Exploration of Reaction Kinetics

Organization

• •
  Mode: laboratory work, work in pairs
• •
  Grading: lab notes, lab performance, and post-lab report
• •
  Safety: goggles, closed-toe shoes, long pants/skirt, and sleeves; latex or nitrile gloves recommended

Chemical Classification and Potential Hazards of Chemicals Used in Experiment

CHEM 115 Expt. 12	Chemical Classification													Possibility Of:					NFPA Codes
Reaction Kinetics	Poison A	Flammable Gas	Flammable Liquid	Combustable Liquid	Reacts with Water	Oxidizer	Organic Peroxide	Poison B	Corrosive Acid	Corrosive Base	Irritating or Harmful	Misc. Hazard	No Hazard	Fire	Sudden Release of Pressure	Reactive	Immediate (Acute) Health Hazard	Delayed (Chronic) Health Hazard	Fire	Health	Reactivity	Special Precautions
Oxalic Acid, 0.20 M								X			X						X		0	2	1
Potassium Permanganate, 0.0015 M						X													0	2	2
Sulfuric Acid, 2.0 M									X		X					X			0	2	1	COR

I: Background

In order for a reaction to take place, the molecules, atoms, or ions involved in that reaction must collide in the proper orientation and with sufficient energy. Only a small fraction of the collisions that take place result in a reaction occurring. The rate of a reaction is therefore directly influenced by the number of collisions that occur, and the energy of those collisions. By manipulating either the collision frequency or the collision energy, the rate of a chemical reaction can be manipulated. In the previous exercise, you manipulated the rate of a reaction by changing the concentration of reactant. Would this change have an effect on collision energy or collision frequency? What very simple thing can be done in a laboratory to manipulate the energy of chemicals in a reaction flask? Numerous factors influence the rates of chemical reactions. The most fundamental factor is the nature of the reactants and products. Some substances are extremely unstable, spontaneously decaying in microseconds or less. Others are extremely stable, and will react only under extreme conditions, and then very slowly. For a particular chemical reaction, many factors can influence the rate of reaction. Some of the factors that often, but not always, influence the rate of a reaction are as follows.

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  the concentrations of the reactants and products
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  the concentrations of spectator ions (ionic strength of the solution)
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  the concentration or surface area of a catalyst
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  the pressure of the reaction vessel
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  the temperature at which the reaction is performed

In this inquiry, you will investigate the influence of some of these factors. The reduction of permanganate ion by oxalic acid in acidic solution will be studied. The overall reaction can be written as follows. While the stoichiometry of this reaction suggest the involvement of 13 species reacting, in reality, it is highly unlikely that 13 species will collide in solution in the proper orientation and instantly be converted to 26 new species. In reality, this net reaction occurs as a series of small chemical steps involving 1, 2, or at most, 3 species at a time. Each of these small chemical steps is called an elementary reaction, and the series of elementary reactions by which reactants are converted to products is called the reaction mechanism. While the goals of this experiment are simply to explore the factors that influence the rate of this redox reaction, kinetics investigations similar to those performed here are used to gain a detailed understanding of the actual, step-by-step mechanism of the reaction.

Reaction Mechanisms

As an analogy to the concept of a mechanism in chemistry, imagine buying a bookcase that requires assembly (Figure 1). The overall reaction that takes place is the conversion of 3 shelves and 2 sides into one bookcase, but that reaction doesn't take place in one step. Instead, the assembly proceeds one step at a time using only two or three pieces for each step, and the steps take place in a specific order. The assembly instructions provide the individual steps, or elementary reactions, that are required to convert the pieces of the bookcase into the finished bookcase. Partially assembled bits of the bookcase made after one step act as intermediates that are needed for a later step in the mechanism. Rates of reactions occurring in solution are typically reported with the intensive unit, M · s^–1; that is, moles per liter per second. The rate can be defined in terms of the disappearance of a reactant, or the appearance of a product. Because the permanganate ion has an intense color, and all other reactants and products are colorless, you will monitor the rate of the reaction by observing the time required for a particular concentration of permanganate ion to decrease to zero. For most of your experiments, the initial concentration of permanganate ion is the same, so the relative rates of reaction can be determined simply by comparing the time required for all of the permanganate ion to react.

