chapter 12 chemical kinetics.notebook · 2016. 9. 21. · chapter 12 chemical kinetics.notebook 3...

Chapter 12 Chemical Kinetics.notebook

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12.1 Reaction RatesWhen you finish this section you will be able to use concentration vs. time data to determine relative reaction rates.

Chemical Kinetics the area in chemistry that concerns reaction rates

The rate of a reaction is defined as:

Rate = change in concentration of a substancechange in time

Rate of Rate of 2(Rate of consumption = production = production

of NO2 of NO of O2)

Δ[NO2] = Δ[NO] = 2(Δ[O2]) Δt Δt Δt

Starting with a flask of nitrogen dioxide at 300 degrees C, concentrations of nitrogen dioxide, nitric oxide, and oxygen are plotted versus time

2NO2(g) → 2NO(g) + O2(g)

Notice that the rate of NO2 decomposition decreases with time until equilibrium is reached. The rate of a reaction is not constant.


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Let’s look at data for the reaction of hydrogen and oxygen to give water,2H2(g) + O2(g) → 2H2O(l)

The rate of disappearance of H2 must be twice the rate for O2, and equal to the appearance of H2O. This is because the coefficients of the chemical equation tell you the relative rates. That is:

Δ[H2] = 2Δ[O2] = Δ[H2O]Δt Δt Δt

Using the data from our reaction, the rate of reaction of hydrogen between 0.004 and 0.008 seconds is:

rate = Δ[H2] = (0.01250.025)= (0.0125) Δt (0.0080.004) 0.004

where “[ ]” means concentration in moles/liters This says that for every mole of O2 that reacts, per unit time, 2 moles of H2 react, and 2 moles of H2O are produced. Note that the rate of disappearance of the reactants is given a negative (”“) sig. That implies that we are losing reactants while gaining products (“+” rate). Note that the standard is, that all reaction rates should be positive.

We are losing hydrogen at a rate of 3.1 mol/L s.

Apr 810:28 PM

Example 12.1A Reaction RatesAnswer the following questions using the hydrogen and oxygen rate data just presented.a. Based on the coefficients of the chemical equation alone, what is the rate of oxygen reaction between 0.004 and 0.008 seconds?b. What is the rate of water production during the same time period?c. Use the concentration vs. time data to prove your answers.

Example 12.1B Practice with Reaction RatesGiven the hypothetical reaction:

2A + nB → qC + rD

If Δ[A] = 0.050 mol/L s Δt

Δ[B] = 0.150 mol/L s Δt

Δ[C] = 0.075 mol/L s Δt

Δ[D] = 0.025 mol/L s Δt

What are the coefficients n, q, and r?


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Examples:1. What is the average rate of the reaction between 10 s and 20 s based on the table of data shown below. A à B

time (s) [A] (mol/L) 0.0 0.200 5.0 0.140 10.0 0.100 15.0 0.071 20.0 0.050

2. For the reaction, 2N2O5(g) à 4NO2(g) + O2(g), the rate of formation of NO2(g) is 4.0 x 103 mol L1s1. (a) Calculate the rate of disappearance of N2O5(g)

(b) Calculate the rate of appearance of O2(g).

3. Consider the following reaction: 3A à 2B

The average rate of appearance of B is given by Δ[B] / Δ t. How is the average rate of appearance of B related to the average rate of disappearance of A?

For the reaction: 2SO2(g) à 2SO(g) + O2(g), the disappearance of SO2(g) was measured to monitor the progress of the reaction.

Time (s) [SO2] (mol/L) 0.00 0.526 10.0 0.418 20.0 0.331 30.0 0.263 40.0 0.209 50.0 0.166 60.0 0.132

4. Calculate the average rate of disappearance of SO2 between 10.0 and 30.0 seconds.


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Examples5. Given the data below for the reaction, yY + zZ à products

Experiment Initial Number [Y]/M [Z]/M Rate (M/s) 1 0.100 0.100 4.0 x 105 2 0.200 0.100 1.6 x 104 3 0.100 0.200 8.0 x 105

What is the rate Law?

What is order of the reaction overall?

Calculate the rate constant? (Include units)

6. The reaction 2NO(g) + 2H2(g) à N2(g) + 2H2O(g) is found experimentally to be second order in NO(g) and firstorder in H2(g).

(a) Write the rate law for the reaction.

(b) What is the overall order of the reaction?

(c) What are the units for the rate constant k?

(d) If [NO] is doubled (while keeping [H2] constant), by what factor will the reaction rate increase?

(e) If [H2] is doubled (while keeping [NO] constant), by what factor will the reaction rate increase?

Apr 89:10 PM

12.2 Rate Laws: An IntroductionAnswer the following review questions after you have read p.5324. 1. What simplifying assumption do we make regarding the rates forward and reverse reactions in our study of rate laws?2. Define rate law.3. What are “k” and “n”?4. Why don’t products appear in the rate law?5. Why is it important to specify the component whose rate we are describing with the rate law?6. Give a couple of practical examples of the importance of determining rate laws.

