Materials Included In Kit

Additional Materials Required

Prelab Preparation

Activities B and C

1. Cut copper wire into 20-cm lengths. Show students how to build “cages” to suspend the magnesium ribbon in solution or make one for each station prior to the lab. {12638_Preparation_Figure_5}
2. Cut magnesium ribbon into 12-cm lengths.

Activity C
Three Water Baths: Prepare water baths at different temperatures as follows:

1. Add crushed ice and water to a 400-mL beaker for a bath at about 0 ºC.
2. Add 250 mL of water to a 400-mL beaker and heat it on a hot plate on a lower setting to prepare a hot water bath around 50 ºC.
3. Add 250 mL of room temperature water (20–25 °C) to a 400-mL beaker. Water baths may be shared by both groups.

Safety Precautions

Acetic acid and hydrochloric acid solutions are toxic and corrosive. Avoid contact with skin and eyes. The hydrogen peroxide solution is an oxidizer and a skin and eye irritant. Manganese dioxide is a body tissue irritant and a strong oxidant; avoid contact with organic material. Magnesium ribbon is a flammable solid; keep either dry sand or Flinn Class D Fire Extinguisher Powder available in case of fire. Sodium iodide solution is slightly toxic. Zinc dust from zinc granules could be flammable. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please review current Safety Data Sheets for additional safety, handling and disposal information. Remind students to wash hands thoroughly with soap and water before leaving the lab.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. Excess acetic acid and hydrochloric acid solutions may be neutralized according to Flinn Suggested Disposal Method #24b. Hydrogen peroxide may be reduced according to Flinn Suggested Disposal Method #22a. Magnesium ribbon, manganese dioxide, potassium iodide, yeast, and zinc may be thrown away according to Flinn Suggested Disposal Method #26a. Potassium iodide and sodium iodide solutions may be flushed down the drain with excess water according to Flinn Suggested Disposal Method #26b.

Lab Hints

• For best results, set up two stations for each activity throughout the lab. This will allow 10 groups of students to rotate through five activity stations in a 50-minute lab period, if needed. A double lab period (two 50-minute class periods) will allow time for both a review of basic kinetics concepts before lab and for a collaborative class discussion after lab.
• The activities may be completed in any order. Also, since each activity is a self-contained unit, the lab may be set up with as many or as few of the activities as the teacher desires. Students should need only 7–8 minutes per station—keep the pace fairly brisk to avoid dawdling. Questions in the Observations and Results worksheet may be answered during downtime between stations.
• Prelab Preparation is an essential component of lab safety, and it is also critical for success in the lab. (Standing in front of the lab station is not a good time for students to be reading the activity for the first time.) Having students complete the written prelab assignment and reviewing the safety precautions for this lab will help teachers ensure that students are prepared for and can work safely in the lab.
• For Activities B and C, build copper wire “cages” to keep the magnesium suspended in the hydrochloric acid and to prevent it from floating. The copper wire will not react with the hydrochloric acid.
• In Activities B and C, reaction times are best measured based on the disappearance of the magnesium metal, especially at higher temperatures. Above 50 °C, evolution of gas bubbles was observed even after all the metal had reacted. This may be due to “outgassing” of dissolved oxygen or hydrogen at higher temperatures. Indeed, it was found that the product mixture obtained from a room temperature run produced bubbles when placed in a 50 °C bath.
• In Activity C, the experimental design seems very simple and straightforward. Upon closer examination, however, several factors emerge that have a bearing on the results. In order to isolate the effect of temperature on the reaction rate, it is desirable to carry out the reactions under conditions where the concentration of hydrochloric acid will not change significantly over the course of the reaction. Using 4-cm (0.03-g) strips of magnesium ribbon corresponds to 0.0012 moles of magnesium metal reacting. The amount of hydrochloric acid consumed in the reaction is twice the number of moles of magnesium, or 0.0024 moles. If the volume of 1.0 M hydrochloric acid used is 18 mL, the initial number of moles of HCl present is 0.018 moles, and the amount of HCl consumed is (0.0024/0.018) x 100, or 13% of the total. This is greater than the 5–10% “extent-of-reaction” generally advised for the method of initial rates—the reactant concentration will not be a controlled variable. The surface area of the magnesium metal also changes during the course of the reaction.

