Chemical Kinetics In kinetics we study the rate at which a chemical process occurs. Besides information about the speed at which reactions occur, kinetics

ChemicalKinetics

Kinetics

• In kinetics we study the rate at which a chemical process occurs.

• Besides information about the speed at which reactions occur, kinetics also sheds light on the reaction mechanism (exactly how the reaction occurs).

© 2012 Pearson Education, Inc.

ChemicalKinetics

The Collision Theory• In a chemical reaction, bonds are

broken(endothermic) and new bonds are formed (exothermic).

• There are three conditions that need to be met in order for a reaction to take place


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The Three Conditions of the Collision Theory

• 1. The reactants need to come into contact (collide)


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2. The collision must occur with a certain minimum energy to form an activated complex (activation energy Eact)

• Must have enough energy to overcome the electron/electron repulsion of the valence shell electrons of the reacting species


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Activation Energy• In other words, there is a minimum amount of energy

required for reaction: the activation energy, Ea.

• Just as a ball cannot get over a hill if it does not roll up the hill with enough energy, a reaction cannot occur unless the molecules possess sufficient energy to get over the activation-energy barrier.


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Reaction Coordinate Diagrams• The diagram shows the

energy of the reactants and products (and, therefore, E).

• The high point on the diagram is the transition state.

• The species present at the transition state is called the activated complex.

• The energy gap between the reactants and the activated complex is the activation-energy barrier.


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• 3. The molecules that are colliding must have the correct orientation


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Proper orientation

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The Collision Model

Furthermore, molecules must collide with the correct orientation and with enough energy to cause bond breakage and formation.


ChemicalKinetics

The Collision Model

Furthermore, molecules must collide with the correct orientation and with enough energy to cause bond breakage and formation.


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5 Factors that affect Reaction Rate

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5 Factors That Affect Reaction Rates

• 1. Concentration of reactants.– As the

concentration of reactants increases, so does the likelihood that reactant molecules will collide.


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2. Nature of the reactants– Some reactant molecules react in a hurry, others

very slowly.– Physical state– gasoline (l) vs gasoline (g) or

K2SO4 (s) + Ba(NO3)(S) no rxn, but they will react when aqueous

– Chemical identity: What is reacting? Oppositely charged ions react very rapidly. Or metallic sodium reacts more rapidly than metallic magnesium. Or nature of bonds in reactants


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3. Temperature– At higher temperatures, reactant

molecules have more kinetic energy, move faster, and collide more often and with greater energy.

– As a guideline in many reactions a 10 C rise in temp will double reaction rate


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Maxwell–Boltzmann Distributions

• Temperature is defined as a measure of the average kinetic energy of the molecules in a sample.

• At any temperature there is a wide distribution of kinetic energies.


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Maxwell–Boltzmann Distributions

• As the temperature increases, the curve flattens and broadens.

• Thus, at higher temperatures, a larger population of molecules has higher energy., and the particles have the activation energy on collision

• This allows for a greater number of successful collisions.


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4. Presence of a catalyst.– Catalysts speed up reactions by changing

the pathway (place reactants in proper orientation—”plays matchmaker”) of the reaction.

– Catalysts are not consumed during the course of the reaction.

– lowers the activation energy of a reaction


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Catalysts• Catalysts increase the rate of a reaction by

decreasing the activation energy of the reaction.• Catalysts change the mechanism by which the

process occurs.• The H Remains unchanged


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Catalysts

One way a catalyst can speed up a reaction is by holding the reactants together and helping bonds to break.


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• 5. Surface Area of the reactants- Greater surface area, increases the chance of collision—more speed

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Reaction Rates

Rates of reactions can be determined by monitoring the change in concentration of either reactants or products as a function of time.


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• Rate= change in concentration of a species

• Time interval

[A]

t

• Rate of reactant is always negative

• Rate of product is always positive

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Reaction Rates

In this reaction, the concentration of butyl chloride, C4H9Cl, was measured at various times.

The speed of a reaction is expressed in terms of “its rate”, some measurable quantity changes with time


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Reaction Rates

The average rate of the reaction over each interval is the change in concentration [Molarity] divided by the change in time:

C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq)

Average rate =[C4H9Cl]

t


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Reaction Rates C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq)

Is average rate constant??:


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Reaction Rates

• Note that the average rate decreases as the reaction proceeds.

• This is because as the reaction goes forward, there are fewer collisions between reactant molecules.



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Reaction Rates

• A plot of [C4H9Cl] versus time for this reaction yields a curve like this.

