Materials Included In Kit

Additional Materials Required

Safety Precautions

Hydrochloric acid solution is moderately toxic by ingestion and inhalation. It is corrosive to eyes and skin. Sodium thiosulfite is a body tissue irritant. The sulfur produced in this reaction has low toxicity, but may be a skin and mucous membrane irritant. Aqueous sulfur dioxide is generated in this reaction, which is a skin and eye irritant. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please consult current Safety Data Sheets for additional safety, handling and disposal information.

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. Have students empty their six-well reaction plates into one large collection container. Filter the collected solution and dispose of the solid in a landfill according to Flinn Suggested Disposal Method #26a. Neutralize and dispose of the filtrate according to Flinn Suggested Disposal Method #24b.

Teacher Tips

• Make sure that the students empty the well plates and clean them thoroughly using cotton swabs or paper towels immediately after they have finished with each data set. If the plates are allowed to sit with the sulfur precipitate in them, the precipitate will begin to deposit on the bottom of the well. It will then be difficult to thoroughly clean the wells.
• It is important to add the contents of the second syringe-full of solution with a reasonable amount of force so that good mixing occurs between the solutions in the well. But, if the syringe is emptied with too much force, the solutions will spill out of the desired well. Therefore, urge students to find an emptying rate that ensures good mixing without spillage.
• As the beakers containing the solutions become emptied, it will be harder to fill the syringes without introducing air bubbles into them. If this occurs, have students tilt the beakers so that the liquid falls together and can be drawn up into the syringe without air bubbles.
• Make sure that the students empty the well plates and clean them thoroughly using cotton swabs or paper towels immediately after they have finished with each data set. If the plates are allowed to sit with the sulfur precipitate in them, the precipitate will begin to deposit on the bottom of the well. It will then be difficult to thoroughly clean the wells.
• It is important to add the contents of the second syringe-full of solution with a reasonable amount of force so that good mixing occurs between the solutions in the well. But, if the syringe is emptied with too much force, the solutions will spill out of the desired well. Therefore, urge students to find an emptying rate that ensures good mixing without spillage.
• As the beakers containing the solutions become emptied, it will be harder to fill the syringes without introducing air bubbles into them. If this occurs, have students tilt the beakers so that the liquid falls together and can be drawn up into the syringe without air bubbles.

Answers to Prelab Questions

Read the Background section of this lab, then answer the following questions using the information given. Show all work.

Hypothetical Reaction: 2A + B → C + D
General Rate Law: Rate = k [A]ⁿ[B]^m

1. Calculate the reaction rate in sec^–1 for each experiment.

   Each of the reaction times must be inverted to obtain the reaction rates. Values are listed in the table.

2. Determine the order with respect to A.

   Compare the concentrations of A and the reaction rates for experiments 1, 2 and 3. Comparing experiments 1 and 2, as [A] is doubled, the rate doubles also. Comparing experiments 1 and 3, as [A] is tripled, the rate triples also. Therefore, the rate is directly proportional to [A], and hence, the order of reactant A is 1.

3. Determine the order with respect to B.

   To determine the order with respect to reactant B, compare experiments 2, 4 and 5. Comparing experiments 2 and 4, as [B] is doubled, the rate increases by a factor of 4, or 22. Comparing experiments 2 and 5, as [B] is tripled, the rate increases by a factor of 9, or 32. Therefore, the order for reactant B must be 2 to satisfy the proportionality relationship between [B] and the rate.

4. Calculate the overall order of the reaction.

   The overall reaction order is calculated by adding the orders of the individual reactants. Hence, the overall reaction order is 1 + 2 = 3.

5. Determine the rate law.

   Plugging the orders into the general rate law given in the problem, the rate law for this hypothetical reaction is: Rate = k [A][B]².

6. Calculate the units on the rate constant, k.

   The units on the rate are sec^–1, the units on [A] are M, and the units on [B]² are M². Therefore, the units on k must be M^–3 sec^–1.

Sample Data

Date Table 1

Data Table 2

Answers to Questions

1. Calculate the concentrations of the HCl solutions before mixing them with the Na₂S₂O₃ solution in Part A. Remember tocalculate the concentration of the HCl solutions that were diluted by mixing water and HCl in the syringe using the equation M₁V₁ = M₂V₂ where M is molarity and V is volume. Record concentrations in the column labeled “[HCl] initial concentration” in Data Table 1. Show at least one sample calculation below. Record the concentration of the Na₂S₂O₃ solution used in Part A in the column labeled “[Na₂S₂O₃] initial concentration” in Data Table 1. Repeat for Part B and Data Table 2.