Teacher Tips

• An interesting application of the role of chloride ions as catalysts in reactions at metal surfaces can be found in the design of Flameless Ration Heaters (FRH), Flinn Catalog No. AP8695. An FRH consists of magnesium metal embedded in a polymer matrix with iron (an activator) and sodium chloride (a catalyst). When the water “pouch” in the FRH is broken, the magnesium metal reacts exothermically with the water, producing enough heat to cook a meal.
• The reaction rates reported in the Data and Results Table for Activities B and C are proportional rates obtained by taking the inverse of the reaction time in seconds. Actual reaction rates can be calculated from the number of moles of magnesium that have reacted divided by the reaction time.
• In Activity E, the decomposition reaction of hydrogen peroxide, while spontaneous, requires a catalyst. This points out a common misconception of the meaning of “spontaneous.” In thermodynamics, a spontaneous reaction is one that will occur without outside intervention. This does not mean that the reaction is fast!

Answers to Prelab Questions

In Activity A, the effect the surface area of a solid reactant on the reaction rate will be studied by observing the reactions of granular zinc and zinc shot with 2 M hydrochloric acid.

In Activity B, the affect the concentration of a reactant has on the reaction rate will be studied by comparing the reaction of magnesium metal with hydrochloric acid at three different solution concentrations.

In Activity C, the reaction rate of magnesium metal with hydrochloric acid will be studied by observing the reaction at three different temperatures.

In Activity D, two different acids are reacted with granular zinc to study the effect the nature of the reactant has on the overall reaction rate.

In Activity E, the decomposition reaction rate of hydrogen peroxide will be studied by observing the catalytic effect of a series of chemical additives.

Sample Data

Activity B. Reactant Concentration

Activity C. Temperature

Answers to Questions

Activity A. Surface Area

1. Describe the observations and compare and contrast the reactions of granular zinc and zinc shot with hydrochloric acid. Be as specific as possible.

   The test tube containing the granular zinc produced hydrogen gas at a much faster rate than the test tube containing the zinc shot.

2. Based on your observations, estimate the relative rates of reaction for granular zinc compared to zinc shot.

   Answers will vary, but students should estimate a rate 2 or more times faster with the granulated zinc.

3. Describe quantitatively (explain) which form of zinc has the greater surface area.

   The majority of the mass of the zinc shot lies underneath its surface. Granulated zinc particles are essentially smaller pieces of the zinc shot. Dividing an object into smaller pieces always increases the surface area.

4. The average shot piece is a cylinder with a diameter of 7.5 mm, a height of 2.0 mm, and a mass of 0.40 g. The average zinc granule is a sphere with a diameter of 0.84 mm. Using the formulas listed, calculate the total surface area of 0.40 g of each form of zinc.

   Surface area of a cylinder = 2πr² + 2πrh
   Surface area of a sphere = 4πr²
   Volume of a sphere = 4/3πr³
   Volume of cylinder = πr²h
   Density of zinc = 7.141 g/mL
   
   Surface area of zinc shot = [2π(7.5/2)² + 2π(7.5/2)(2.0)]mm²
   Surface area of zinc shot = 135 mm²
   Surface area of zinc granule = 4π(0.84/2)² mm²
   Surface area of zinc granule = 2.2 mm² per granule
   
   The mass per granule = volume of granule x density of zinc

   = [4/3π (0.84/2)³ mm³] x 7.141 g/mL x 1 mL/1000 mm³
   = 0.0022 g per granule

   Number of granules in sample size (0.40 g) = 0.40 g/0.0022 g/granule

   = 180 granules

   Surface area of granules in sample size (0.40 g) = 2.2 mm²/granule x 180 granules

   = 400 mm²

5. What is the ratio of the granular surface area to the shot surface area? How does this compare to your estimate of the relative reaction rates between the two?

   Ratio = 400 mm²/135 mm² = 3:1

Activity B. Reactant Concentration

1. As reaction times increase, the rate of the reaction decreases, in other words, the rate of reaction is inversely proportional to the time required for the reaction to go to completion.

   Rate of reaction = k(1/t), where t is the reaction time in seconds.
   Calculate 1/t for the average reaction time for each concentration of HCl and enter the results in the data table. Graph 1/t versus HCl concentration. Place 1/t on the y-axis and HCl concentration on the x-axis.
   Describe the mathematical relationship between the reaction rate and the concentration.
   rate = k(concentration)
   (Equation of a straight line through zero.)

