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Rate Laws and Order of Reaction

Read 6.3 p 372-378

Q 1-6 p 377

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Concentration and Rate Each reaction has its own equation that

gives its rate as a function of reactant concentrations.

this is called its Rate Law

• A rate law is an equation expressing the rate of a reaction in terms of the [molar] of the species involved in the reaction.

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aA + bB cC + dD

• The rate is expressible as rate = k[A]m[B]n

• The powers m and n are kinetic orders,• A and B are the chemical substances• k is the rate constant, different values for

each rxn

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Order of a Reaction• The “order” of a chemical reaction is the number

of chemical [] terms upon which the rate depends

• Rate = k[A] - first order rate; the rate only depends on one [] term

• Rate = k[A]2 - second order rate

• Rate = k[A][B] - second order rate

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Concentration and Rate

Compare Experiments 1 and 2:when [NH4

+] doubles, the initial rate doubles.

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Concentration and Rate

Likewise, compare Experiments 5 and 6: when [NO2

-] doubles, the initial rate doubles.

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Concentration and Rate

This equation is called the rate law, and k is the rate constant.

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Rate Laws• Exponents tell the order of the reaction with

respect to each reactant.• This reaction is

First-order to [NH4+]

First-order to [NO2−]

• The overall reaction order can be found by adding the exponents on the reactants in the rate law.

• This reaction is second-order overall.

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Recall………………………..• A rate law shows the relationship between the reaction

rate and the concentrations of reactants. For gas-phase reactants use Partial Pressure instead of [A].

• k is a constant that has a specific value for each reaction.• The value of k is determined experimentally.

“Constant” is relative here- k is unique for each rxnk changes with T

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Temperature and Rate

• Generally, as temperature increases, so does the reaction rate.

• This is because k is temperature dependent.

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Determining Rxn Order from Experimental Data – Ex.1

• Four experiments were conducted to discover how the initial rate of consumption of BrO3

- ions in the rxn varies as the concentrations of the reactants are changed.

• Use the experimental data in the table below to determine the order of the reaction with respect to each reactant and the overall order.

• Write the rate law for the rxn and determing the value of k.

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Data – Ex. 1

Initial [ ]mol L-1

Initial Rate

Experiment BrO3- Br- H+ mol L-1 s-1

1 0.10 0.10 0.10 1.2x10-3

2 0.20 0.10 0.10 2.4x10-3

3 0.10 0.30 0.10 3.5x10-3

4 0.20 0.10 0.15 5.4x10-3

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Homework

• Pg. 377 # 1-6

• Lab Exercise 6.4.1 The Sulfur Clock “determine the order of the reaction” p.404-405

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Review of Chemical KineticsFirst order Second order Second order

Rate Laws

Graph of rate vs concn

Linear Exponential Exponential

-dependent on one molecule

-dependent on two molecules

-dependent on two molecules

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First-Order ProcessesConsider the process in which methyl isonitrile is converted to acetonitrile.

CH3NC CH3CN

How do we know this is a first order rxn?

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Reaction MechanismsRead 6.4 p 383-390

The sequence of events that describes the actual process by which reactants become products is called the reaction mechanism.

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Reaction Mechanisms

• Reactions may occur all at once or through several discrete steps.

• Each of these processes is known as an elementary reaction or elementary process.

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Reaction Mechanisms

• The molecularity of a process tells how many molecules are involved in the process.

• The rate law for an elementary step is written directly from that step.

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Multistep Mechanisms

• In a multistep process, one of the steps will be slower than all others.

• The overall reaction cannot occur faster than this slowest, rate-determining step.

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Slow Initial Step

• The rate law for this reaction is found experimentally to be

Rate = k [NO2]2

• CO is necessary for this reaction to occur, but the rate of the reaction does not depend on its concentration.

• This suggests the reaction occurs in two steps.

NO2 (g) + CO (g) NO (g) + CO2 (g)

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Slow Initial Step• A proposed mechanism for this reaction is

Step 1: NO2 + NO2 NO3 + NO (slow)

Step 2: NO3 + CO NO2 + CO2 (fast)

• The NO3 intermediate is consumed in the second step.

• As CO is not involved in the slow, rate-determining step, it does not appear in the rate law.

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Fast Initial Step

• The rate law for this reaction is found (experimentally) to be

• Because termolecular (= trimolecular) processes are rare, this rate law suggests a two-step mechanism.

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Fast Initial Step• A proposed mechanism is

{Step 1 is an equilibrium- it includes the forward and reverse reactions.}

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Fast Initial Step

• [NO] [Br2] = [NOBr2]

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Questions

• Q 2 p 390• Q 1-3 p 391

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Unit Review

• Chapter 5: Self Quiz Q 1-18 p 355 Q 2,4,6,9,14,16

• Chapter 6: Self Quiz 1-18 p 407, Q 4,7,910,13-16 p 408-409