   See data tables in Sample Data. The concentrations are determined using the formula M₁V₁ = M₂V₂. For example, for well 3 in Data Table 1, the HCl initial concentration is calculated as follows:

   {11828_Answers_Equation_1}
2. Calculate the concentrations of the HCl solutions and Na₂S₂O₃ solutions after mixing them together for each well. Record these concentrations in the columns labeled “final concentration” in both data tables. Show at least one sample calculation below. To calculate these concentrations, remember that mixing the solutions dilutes each of them. Use the formula M₁V₁ = M₂V₂ to calculate the final concentrations after dilution.

   See data tables in Sample Data. The concentrations are determined using the formula M₁V₁ = M₂V₂, where the final volume is 5 mL in each case. For example, in well 2 in Data Table 1, the calculation of the final concentration of HCl is as follows:

   {11828_Answers_Equation_4}
3. Calculate and record the reaction rate in sec^–1 for each well in the data tables.

   See data tables in Sample Data. The reaction rate in sec^–1 is the inverse of the average reaction time in sec.

4. Based on your data in Data Table 1, does the reaction rate depend on the HCl concentration? If so, how? What is the reaction order for HCl?

   As the HCl concentration is varied the reaction rate does not vary to any significant extent (allowing for experimental error). Therefore, the reaction order for HCl is zero.

5. Based on your data in Data Table 2, does the reaction rate depend on the Na₂S₂O₃ concentration? If so, how? What is the reaction order for Na₂S₂O₃?

   Comparing wells 1 and 2, as the Na₂S₂O₃ concentration is halved, the reaction rate decreases by a factor of 2. Comparing wells 1 and 3, as the Na₂S₂O₃ concentration is decreased by a factor of 3, the reaction rate also decreases by a factor of 3. Therefore, the reaction order for Na₂S₂O₃ must be first order.

6. Plot a graph showing the reaction rate vs. final [HCl] using the data from Data Table 1. What is the shape of the resulting plot?
   {11828_Answers_Figure_1}

   The graph for HCl is approximately a straight line with a slope of approximately zero.

7. On the same graph, plot the reaction rate vs. final [Na₂S₂O₃] using the data from Data Table 2. What is the shape of the resulting plot?

   The graph for Na₂S₂O₃ is a straight line with a positive slope. The variables are linearly related and directly proportional.

8. What is the overall order of the reaction?

   The overall reaction order is first order. This is determined by adding the reaction orders of each of the reactants (1 + 0 = 1).

9. Write the rate law for this reaction.

   Rate = k [Na₂S₂O₃].

10. Based on the rate law you determined, what are the units on the rate constant, k?

    The units on the rate are sec^–1 and the units on [Na₂S₂O₃] are M. Therefore, the rate constant must have units of (sec^–1 M^–1).

References

Flinn Publication No. 13154, The Aloha Chemical Sunset; Flinn Scientific: Batavia, IL, 2005.

Introduction

In this laboratory activity, the effect of reactant concentration on the rate of the reaction between sodium thiosulfate and hydrochloric acid will be studied. The data will then be used to determine the order of each reactant and the rate law for the reaction.

Concepts

• Kinetics
• Rate laws
• Reaction order
• Concentration vs. reaction rate

Background

Reaction Rates
The rate of a chemical reaction is a measure of how fast the reaction occurs. Some chemical reactions occur as soon as the reactants come into contact with each other. Two examples of this type of fast reaction are acid–base reactions or the decomposition of hydrogen peroxide with a catalyst. Other chemical reactions may take years to occur such as the oxidation of iron. Fast or slow processes can be made to occur faster or slower, depending on the reaction conditions. Some of the factors that affect the reaction rate are the nature of the reactants, temperature, surface area, concentration, and the addition of a catalyst.

