experiment 23 pre laboratory assignment

Experiment 23 – Factors Affecting Reaction Rate General Chemistry I Lab 1112.55

 University

Instructor:

Student Name:

Partners Name:

Date

Objective

In this experiment, the student will study and observe the factors that affect the reaction rate of certain chemicals.

Introduction

The student will gain knowledge of the term chemical kinetics, which studies the rate of chemical reactions, how the rate of chemical reactions is controlled, and the path or mechanism by which the reaction occurs from the reactants to the products. Fortunately, the reaction rates can vary from very fast to very slow. This can usually be expressed as a change in the concentration of a reactant. As the increase in the concentration increases, the faster the rate of reaction becomes.

Other frameworks could be used to alter the concentration of a chemical reaction. For example, when the chemical reaction changes its color, temperature, pH, gas evolution, odor, and conductivity, it is an indication of a chemical reaction change.

In this experiment, the student will explore “four out of five factors” that can control the reaction rates that are affected. These are: presence of catalysts, concentration of the reactions (except for surface area of the reaction), and nature of reactants and temperature of the chemical system. Some chemicals can undergo rapid changes because of their nature to be more reactive.

In addition, the rate of the chemical reaction doubles for each 10 °C increase in temperature. The additional heat raises the number of collisions between the molecules of the reactants, and also raises the energy of the molecules. The presence of catalysts can be seen when the pathway of the chemical reaction reroutes so that the “alternate” path has a lower activation energy than the unanalyzed reaction. The increase in the concentration of reactants increases the reaction rate. As a result, there is a higher number of reactant molecules.

Procedures

Nature of the Reactants

Different acids affect reaction rates

Half-full set of four labeled small test tubes with 3M H2SO4, 6M HCl, 6M

CH3COOH, and 6M H3PO4

• Submerge a 1-cm strip of magnesium ribbon into each test tube
• Compare the reaction rates and record your observations
• Different metals affect reaction rates
• Half-fill a set of three labeled small test tubes with 6M HCl
• Submerge 1-cm strips of zinc, magnesium, and copper separately

Into the test tubes

• Compare the reaction rates of each metal in HCl
• Record your observations
• Temperature of the Reaction: Hydrochloric Acid–Sodium Thiosulfate Reaction System
• Prepare the solutions
• Pipet 2 mL of 0.1M Na2S2O3 into each of a set of three 150-mm
• Second set of three 150-mm test tubes
• Pipet 2 mL of 0.1M HCl
• Label each set of test tubes
• The first pair of Na2S2O3–HCl pair test tubes is to be combined at room

Temperature in Part B.2

Place a second pair of Na2S2O3–HCl pair test tubes in an ice water bath

For Part B.3

• A third pair of Na2S2O3–HCl pair test tubes in a warm water bath
• (<60°C) for Part B.4
• Allow each pair of test tubes to establish thermal equilibrium (~5
• Minutes) before continuing to Parts B.3 and 4
• Record the time for reaction at room temperature.
• Be prepared to start time for monitoring the reaction rate.
• Combine the first pair of Na2S2O3–HCl pair test tubes

START TIME

• Agitate the mixture for several seconds
• STOP TIME when the cloudiness of the sulfur appears
• Record the time lapse and room temperature, using all certain digits plus
• One uncertain digit.
• Record the time for reaction at the lower temperature
• From the ice bath, pour the HCl solution into the Na2S2O3 solution

START TIME

• Agitate the mixture for several seconds
• Return the reaction mixture to the ice bath
• STOP TIME when the cloudiness of the sulfur appears
• Record the time lapse for the reaction and the temperature of the bath,
• Using all certain digits plus one uncertain digit
• Record the time for reaction at the higher temperature
• From the warm water bath, pour the HCl solution into the Na2S2O3

Solution

• Proceed as in Parts B.2 and B.3.
• Record the appropriate data
• Repeat any of the above reactions as deemed necessary
• Plot the data.
• Plot temperature (y-axis) versus time (x-axis) using appropriate software
• Interpret your data as suggested on the Report Sheet.

SKIP PART C

• Presence of Catalysts
• Add a catalyst
• Place approximate 2mL of a 3% H2O2 in a clean small test tube
• Add 1 or 2 crystals of MnO2 to the solution
• Observe
• Note its instability

SKIP PART E

• Concentration of Reactants: Iodic Acid-Sulfurous Acid System.
• Prepare the test solutions
• Review the preparation of the test solutions in Table 23.1
• Set up five, clean and labeled test tubes in a test tube rack
• Measure the volumes of the 0.01M HIO3, starch, and water with dropping

(Or Beral) pipets

Calibrate the HIO3 dropping pipet to determine the volume (mL)

Per drop

• Calibrate a second dropping (or Beral) pipet with water to determine the
• Number of mL per drop
• Calibrate a third dropping (or Beral) pipet for the 0.01M H2SO3 solution that delivers 1 mL
• Mark the level on the pipet so that quick delivery of 1 mL of the
• H2SO3 solution to each test tube can be made
• Alternatively, use a calibrated 1-mL Beral pipet.
• Record the time for the reaction.
• Place a sheet of white paper beside the test tube
• As one student quickly transfers 1.0 mL of the 0.01M H2SO3 to the
• Respective test tube
• The other notes the time
• Immediately agitate the test tube
• Record the time lapse (seconds) for the deep-blue I3–•starch complex to

Appear

Complete remaining solutions

• Repeat Part F.2 for the remaining reaction mixtures in table 23.1
• Repeat any trials if necessary
• Plot the data
• Using appropriate software, plot for each solution the initial concentration of iodic acid [HIO3]0 (y-axis), versus the time in seconds (x-axis) for the reaction

Data Sheet

Calculations

Results/Discussion

I believe the results from this experiment were fairly successful. I feel that the experimental procedures were used and carried out appropriately. Therefore, excellent results were obtained in the experiment. I’d be so bold as to say it was probably only Part F., the concentration of reactants. Results of that portion of the experiment may be incorrect because the drops may have been miscounted and/or the calculations may have been incorrect. The remainder of the experiment was successfully produced, however.

Conclusion

The aim of the experiment was achieved and documented. I learned in this experiment how to find the factors that affect the reaction rates by using the four or five factors: nature of the reactants, temperature of the chemical system, presence of a catalyst, and concentration of the reactants. As I observed these factors, I noticed how, depending on which factor was used, some reacted faster in one factor (nature of the reactants) than in a different factor (temperature of the chemical system).

Laboratory Questions

Why does the reaction rate of virtually all reactions increase with an increase in temperature? If you were to make a glass of sweetened iced tea the old-fashioned way, by adding sugar and ice cubes to a glass of hot tea, which would you add first?

1.0mL volume of H2SO3 is added to a mixture of 6 drops of .01M HIO3 14 drops of DI water and 1 drop of starch solution. Assuming 20 drop/mL for all solutions, determine the initial molar concentration of [HIO3]

Reaction of Mg with HCl produces hydrogen gas, as we seen in the class. Predict the gas produced by the reaction of calcium carbonate with HCl and write the balanced equation of the reaction.

If the same volume of 4M HCl and H2SO4 is used in the reactions with given amount of Mg, which acid will react more slowly or they both will have the same reaction rate. Explain.

How does a catalyst work?

It speeds the reaction by increasing the temperature of the reaction.

It speeds the reaction by increasing the concentration of the reactants.

It speeds the reaction by lowering the activation energy of the reaction

It speeds the reaction by being consumed during the reaction