Example 12.3B Rate Law DeterminationDetermine the rate law, and solve for the order and value of the rate constant for the reaction.

C + D + E → Productsgiven the following data


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12.6 Reaction MechanismsWhen you finish this section you will be able to evaluate the validity of simple reactions mechanisms.

• Reaction mechanisms – the series of steps by which a chemical reaction occurs.

Consider the reaction: NO2(g) + CO(g) → NO(g) + CO2(g)

The mechanism involves the following steps:

NO2(g) + NO2(g) → NO3(g) + NO(g)

NO3(g) + CO(g) → NO2(g) + CO2(g)

A balanced reaction does not tell us how the reactants become products.

k1

k2

NO3 is an intermediate – a species that is neither a reactant nor product. It is formed and consumed in the reaction sequence and so never seen as a product.

Molecularity – the number of species that collide to produce the reaction indicated by an elementary step. Uni and bimolecular are two important molecularities. (why isn’t a termolecular molecularity likely?)

Apr 168:53 AM

Ratedetermining step – the slowest step. The reaction can only go as fast as the ratedetermining step.

NO2(g) + NO2(g) → NO3(g) + NO(g)

NO3(g) + CO(g) → NO2(g) + CO2(g)

Slow (rate determining)

Fast

Main ideas:• a chemical equation only tells us what reactants become what products. It does not tell us how these changes occur. This is the goal of reaction mechanism studies.• the rate law is determined by the ratedetermining step. This step may be in one direction or reversible (both directions). Your goal is to be able to determine whether mechanisms are consistent with experimentally determined rate laws.

Review Sample 12.6 in your textbook, then try the following examples.

Reaction mechanism is a series of elementary steps that must satisfy two requirements:1. The sum of the elementary steps must give the overall balance equation for the reaction.2. The mechanism must agree with the experimentally determined rate law.

NO + H2 → N + H2O (slow)N + NO → N2O (fast)N2O + H2 → N2 + H2O (fast)

Is this mechanism consistent with the observed rate law?

The balanced equation for the reaction of nitric oxide with hydrogen is:

2NO + 2H2 → 2H2O + N2

The experimentallydetermined rate law is: Rate = k[NO]2[H2]The following mechanism has been proposed

Example 12.6A Reaction Mechanisms

k3k2k1


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12.7 A model for Chemical KineticsWhen you finish this section you will be able to state the criteria for reaction according to the collision model.

(a) The change in potential as a function of reaction progress

(b) A molecular representation of the reaction

There are two empirical facts about the rates of chemical reactions. All other things being equal.1. The more concentrated your reactants, the faster the reaction.2. The higher the temperature, the faster the reaction.

Apr 1610:39 AM

Collision model molecules must collide to react.The collision theory says that there are four steps involved in having a successful reaction.

1. Molecules must move toward each other (through random motion)2. Molecules must it with the proper orientation (see Figure below).3. Molecules must hit with sufficiently high energy.4. Molecules will separate after reaction occurs.

Several possible orientations for a collision between two BrNO molecules. (a) and (b) can lead to a reaction, but orientation (c) can not.


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12.8 CatalysisThe following questions will help you review the material in this section1. Define catalysis.2. What are enzymes?3. Why are catalysts more useful for increasing the reaction rate than raising the temperature?4. How do catalysts work? How does this reflect the collision theory model?5. What is a heterogeneous catalyst? Give some examples6. Discuss the role of catalysis in the conversion of sulfur dioxide to acid rain.7. What is a homogeneous catalyst? Give an example.8. How does nitric oxide act as a catalyst in the production of ozone?9. How can chlorine catalyze the decomposition of ozone?

Apr 231:07 PM


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The overall rate of the reaction is determined, for the most part, by the transition state of highest energy in the pathway. This transition state, which is usually the slowest step, controls the rate of reaction and is thus called the ratedetermining step of the mechanism.

Most reactions require the addition of energy. Energy is needed for molecules to pass over the energy barriers that separate them from becoming reaction products. These energy barriers are called the activation energy, or enthalpy of activation, of the reactions. At room temperature, most molecules have insufficient kinetic energy to overcome the activation energy barrier so a reaction can occur. The average kinetic energy of molecules can be increased by increasing their temperature. The higher the temperature, the greater the fraction of reactant molecules that have sufficient energy to pass over the activation energy barrier. Thus, the rate of a reaction increases with increasing temperature.The rate of a reaction also depends on the number of interactions between reactant molecules. Interactions increase in solutions of greater concentrations of reactants, so a reaction rate is directly proportional to the concentration of the reactants. The proportionality constant is called the rate constant for the reaction. Not every collision is effective in producing bond breakage and formation. For a collision to be effective, the molecules must have sufficient energy content as well as proper alignment. If all collisions were effective, every reaction would proceed with explosive force.

chapter 12 chemical kinetics.notebook · 2016. 9. 21. · chapter 12 chemical kinetics.notebook 3...

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