Activation Energy

In every chemical reaction, the reactants must overcome an energy barrier in order to react. This energy barrier is called the activation energy and is abbreviated E_a. Shown in Figure 2 are two reaction progress diagrams. The energy of the chemical species as the reaction progresses is traced on the y-axis. The reactants must gain energy to reach the transition state, after which it is possible for the reactants to be converted to products. Both the activation energy and the net change in energy can be deduced from Figure 2. Imagine bowling or playing pool. To knock down the most pins, or to have a powerful break, you want to hit the pins or the rack with a lot of energy and speed. This is similar to the effect that temperature has on the energy of collisions between molecules. The energy of the molecules increases with increasing temperature (and decreases with decreasing temperature). At higher temperatures, molecules collide with much more energy than at lower temperatures, and a greater proportion of the collisions take place with energy equal to or greater than the activation energy. This means that more reactants collide with enough energy to become products, resulting in an increase in the rate of the reaction. As a matter of fact, an increase in temperature of only 10°, from 25°C to 35°C, will double the rate of the reaction.

Effect of Temperature on Kinetics

While an increase in collision frequency also results in higher temperature, it is not the number of collisions that determines the rate of the reaction—it is the energy of the collision that determines the rate of the reaction. You can hit a pool ball softly over and over and over again, requiring many collisions before you sink it, or you can hit it hard once and sink it. It is the energy of the collision that matters most, and temperature is the way to control the energy of the collision. As the temperature increases, the number of collisions occurring with energy equal to E_a increases exponentially, and that is why the rate of the reaction increases with increasing temperature.

Mechanisms and Rate-Limiting Steps

Remember the bookcase building mechanism at the start of this exercise? Each one of the individual steps in that mechanism is associated with its own rate, its own reaction progress diagram, and its own E_a. In a chemical reaction, each step in a mechanism must occur in the proper sequence, because the products of an earlier step will be an intermediate that is needed in a later step. Some of the steps in the mechanism are faster than the others, but one of the steps in the mechanism will be the slowest of all of the steps. The rate of the overall reaction can be no faster than the rate of the rate-limiting step, because the slowest step has to happen before the rest of the steps can happen. For the multi-step reaction mechanism that must exist for the reaction studied in this exercise, the observed rate of the reaction will depend only on the reactant(s) involved in the rate-limiting step. In this exercise, the experimental data suggests that the rate-limiting step of this reaction involves H₃O⁺ but not oxalic acid. If you were a chemist investigating the mechanism of this reaction, you would have valuable information following these two experiments; you would know there were at least two steps in the mechanism—the rate-limiting step involving H₃O⁺, and a faster step involving oxalic acid.

II: Exercises

Part A: Getting Ready to Perform the Experiments

Part B: Exploring the Redox Reaction

Exploring Ways to Increase the Rate of a Chemical Reaction

General Experimental Notes The rest of the lab involves exploring different experimental parameters on the rate of the reaction explored in Part B. For each of the following experiments, you will use three 4-inch test tubes that have been cleaned, distilled-water rinsed, dried, and labeled. Support the test tubes in a test tube rack, and place a piece of white paper underneath and behind the test tubes. Unlike Part B, you must carefully record the time it takes for the entire permanganate ion sample to be reduced by the oxalic acid. The recommended procedure is as follows.

Part C: Exploring the Effect of the Concentration of H₃O⁺(aq)

• a.
  Is H₃O⁺ a reactant or a product in the reaction? What effect did increasing the amount of H₃O⁺ have on the rate of the reaction?
• b.
  Which test tube had the slowest reaction rate, and which had the fastest reaction rate? How does the reaction time relate to rate of reaction—do long reaction times reflect slow reaction rates or fast reaction rates?