2. How much did the rate of reaction change when the concentration of HCl was doubled?

   The rate doubled.

3. Using the graph and the definition of reaction rate, estimate the reaction time for the same size piece of magnesium ribbon to react with 1.5 M HCl.
   {12638_Answers_Figure_6}

Activity C. Temperature
Rate of reacton = k(1/t), where t is the reaction time in seconds.

1. Calculate 1/t for the reaction time for each temperature and enter those results in the data table.
2. Convert each temperature to kelvin and enter the values in the data table. (K = °C + 273.15)
3. Graph 1/t on the y-axis versus T(K) on the x-axis for each temperature studied.
   {12638_Answers_Figure_7}
4. Describe the mathematical relationship between the reaction rate and the temperature in kelvins.

   The reaction rate appears to be directly proportional to the temperature, in kelvins, within the temperature range studied. Note: In general, a linear relationship is not expected.

5. The collision theory of reaction rates states that the rate of a reaction depends on the number of collisions between molecules, the average energy of the collisions and the effectiveness of the collisions. (a) How does temperature affect each of these factors in collision theory? (b) Does the effect of temperature on the reaction rate support the collision theory of reaction rates? Explain.

   Increasing the temperature increases the reaction rate, decreasing the temperature decreases the rate, in support of the collision theory of reaction rates. The effect of temperature can be explained in terms of both the number of collisions between molecules and their average energy. Increasing the temperature increases the average kinetic energy of molecules—they move faster. As molecules move faster, the rate of collisions between molecules will increase, thus increasing the reaction rate. More importantly, as the average energy of the colliding molecules increases, more of the colliding molecules have sufficient energy to surpass the activation energy barrier and be converted to products.

Activity D. The Nature of the Reactants

1. What did you observe about the two reactions?

   The rate of gas formation (H₂) was much greater in the test tube containing hydrochloric acid than in the test tube with acetic acid.

2. Based on your observations, estimate the relative rates of reaction between the two forms of acid with zinc.

   Student answers will vary.

3. Zinc reacts with acids according to the following equation:

   Zn(s) + 2HA(aq) → Zn(A)₂(aq) + H₂(g)

   The net ionic reaction is:

   Zn(s) + 2H₃O⁺(aq) → Zn²⁺(aq) + H₂(g)

   Hydrochloric acid is a strong acid and acetic acid is a weak acid. Based on the above information, explain the differences in the two reaction rates.

   A strong acid completely dissociates in water, while a weak acid only partially dissociates. At equal concentrations, hydrochloric acid produces a higher concentration of hydrogen ions, one of the reactants, than does acetic acid. This higher concentration of a reactant produces a higher rate of reaction.

Activity E Observations and Analysis

Part 1

1. What did you observe about the reactions in the three test tubes?

   With manganese dioxide added, the solution turned opaque and a large amount of gas was released. Sodium iodide crystals turned the solution yellow and a large amount of gas was also released. The active yeast also turned the solution opaque, releasing gas bubbles.

2. Based on your observations, estimate the relative rates of reaction for each catalyst.

   Answers will vary.

3. What did you observe with the glowing splint? Would this confirm that a decomposition reaction is taking place? Explain.

   The glowing splint ignited when placed in the mouth of each test tube, indicating the gas produced was oxygen, the product of hydrogen peroxide decomposition.

References

Special thanks to Patricia Mason (retired) Delphi Community H.S., Delphi, IN, and to Kathy Kitzmann, Mercy H.S., Farmington Hills, MI, for providing Flinn with the general idea and specific activity suggestions for “activity stations” lab kits.

Introduction

Many chemical reactions seen in the lab—indicator color changes, formation of a precipitate, evolution of a gas—occur almost immediately upon mixing. Other reactions, however, occur at a slower rate and can be studied by following the progress of the reaction over time. Kinetics is the study of the rates at which the concentrations of reactants and products change in a chemical reaction. What factors determine how fast a chemical reaction will occur?

In general, the greater the rate of a chemical reaction, the less time is needed for a specific amount of reactants to be converted to products. The rate of a reaction can be determined therefore by observing either the disappearance of reactants or the appearance of products as a function of time. Use this set of five “mini-lab” activities to study the effect of surface area, temperature, reactant concentration, the nature of the reactants and the presence of a catalyst have on the rates of chemical reactions.