All rates are measured in terms of the time it takes to complete an event. For example, a car may travel at a rate of 60 miles/hour, a manufacturing plant may produce cars at the rate of 10 cars/day, or a potter may make eight pieces of pottery/week. Each of these rates tells how much time it takes to travel 60 miles, produce 10 cars, or make eight pieces of pottery. In a chemical reaction, the event that is completed is the conversion of reactants to products. A chemical reaction rate is measured in terms of the rate of disappearance of reactants and appearance of products. The reaction is completed when all of one or more reactants have been consumed and converted to products. One way to determine the rate of a chemical reaction is to measure the time from when the reactants are mixed (the start) to when at least one of the reactants has been completely consumed (the end). Some reactions have no convenient visual way to identify the end of the reaction, so advanced instrumentation is needed to monitor the reaction in progress. However, other reactions, such as the one in this laboratory activity, do have convenient ways to identify the end of the reaction such as a color change or precipitate formation.

Reaction Order and Rate Laws
The overall rate of a chemical reaction may depend on the concentrations of one or more of the reactants or it may be independent of the reactant concentrations. Exactly how the rate depends on reactant concentration is expressed in an equation called a rate law. For a general chemical equation, such as

the general rate law would be written as where k is the rate constant, [A] and [B] are the molar concentrations of each of the reactants, and n and m are exponents that determine how the rate depends on the reactant concentrations. The rate constant and the exponents n and m must be determined experimentally—they cannot be determined simply by looking at the balanced chemical equation. The rate constant for a reaction does not depend on the reactant concentrations, but does depend on temperature. The exponents n and m give the order of the reaction. The above reaction is said to be nth order in A and mth order in B. The overall reaction order is obtained by taking the sum n + m. Generally, the exponents n and m are positive whole numbers; however, they may contain fractions or even be negative numbers.

Because the exponents n and m vary from reaction to reaction, rate laws for different reactions take on different forms. The chemical equations and their experimentally determined rate laws listed in Table 1 can be compared to illustrate this fact. From the equations in Table 1, several important points about rate laws can be made.

1. The orders of each reactant in the rate law determine how the rate changes as the concentration of each reactant changes.
   • Zeroth Order. If a reactant has an order of zero, the rate is independent of the concentration of that reactant, and that reactant does not appear in the rate law (because anything raised to the zeroth power is one). For example, in Reaction 3, the order of CO is zero and hence [CO] does not appear in the rate law. Because [CO] does not appear in the rate law, it cannot affect the rate. This means that increasing or decreasing the concentration of a reactant which has an order of zero does not affect the rate of the reaction.
   • First Order. When a reactant is first order, the reactant will appear in the rate law and have an exponent of 1. For example, in Reaction 5, both NO₂ and O₃ have an order of 1. Increasing the concentration of NO₂ or O₃ increases the rate and decreasing the concentration of NO₂ or O₃ decreases the rate. If the concentration of NO₂ is doubled, for instance, the rate will also double. If the concentration of O₃ is halved, the rate will also be cut in half. Hence, the rate is directly proportional to any reactant that is first order.
   • Second Order. For reactants that are second order, they will appear in the rate law as the concentration of the reactant squared. For example, in Reaction 1, [NO] appears in the rate law and is squared. Hence NO has an order of 2. As a result, if the concentration of NO is doubled, this will cause the rate to increase by a factor of 4.
2. The orders n and m in the rate law are not necessarily the same as the coefficients in the balanced chemical equation. For example, in Reaction 2, the exponent for [Br₂] is ½, while its coefficient is 1. Similarly, Reaction 3 is second order in NO₂ and zeroth order in CO, while both of their coefficients in the balanced chemical equation are 1.
3. Reactants can be fractional exponents, even though they are generally positive whole numbers. For example, Reaction 2 is half order in Br₂. This means that if the concentration of Br₂ is doubled, the rate increases by 2^½ or √2.
4. The overall reaction order is the sum of the individual reactant orders. For example, in Reaction 1, the order with respect to NO₂ is 2 and the order with respect to O₂ is 1. Therefore, the overall reaction order is 2 + 1 = 3.
5. The units for the rate in a rate law will always be the same, while the units on the rate constant and the reactant concentrations raised to their powers vary from rate law to rate law.
   • Reaction rates are expressed in sec^–1 (1/sec).
   • The units for each reactant’s concentration depend on the exponent, or order of the reactant. The concentration is always expressed in terms of Molarity, or moles/liter. But, if a reactant has an order of 2 for a given rate law, such as NO₂ in Reaction 3, then that reactant’s concentration will have units of (moles/liter)²
   • The units on the rate constant k must give the correct units for the rate (sec^–1) when multiplied by the reactant concentrations. The units for the rate constant will be the same in reactions with the same overall reaction order, but different between two reactions with different overall reaction orders. To see why, look at Reactions 3 and 5. Both have an overall reaction order of 2 even though they have different forms for the rate law. In Reaction 3, the units for [NO₂]₂ are (moles/liter)². In Reaction 5, the units for [NO₂][O₃] are (moles/liter) x (moles/liter) = (moles/liter)². In each case, because the overall reaction order is 2, the units for the reactant concentration part of the rate law are the same, (moles/liter)². The units for the rate constant for reactions with an overall reaction order of 2 must therefore be liters²/(moles²•sec).

