Measuring Reaction Rate

Austin Peay State University Department of Chemistry Chem 1041

Measuring Reaction Rate

Revision SP11 IB Page 1 of 9

Cautions HCl is corrosive and toxic

Purpose The purpose of this experiment is to determine the experimental rate of a chemical reaction.

Introduction

Kinetics is the area of chemistry that is concerned with the rates of reactions. Chemical processes do not happen

instantaneously, but rather at a certain rate over time. Different reactions occur at varying speeds: the rusting of iron is

reasonably slow, while the combustion of TNT is extremely fast. The rate of a reaction can vary depending on many

factors including concentration, temperature, surface area of the solid, whether a catalyst is present, and pressure.

Reactions will only occur when the reactants come into contact each other and collide with enough energy to result in a

reaction.

Concentration influences reaction rate by altering how often the reactant particles collide. When the

concentration of a solution is raised, the number of particles in a given volume also increases. An increase in the number

of particles per unit volume will consequently increase the probability of collisions occurring between the particles. The

frequency of collisions is directly related to reaction rate; if the particles collide more often, the reaction will occur faster.

Therefore, an increase in the concentration of a solution will increase the number of collisions per unit time and therefore

the rate of reaction will increase.

Temperature affects reaction rate by changing the overall speed of the reactant particles. Temperature is related

to the average kinetic energy of the reactants. When the temperature is increased, the overall kinetic energy will increase.

Increased kinetic energy means the reactants are moving at a higher speed, so the number of collisions and the rate of

reaction will also increase.

A catalyst can increase a reactions rate by providing an alternative path to the products which requires less

energy. Catalysts may take an active part in the reaction, but at the end of the reaction, they will be recovered in the

same form at the beginning of the reaction.

In this experiment, we will be measuring the rate at which calcium carbonate reacts with hydrochloric acid in the

following process.

2 2

We will be varying the concentration of hydrogen ions by using different concentrations of hydrochloric acid. The

rate of reaction will be measured by observing the reaction on a balance. As the carbon dioxide is released, the total

weight of the reaction will change because of carbon dioxide leaving the solution. In this experiment, we assume the rate

of reaction is related to a change in concentration with respect to time. This can be depicted by the following equation.

∆

∆

Where:

t = time (s)

[H+] = molarity of hydrogen ions

k = the rate constant for the reaction.




Procedure

1. Obtain two clean 100 mL graduated cylinders. Using a wax pencil, label the first cylinder “A” (for acid) and the

second cylinder “W” (for water).

2. Using cylinder “A” measure 25.0 mL of the 1.00 M HCl. Make sure to record this solution’s exact concentration on

your data sheets.

3. Using cylinder “W” measure 25.0 mL of distilled water.

4. Pour both solutions into a 250 ml beaker. This will make a 0.50 M HCl solution

5. Tare a weigh boat on a scale and then measure out 1.5 g of limestone (calcium carbonate). Make sure you record

the exact weight.

6. Place the beaker with acid solution on a zeroed scale and record the mass. Add the mass of the calcium

carbonate to the mass of the beaker and solution to find the mass at time zero.

7. Quickly add the calcium carbonate to the mixture and start the timer.

8. Record the weight of the solution every 30 seconds for 5 minutes or until the solution has stopped fizzing.

9. When data collection is complete, the solution may be poured down the drain with an excess of water.

10. Repeat the complete procedure for experiments 2, 3, and 4 using the volumes of acid and water indicated in the

table below. Make sure to rinse and dry the beaker between trials.

Experiment Volume of 1.00 M HCl Volume of DI water

1 25.0 mL 25.0 mL

2 12.5 mL 37.5 mL

3 0.0 mL 50.0 mL

4 50.0 mL 0.0 mL

Waste Disposal

Pour all solutions down the sink with an excess of water.

Clean-Up

Wash all glassware with soap then rinse 3 times with tap water, and once with deionized water.




Data Sheet 1

Name: __________________________________ Lab Partner: ______________________________

Experiment 1 Experiment 2

Volume of Water Volume of Water

Volume of Acid Volume of Acid

Mass of beaker Mass of beaker

Mass of beaker plus solution


Mass of CaCO3 Mass of CaCO3

Mass beaker, solution plus

CaCO3


CaCO3

Clo

ck T

ime

Ela

pse

d

Tim

e (s

)

Mas

s o

f re

acti

on

Mas

s o

f C

O2

pro

du

ced

Mo

le o

f C

O2

pro

du

ced

Clo

ck T

ime

Ela

pse

d

Tim

e (s

)

Mas

s o

f re

acti

on

Mas

s o

f C

O2

pro

du

ced

Mo

le o

f C

O2

pro

du

ced




Data Sheet 2

Experiment 3 Experiment 4

Volume of Water Volume of Water

Volume of Acid Volume of Acid

Mass of beaker Mass of beaker



Mass of CaCO3 Mass of CaCO3


CaCO3


CaCO3

Clo

ck T

ime

Ela

pse

d

Tim

e (s

)

Mas

s o

f re

acti

on

Mas

s o

f C

O2

pro

du

ced

Mo

le o

f C

O2

pro

du

ced

Clo

ck T

ime

Ela

pse

d

Tim

e (s

)

Mas

s o

f re

acti

on

Mas

s o

f C

O2

pro

du

ced

Mo

le o

f C

O2

pro

du

ced




Data Sheet 3

Molarity of HCl used

EXPT 1 EXPT 2 EXPT 3 EXPT 4 Units

Volume of HCl

Volume of water

Concentration of HCl

Concentration of H+

Rate of CO2 production (Slope)

Value of rate constant (k)

Calculations: Show al your work for calculations.

Prepare a graph of moles of CO2 vs. time for each of your trials. Please note, your instructor may require these

graphs to be generated with a spreadsheet program, like Excel.

Using your graph, draw a best-fit line through all data points prior to mixing and calculate the slope of the line.

This slope corresponds to the rate of reaction for that experiment

Use the reaction rate from the previous step and the molarity of the hydrogen ion in the solution to calculate the

rate constant for each experiment.




Calculations




Post-lab Assignment Name: ______________________________

1. How does the concentration of H+ influence the reaction rate? Use your data to explain your answer.

2. Should the rate constants for each experiment be the same? Why or why not?

3. Compare the rate constants you calculated for each experiment. Were these numbers similar to each other?

Why or why not?

4. How could this experiment be improved to minimize error? Provide at least 2 suggestions.




This page was intentionally left blank for double-sided printing,




Pre-lab Assignment Name: _________________________________ 1. Why is the rate of a reaction important in the chemical industry?

2. How will an increase in temperature affect the rate of reaction? Explain your reasoning using an example.

3. Name two other variables besides temperature that effect reaction rate. Describe how each influences reaction rate.

4. During a similar experiment a student produces 5 moles of CO2 per minute at a pH of -1. Determine the rate constant

for the reaction. (Hint: use the pH to determine [H+])

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Measuring Reaction Rate