Concepts

• Kinetics
• Reaction rate
• Single replacement reaction
• Rate law
• Catalyst
• Decomposition reaction

Background

Activity A. Surface Area
When solids react with liquids or gases, only the molecules or atoms on the surface of the solid can react.

In this activity, the effect of surface area on reaction rate will be observed in the reaction of solid zinc metal with a hydrochloric acid solution.

Equal masses of two forms of zinc metal—granular (large surface area) and shot (smaller surface area)—will be added to separate test tubes containing identical volumes of 2 M hydrochloric acid. The relative rates of the two reactions will be compared by observing the rate of formation of the gas bubbles (H₂).

Activity B. Reactant Concentration
Why does a candle burn more brightly in pure oxygen than it does in air? Oxygen is a reactant in the combustion reaction that takes place when the candle burns. The rate of the reaction, and thus the brightness of the flame, depends on the concentration of oxygen. In this activity, the effect of acid concentration on the rate of reaction of magnesium metal with hydrochloric acid will be investigated.

Activity C. Temperature
In this activity, the effect of temperature on the rate of reaction of magnesium metal with hydrochloric acid will be investigated.

Activity D. The Nature of the Reactants
Some substances, by their nature, are more reactive than others. All reactions involve the breaking of old bonds and the formation of new bonds. Molecules with weaker bonds tend to react faster than those with stronger bonds. In this experiment, zinc will react with two different acids. The nature of each acid will determine the relative rates of their reactions.

Activity E. Catalysts
Catalysts are substances that speed up a reaction but are not consumed in the reaction. Catalysts work by lowering the activation energy for the reaction, thus increasing its rate. Hydrogen peroxide undergoes a decomposition reaction to produce water and oxygen.

2H₂O₂(aq) → O₂(g) + 2H₂O(l)

Solutions of hydrogen peroxide can be stored for long periods of time because the decomposition reaction is generally slow in the absence of a catalyst. Observe the effects of various substances on the rate of decomposition of hydrogen peroxide.

Experiment Overview

The purpose of this activity-stations lab is to investigate various factors that may influence the rate of chemical reactions. Five mini-lab activities are set up around the classroom. Each activity focuses on the effect a particular reaction condition has on the reaction rate.

1. Surface Area
2. Reactant Concentration
3. Temperature and Reaction Rates
4. Nature of the Reactants
5. Catalysts

Materials

Prelab Questions

Read the introduction material and Procedure for each activity A–E. Write a brief, one- to two-sentence description of each experiment. Example: In Activity A, the effect the surface area of a solid reactant on the reaction rate will be studied by observing the reaction of granular zinc and zinc shot with 2 M hydrochloric acid.

Safety Precautions

Acetic acid and hydrochloric acid solutions are toxic and corrosive. Avoid contact with skin and eyes. The hydrogen peroxide solution is an oxidizer and a skin and eye irritant. Manganese dioxide is a body tissue irritant and a strong oxidant; avoid contact with organic material. Magnesium ribbon is a flammable solid; keep either dry sand or Flinn Class D Fire Extinguisher Powder available in case of fire. Sodium iodide solution is slightly toxic. Zinc dust from zinc granules could be flammable. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the lab.

Procedure

Activity A. Surface Area

1. Using a graduated cylinder, carefully add 20 mL of 2 M hydrochloric acid to each test tube. Place the two test tubes in the test tube rack.
2. Zero (tare) a weighing dish on the balance and measure the mass of one piece of zinc shot in the weighing dish. Note the mass, then remove the weighing dish.
3. Place the second weighing dish on the balance and tare it. Add zinc granules to the weighing dish until their total mass equals that of the zinc shot.
4. Simultaneously add the zinc samples to the HCl solution in the separate test tubes and observe any changes in the appearance of the reaction mixtures. Compare the rate of bubbles being produced in each test tube for several minutes. Record all observations (Question 1).
5. Wearing gloves, rinse the contents of the test tubes into the large waste beaker provided at the activity station.
6. Answer the questions in the Observations and Analysis section.