Determining a Rate Law
To determine a rate law for a reaction, the following procedure may be followed. First, one reactant’s concentration is held constant while the second reactant’s concentration is varied. For each reaction, the reaction time is measured and recorded. Then, the first reactant’s concentration is varied while the second reactant’s concentration is held constant. Again, the reaction time for each reaction is measured and recorded. The reaction rate for each reaction is calculated by taking the inverse of the reaction time. The data is then analyzed to determine the order of each reactant and the rate law.

The following steps outline the procedure to estimate the rate law for the reaction between nitric oxide and oxygen gas. The following is a sample set of data that was obtained by performing the reaction five times and varying the concentrations of the reactants as indicated. In each case, the reaction time was measured, then inverted to find the reaction rate. Step 1. Determining the order with respect to O₂. In the first three trials, the concentration of NO is constant while the concentration of O₂ is varied. Therefore, any change in the rate will be solely due to the change in the concentration of O₂. To determine the order with respect to O₂, the [NO]ⁿ part of the rate law can be ignored because it does not vary. The rate constant k can also be ignored because it also does not vary. As a result, the only part of the rate law necessary for determining m is The symbol ∝ means “proportional to.” With this symbol, the goal is to find the factor by which [O₂] is increased and the factor by which the rate is increased from one trial to the next. From the data table we see that as the concentration of O₂ is doubled, the rate doubles also. The factors by which the concentration of O₂ and the rate are increased are then inserted into the proportionality equation. The exponent m is calculated such that the proportionality equation is satisfied. In this case, m must be 1 for the equation to be satisfied. This means that the reaction is first order in O₂.

Step 2. Determining the order with respect to NO. In trials 2, 4 and 5 the concentration of O₂ is held constant while the concentration of NO is varied. Therefore, any change in the rate will be solely due to the change in the NO concentration. In the rate law, [O₂]^m and the rate constant k can be ignored because they do not vary. As a result, the part of the rate law necessary for determining n is Comparing trials 2, 4 and 5, it can be observed that as the concentration of NO is doubled, the rate increases by a factor of 4. Plugging these values into the proportionality equation gives The value of n must therefore be 2 to satisfy the proportionality equation. This means that the reaction is second order in NO.

Step 3. Determining the Rate Law. Plugging the orders for each reactant into the general rate law given in Equation 4, the rate law for the reaction is determined to be:

Experiment Overview

In this laboratory experiment, experimental data will be used to determine the rate law for the reaction between hydrochloric acid, HCl, and sodium thiosulfate, Na₂S₂O₃.

To determine a rate law for this reaction, the following outline will be followed. First, the HCl concentration is held constant while the Na₂S₂O₃ concentration is varied. For each reaction, the reaction time is measured and recorded. Then, the HCl concentration is varied while the Na₂S₂O₃ concentration is held constant. Again, the reaction time for each reaction is measured and recorded. The reaction is timed by monitoring the appearance of the solid black product which is sulfur, S. As solid sulfur is formed, the reaction mixture will become dark and clouded with the black precipitate. The reaction time will be measured by noting the time at which you can no longer see through the solution.

From the general form for a rate law given in equation 2, the general rate law for the reaction between HCl and Na₂S₂O₃ is written as shown in Equation 11. By determining how the reaction rate varies as the concentration of each reactant is varied, the orders n and m, and hence the rate law, will be determined.

Materials

Prelab Questions

See page 1 of the Student PDF.