• c.
  By what factor did the concentration of H₃O⁺ change from C–1 to C–2. By what factor did the rate of the reaction change from C–1 to C–2?
• d.
  If the rate of a reaction changes by a substantial amount (factor of 2 or more) when the concentration of a reactant changes, then the rate of the reaction is considered to be dependent on the concentration of that reactant. Would you say that the rate of the reaction depends on the concentration of H₃O⁺ in solution?
• e.
  What are the molecules and ions doing in solution—are they static or moving around?
• f.
  Can ions or molecules that are isolated from one another react? What must happen for them to react?
• g.
  Why should the permanganate disappear faster if the concentration of H₃O⁺(aq) is higher?
• h.
  Devise a general, qualitative statement about the effect that an increase in the concentration of a reactant has on the rate of a reaction, and an explanation of that effect.

Part D: Exploring the Effect of Temperature

Part E: Exploring the Effect of Copper Metal

Part F: Rates of Reaction and Activation Energy

In this exercise, you observed the effect of the concentration of H₃O⁺, temperature, and Cu(s) on the rate of reaction. This was a macroscopic observation based on a change in color that occurred in the test tube. This macroscopic observation tells you what happened to the rate of a reaction as the function of changing amounts or reactants, temperature, or adding Cu, but it doesn't tell you how these changes influence the rate of the reaction. To understand how these changes affect rate, you must think at a molecular level, in terms of collisions between the molecules in the test tubes. In the background for this exercise, you were told that for a reaction to take place, the molecules or ions involved in the reaction must collide in the proper orientation and with sufficient energy. The change in rate that you measured was a macroscopic observation. The relationship between the frequency of collisions and concentration of reactants was explored in Part C; and in questions e–h, you thought about how a change in H₃O⁺ affected rate at a molecular level. In this section, you will explore the relationship between the rate of a reaction and the energy of a collision, and the effect a catalyst has on the rate of a reaction. Consider the following figure. The darker curve represents the uncatalyzed reaction; the lighter line represents the catalyzed reaction. It is important to realize that it is not the number of steps, but the E_a of a particular step, that determines the rate. Although the catalyzed pathway involves more steps, each step is much faster than the uncatalyzed pathway, so the reaction will be faster overall in the presence of a catalyst. The catalyst changes the mechanism of the reaction to one with lower energy(ies) of activation, resulting in a faster rate of reaction. Some characteristics of catalysts are that they are not consumed during the course of a reaction, and they do not appear in the overall equation of the reaction being studied.

• 1.
  In Figure 2, which reaction requires more energy to reach the transition state? Recreate the diagrams in your notebook. On each diagram, draw a line indicating where you measured this energy, in order to make the comparison. This energy is called the activation energy, E_a, and it represents the energy barrier between reactants and products.
• 2.
  Which reaction in Figure 2 do you think would be faster—one with a larger energy barrier or one with a smaller energy barrier? Explain why you think so. Given this, which reaction profile corresponds to a faster reaction, and which corresponds to a slower reaction?
• 3.
  For a reaction to take place, the species involved must collide with sufficient energy. What is the minimum amount of energy needed for a collision to result in a reaction?
• 4.
  Devise a general statement about the relationship between E_a and the rate of a reaction.
• 5.
  Describe the effect of temperature on the energy of collisions, and explain how this effect changes the rate of a reaction.

In Figure 3, the path of the reaction in the absence of a catalyst is shown as a dark line, and the path in the presence of a catalyst is shown as a lighter line.

• 6.
  Describe the differences in appearance between the catalyzed and uncatalyzed pathways in Figure 3, including relative energies and number of energy barriers.
• 7.
  Which reaction will be faster—the catalyzed pathway or the uncatalyzed pathway? Why?
• 8.
  Of the variables studied during this experiment, H₃O⁺(aq), C₂H₄O₂(aq), Cu(s), and temperature, which was the catalyst? Why do you think so?
• 9.
  Describe the effect of a catalyst on the rate of a reaction, and explain how a catalyst changes the rate of a reaction.

Copyright © 2016 Advanced Instructional Systems Inc. dba WebAssign, a Virginia Corporation and Cengage Company, and San Francisco State University Department of Chemistry and Biochemistry Faculty, Fall 2016 Update by Nancy Gerber | Credits