) 0.230


Average mass of 1 drop of water (Trial 1-3) .0446 mL


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Part 2.Data Table. Determine the Reaction Rate and Calculate the Rate Law.

Time, seconds

TempC

Reaction RateM/s

Initial Concentrations, M

[I-] [BrO3

-] [H

+]

ExperimentNo.

Trial1

Trial2

Average

1 22 24.1 23.05 24.6C 6.0310-

.0010 .0067 .0167

2 17.6 17.4 17.5 24.6C 7.9410-

.0020 .0067 .0167

3 8.5 8.7 8.6 24.6C 1.6110-

.0030 .0067 .0167

4 19.3 22.7 21 24.6C 6.6210-

.0010 .0133 .0167

5 13.7 11.3 12.5 24.6C 1.1110-

.0010 .0200 .0167

6 7.6 8.4 8 24.6C 1.7410-

.0010 .0067 .0333

7 4.1 4.3 4.2 24.6C 3.3110-

.0010 .0067 .0500

8 6.1 4.9 5.5 24.6C 2.5310-

.0015 .0100 .0250

Part 3.Data Table. Determine the Activation Energy.

Time of Reaction, seconds

Approxi

mateTemperat

ure, C

Measured

TemperatureC

Temperature

, K

1/Temperature

K-1

Trial 1 Trial 2 Average

Time

Rate of

ReactionM/s

Rate

Constant, k(with

units)

Natu

ralLog

k

0 C 0.7 C 273.7 3.6510-

1:02:00 0:56:0 0:59:0 2.3510-

126 M- s- 4.83

10 C 10..1 C 283.1 3.5310-

0:30:01 0:30:02 0:30:15 4.6110-

247 M- s- 5.50

20 C 21.5 C 294.5 3.3910-

0:26:02 0:30:01 0:28:15 4.9410-

264 M-3s-1 5.57

40 C 38.5 C 311.5 3.2110-

0:10:09 0:10:01 0:10:50 1.3210-

706 M-3s-1 6.55


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Part 4.Data Table. Observe the Effect of a Catalyst on the Rate

Reaction Time, seconds

Trial 1 17.5 seconds

Trial 2 17.9 seconds

Calculations:

Find the Rate Constant

Experiment 1 2 3 4 5 6 7

Value of K322 M- s- 212 M- s- 287 M- s- 178 M- s- 199 M- s- 234 M- s- 197 M- s-

Average

Value

232 M-3s-1

Conclusion

From the results of the lab, it can be concluded that the reaction conditions affect reaction

rates. An experiment was conducted at specific concentrations of each of the reactants and then

used to measure the reaction rate. Then, the concentration of one reactant is changed and thus

shows how the reaction rate changes. It was repeated for each reactant and it allowed thecalculation of the order of each reactant. Once the orders were known, which resulted in the rate

law: Rate = k [I-] [BrO3

-] [H

+]

2, the value of the rate constant was calculated in which averaged

out to be 232M-3

s-1

. By using a separate experiment, the activation energy was also determined

by measuring the time it took for a reaction to take place at different temperatures. Data was

recorded and used to find the natural log of k and temperature in the Kelvin scale. The points

were graphed and the slope of the best fit line was found and put into an equation which resulted

in the activation energy of 27.8 kJ. Catalysts also affect reaction rates. Catalysts (Copper II

Nitrate) was added in the same conditions of Experiment 1 and resulted in a faster reaction of

17.5 and 17.9 seconds (Trial 1 and 2, respectively) than the 23 seconds recorded in the original

experiment 1.

Theory Analysis

The rate law for the lab resulted to be: Rate = k [I-] [BrO3

-] [H

+]2. The rate law was found

to see which of the reactants will have an impact on a reaction rate. For example, when

Experiment 1 and Experiment 7 are compared, it is obvious that Experiment 7 reacts a lot faster

than Experiment 1. This is because Experiment 7 had higher volumes of HCl in the solution and


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since [Cl-] is included in the rate law it will thus affect the reaction rate. Experiment 1 also had

HCl but less of it due to the fact that water took in some of its place and since water is not in the

rate law, the addition of water did not affect the reaction rate and with less volumes of the

reactants that appear in the rate law, it does not proceed faster than Experiment 7. In Part 3.

Data Table. Determine the Activation Energy it can be seen that the higher the temperature, the

faster the reaction will take place. This is due to the fact that an increase of temperature increases

the chance of sufficient energy for a reaction and chance of collision at a reactive site, and vice

versa. Both are factors that affect the rate of chemical reactions. Lastly, catalysts also speed up

the rate of chemical reactions. Without a catalyst, experiment 1 takes an average of about 23.05

seconds to react. When Copper II Nitrate (catalyst) is added to Experiment 1, the reaction speeds

up by taking an average of about 17.7 seconds to react. This occurs because the catalyst changed

the mechanism of the reaction providing an alternate pathway that decreases the activation

energy and thus making the reaction go faster.

Questions

1. The higher the concentration, the greater the chance of collision and the faster the reaction rate.If the concentration decreases, there is a less chance of collision and therefore the slower the

reaction rate.

2. Determine the rate law by comparing two trials where the concentration of only one reactantchanges. Find the quotient of the rate of reaction and the quotient of the concentrations of the

changing chemical and solve for the exponent that would come with the concentrations. Use

logarithms if needed to solve for the exponent and the exponent would determine the order it is

in the rate law. This has to be done for every chemical involved in the reaction, individually.

3. It is reasonable since there was a few factors that could have affected the lab dataincluding air bubbles, the starch solution (not completely dissolved), and dirty wells.

4. As the temperature increases, there is a greater chance of sufficient energy for thereaction and chance of collision at the reactive site so therefore faster reaction rates. As

the temperature decreases, there is not sufficient energy for the reaction and therefore less

chance of collision at the reactive site so therefore slower reaction rates.

5. To find the activation energy, the slope of the best fit line ofln kvs. (1/Temperature K-1)needs to be determined. Once it is found, it can be put in the equation: slope = -Ea/R.

Afterwards, solve for Ea and that would equal to the activation energy.

6. The reaction rate is defined as how fast a reaction takes place. Unlike the reaction rate,the specific constant rate is an experimentally determined constant, which is different fordifferent reactions and changes only with temperature

7. A catalyst is a substance that speeds up the rate of a reaction without being consumed inthe reaction. When a catalyst is added to a reaction, it provides an alternate pathway that

decreases the activation energy and thus making the reaction go faster.

8. Skip.


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9. My data was not completely accurate since there was a 16% difference between rates ofreaction between Experiment 8 and the data.

10.By using cleaner wells, it reduces the chances of the chemicals reacting with otherunwanted chemicals in the dirty wells. Air bubbles could have changed the amount of

concentration that was needed by using less of the concentration than of the ones that was

actually needed to put in the experiments. This could have affected the data dramatically

since only a little amount of volume of chemicals was used in the lab so less air bubbles

could improve the results of the lab. The starch not dissolving well in water could have

given different concentrations then the one listed in the lab affecting the lab dramatically

also. If the starch was completely dissolved, it could have produced better lab results.

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Lab 3 Kinetics