Safety Precautions

Hydrochloric acid solution is moderately toxic by ingestion and inhalation. It is corrosive to eyes and skin. Sodium thiosulfate is a body tissue irritant. The sulfur produced in this reaction has low toxicity, but may be a skin and mucous membrane irritant. Aqueous sulfur dioxide is generated in this reaction, which is a skin and eye irritant. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Procedure

Prelab Preparation

1. Label one small beaker “HCl.” Add 30 mL of 1.0 M hydrochloric acid, HCl, solution to this beaker.
2. Label another small beaker “Na₂S₂O₃.” Pour 30 mL of 0.30 M sodium thiosulfate, Na₂S₂O₃, solution into this beaker.
3. Label a third small beaker “water.” Pour 25 mL of distilled or deionized water into this beaker.
4. Place a piece of white paper underneath the six-well reaction plate. Now remove the reaction plate and draw six black “+” signs in the spots directly underneath each well. Replace the six-well reaction plate and verify that the “+” signs can be seen through each well.

Part A. Varying the Concentration of HCl

5. Fill the 3-mL syringe up to the 2-mL mark with distilled water by submerging the syringe in the “water” beaker and drawing water into the syringe until the plunger is at the 2-mL mark. Make sure that there are no air bubbles in the syringe. (If the syringe has a tip cover, remove it before filling the syringe.)
6. Now submerge the syringe in the “HCl” beaker and draw 1 mL of HCl solution into the syringe so that the plunger sits at the 3-mL mark.
7. Empty the syringe into the well 3 of the six-well reaction plate (see Figure 1).
   {11828_Procedure_Figure_1}
8. Follow the same filling and emptying procedure as outlined in Steps 5–7, but use the amounts of water and HCl solution indicated below for wells 1 and 2. Fill the wells in reverse order (well 2 next, then well 1). By filling the syringe with the most dilute mixture first and working up to the most concentrated, the syringe does not need to be rinsed between fillings (see Figure 1).
9. Rinse the syringe thoroughly with water. Fill the 3-mL syringe to the 2-mL mark with the Na₂S₂O₃ solution. Prepare to start the timer. Empty the syringe into well 1. Time the reaction with a stopwatch or timer by measuring the time from which the solution was added until the black “+” sign can no longer be seen through the solution. Record the exact time in seconds in Data Table 1.
10. Repeat step 9 for wells 2 and 3, adding 2 mL of the Na₂S₂O₃ solution to each well. Carefully time each reaction with a stopwatch or timer by measuring the time from which the solution was added until the black “+” sign can no longer be seen through the solution. Record the exact time in seconds in Data Table 1.
11. Repeat steps 5–10 in wells 4–6 to obtain a second set of data.
12. Empty the six-well reaction plate into the collection container provided by your teacher. Rinse and dry each of the wells with soap and water. Use a cotton swab or a paper towel to thoroughly clean and dry each well.

Part B. Varying the Concentration of Na₂S₂O₃

13. Fill the 3-mL syringe up to the 2-mL mark with distilled water by submerging the syringe in the “water” beaker and drawing water into the syringe until the bottom of the plunger is at the 2-mL mark. Make sure there are no air bubbles in the syringe.
14. Now submerge the syringe in the “Na₂S₂O₃” beaker and draw 1 mL of the Na₂S₂O₃ solution into the syringe so that the plunger sits at the 3-mL mark.
15. Empty the syringe into well 3 of the six-well reaction plate (see Figure 2).
    {11828_Procedure_Figure_2}
16. Follow the same filling and emptying procedure as outlined in steps 13–15, but use the amounts of water and Na₂S₂O₃ solution indicated below for wells 1 and 2. Fill the Wells in reverse order (well 2 next, then well 1). By filling the syringe with the most dilute mixture first and working up to the most concentrated, the syringe does not need to be rinsed between fillings (see Figure 2).
17. Rinse the syringe thoroughly with water. Fill the 3-mL syringe to the 2-mL mark with the HCl solution. Prepare to start the timer. Empty the syringe into well 1. Time the reaction with a stopwatch or timer by measuring the time from which the solution was added until the black “+” sign can no longer be seen through the solution. Record the exact time in seconds in Data Table 2.
18. Repeat step 17 for wells 2 and 3, adding 2 mL of the HCl solution to each well. Carefully time each reaction with a stopwatch or timer by measuring the time from which the solution was added until the black “+” sign can no longer be seen through the solution. Record the exact time in seconds in Data Table 2.
19. Repeat steps 13–18 to obtain a second set of data.
20. Empty the six-well reaction plate into the collection container provided by your teacher. Rinse and dry each of the wells with soap and water. Use a cotton swab or paper towel to thoroughly clean and dry each well.

Student Worksheet PDF

11828_Student1.pdf