• The slope of a line tangent to the curve at any point is the instantaneous rate at that time.



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Reaction Rates

• All reactions slow down over time.

• Therefore, the best indicator of the rate of a reaction is the instantaneous rate near the beginning of the reaction.



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Reaction Rates

• In this reaction, the ratio of C4H9Cl to C4H9OH is 1:1.

• Thus, the rate of disappearance of C4H9Cl is the same as the rate of appearance of C4H9OH.


Rate =[C4H9Cl]

t=

[C4H9OH]t


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• Graphing data of an experiment will show average reaction rate.

• Instantaneous rate = slope (rise/run) of line tangent of to the curve at that point!!

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Example Problem

• Examine the graph above and calculate the AVERAGE rate at which [NO2] changes in the first 50 seconds

• What is the instantaneous rate and 600 seconds?

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Relative Reaction Rates

• We can consider the appearance of product along with the disappearance of reactants. What if the ratio is not 1:1??

• The reactants concentration is declining while the products’ are increasing

2 HI(g) H2(g) + I2(g)


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Reaction Rates and Stoichiometry

• We look at the Relative Rates of Reaction

2 HI(g) H2(g) + I2(g)

= 12

[HI]= 1

1[H2]t

11

=


In such a case

t t

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Relative Reaction Rates and Stoichiometry

• To generalize, then, for the reaction• Make sure reactants have neg sign and watch stoich

aA + bB cC + dD

Rate = 1a

[A]t =

1b

[B]t =

1c

[C]t

1d

[D]t=


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Example:

• What are the relative rates of change in concentration of products and reactant in the decomposition of nitrosyl chloride?

• 2NOCl (g) 2 NO (g) + Cl2 (g)

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Differential Rate Law

• Considers only the relation between the reaction rate and the concentration of reactants given by a mathematical equation

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Differential Rate Law

To find the exact relation between rate and concentration, we must conduct experiments and collect information.


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Concentration and Rate

If we compare Experiments 1 and 2, we see that when [NH4

+] doubles, the initial rate doubles.

NH4+(aq) + NO2

(aq) N2(g) + 2 H2O(l)


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Concentration and Rate

Likewise, when we compare Experiments 5

and 6, we see that when [NO2] doubles, the

initial rate doubles.

NH4+(aq) + NO2

(aq) N2(g) + 2 H2O(l)


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Concentration and Rate• This means

Rate [NH4+]

Rate [NO2]

Rate [NH4+] [NO2

]which, when written as an equation, becomes

Rate = k [NH4+] [NO2

]• This equation is called the rate law, and k is

the rate constant.

Therefore,


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Rate Laws• A rate law shows the relationship between the

reaction rate and the concentrations of reactants.• The exponents tell the order of the reaction with

respect to each reactant.• Since the rate law is

Rate = k[NH4+] [NO2

]

the reaction is

First-order in [NH4+]

and

First-order in [NO2]


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Rate Laws

Rate = k[NH4+] [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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The rate constant k

• Is temperature dependent

• Must be determined by experimentation

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The formula for the differential rate law

C

aA + bB xX

Where C is the catalyst

Initial rxn rate= k[A]m [B]n [C]p

Where m=order of reaction for reactant A,

n=order for reactant B,p=order of catalyst

Exponents can be 0,whole numbers or fractions and must be determined experimentally. The exponent is 0 it has no effect on the rate.

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BrO3-(aq)

+ 5 Br- (aq) + 6 H+(aq) 3 Br2 (aq) + 3 H2O

Initial Concentrations Rate in M per unit

time

Mixture [BrO3 ]/ M‑ [Br-] / M [H+] /M

A 0.0050 0.025 0.030 10B 0.010 0.025 0.030 20C 0.010 0.050 0.030 40D 0.010 0.050 0.060 160

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Another example

Initial Concentrations Rate in mol L-1 hr-1Experiment [A] [B]

1 0.50 0.20 0.50 x 10-2

2 0.75 0.20 0.50 x 10-2

3 1.00 0.20 0.50 x 10-2

4 0.50 0.40 1.00 x 10-2

5 0.50 0.60 1.50 x 10-2

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Experiment Starting [X] /M

Starting [Y]/M

Rate/Mmin-1

1 0.00500 0.0250 1.002 0.0100 0.0250 4.003 0.00500 0.0125 1.00

Finding order of a reaction with algebra

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Chemical Kinetics In kinetics we study the rate at which a chemical process occurs. Besides information about the speed at which reactions occur, kinetics