Activity B. Reactant Concentration

1. Place three large test tubes in a test tube rack.
2. Using a graduated cylinder, add 18 mL of 0.5 M hydrochloric acid to the first large test tube.
3. Add 18 mL of 1.0 M HCl to the middle test tube.
4. Add 18 mL of 2.0 M HCl to the last test tube.
5. Obtain a 12-cm strip of magnesium ribbon.
6. Using scissors, cut the magnesium ribbon into three, 4.0-cm long pieces. Note: Be as precise as possible. The reaction time will depend on the amount of magnesium reacting in each trial.
7. Twist and fold one end of the 20-cm length of copper wire around a pencil to make a small “cage” into which the magnesium ribbon may be inserted. The other end of the wire must be long enough so that the wire will hang over the side of the test tube and the cage will be below the liquid level marked on the test tube (see Figure 1).
   {12638_Procedure_Figure_1}
8. Fit one piece of magnesium ribbon loosely through a copper wire cage so the magnesium will be held in place but not wrapped around too tightly.
9. Suspend the copper wire cage and the piece of magnesium in the 0.5 M hydrochloric acid solution in test tube 1 and immediately start timing.
10. Measure the time until the magnesium metal has completely reacted and the solution stops bubbling. Record the reaction time in the data table.
11. Remove the copper wire “cage” from the test tube and rinse with a small amount of distilled water. Pat dry with a paper towel.
12. Repeat steps 8–11 with the 1.0 M HCl test tube.
13. Repeat steps 8–11 with the test tube containing 2.0 M HCl.

Activity C. Temperature

1. Place the three large test tubes in the test tube rack.
2. Using a graduated cylinder, add 18 mL of 1 M hydrochloric acid to each of three large test tubes.
3. With a test tube clamp attached to each test tube, place one test tube in a beaker of room temperature water, one test tube in the ice-water bath, and one test tube in the hot-water bath.
4. Allow the test tubes to sit in their respective water baths for three minutes to reach thermal equilibrium.
5. Obtain a 12-cm strip of magnesium ribbon.
6. Using scissors, cut the magnesium ribbon into three, 4.0-cm long pieces. Note: Be as precise as possible. The reaction time will depend on the amount of magnesium reacting in each trial.
7. Twist and fold one end of the 20-cm length of copper wire around a pencil to make a small “cage” into which the magnesium ribbon may be inserted. The other end of the wire must be long enough so that the wire will hang over the side of the test tube and the cage will be below the liquid level marked on the test tube (see Figure 2).
   {12638_Procedure_Figure_2}
8. Measure and record the temperature of the room temperature water bath.
9. Fit one piece of magnesium ribbon loosely through a copper wire cage so the magnesium will be held in place but not wrapped around too tightly.
10. Suspend the copper wire cage and the piece of magnesium in the hydrochloric acid solution in the room temperature test tube and immediately start timing.
11. Measure the time until the metal has disappeared and the solution stops bubbling. Record the reaction time in the data table.
12. Remove the copper wire cage from the test tube and rinse with a small amount of distilled or deionized water. Pat dry with a paper towel.
13. Repeat steps 8–12 with each test tube in the ice-water and hot-water baths.

Activity D. The Nature of the Reactants

1. Place the two test tubes in the test tube rack.
2. Using a graduated cylinder, add 18 mL of 2 M hydrochloric acid solution to the first test tube.
3. Add 18 mL of 2 M acetic acid solution to the second test tube.
4. Mass 0.2 g of granular zinc in each of two weighing dishes.
5. Add the 0.2 g of zinc simultaneously to each of the test tubes.
6. Record your observation of the reaction rates and answer the questions in the Observations and Analysis section.

Activity E. Catalysts Examples of Catalysts

1. Place three test tubes in the test tube rack.
2. Use the graduated cylinder to add 5 mL of 3% hydrogen peroxide solution to each test tube.
3. In three small weighing dishes, add a pinch (spatula tip amount) of manganese dioxide (MnO₂) to one, sodium iodide (NaI) to another and dry active yeast to the last weighing dish.
4. Add the solids to separate test tubes of hydrogen peroxide.
5. Observe any signs of a chemical reaction in each test tube and record all observations. Be as descriptive and specific as possible.
6. Light a wooden splint, blow out the flame and then place the glowing end just inside the mouth of the first test tube.
7. Repeat step 6 for each test tube.
8. Record all observations.

Student Worksheet PDF

12638_Student1.pdf