Chemical Kinetics
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Chemical Kinetics. A B. rate =. D [A]. D [B]. rate = -. D t. D t. Chemical Kinetics. Thermodynamics – does a reaction take place? Kinetics – how fast does a reaction proceed?. Reaction rate is the change in the concentration of a reactant or a product with time ( M /s).

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Chemical Kinetics

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Chemical kinetics

Chemical Kinetics


Chemical kinetics

A B

rate =

D[A]

D[B]

rate = -

Dt

Dt

Chemical Kinetics

Thermodynamics – does a reaction take place?

Kinetics – how fast does a reaction proceed?

Reaction rate is the change in the concentration of a reactant or a product with time (M/s).

D[A] = change in concentration of A over

time period Dt

D[B] = change in concentration of B over

time period Dt

Because [A] decreases with time, D[A] is negative.

13.1


Chemical kinetics

A B

time

rate =

D[A]

D[B]

rate = -

Dt

Dt

13.1


Chemical kinetics

Br2(aq) + HCOOH (aq) 2Br-(aq) + 2H+(aq) + CO2(g)

time

Br2(aq)

393 nm

393 nm

Detector

light

D[Br2] aDAbsorption

13.1


Chemical kinetics

Br2(aq) + HCOOH (aq) 2Br-(aq) + 2H+(aq) + CO2(g)

slope of

tangent

slope of

tangent

slope of

tangent

[Br2]final – [Br2]initial

D[Br2]

average rate = -

= -

Dt

tfinal - tinitial

instantaneous rate = rate for specific instance in time

13.1


Chemical kinetics

rate

k =

[Br2]

rate a [Br2]

rate = k [Br2]

= rate constant

= 3.50 x 10-3 s-1

13.1


Chemical kinetics

2H2O2 (aq) 2H2O (l) + O2 (g)

[O2] = P

n

V

1

1

D[O2]

P = RT = [O2]RT

RT

RT

DP

rate =

=

Dt

Dt

measure DP over time

PV = nRT

13.1


Chemical kinetics

2H2O2 (aq) 2H2O (l) + O2 (g)

13.1


Chemical kinetics

2A B

aA + bB cC + dD

rate = -

=

=

rate = -

= -

D[D]

D[B]

D[A]

D[B]

D[C]

D[A]

rate =

1

1

1

1

1

Dt

Dt

Dt

Dt

Dt

Dt

c

d

a

2

b

Reaction Rates and Stoichiometry

Two moles of A disappear for each mole of B that is formed.

13.1


Chemical kinetics

Write the rate expression for the following reaction:

CH4(g) + 2O2(g) CO2(g) + 2H2O (g)

D[CO2]

=

Dt

D[CH4]

rate = -

Dt

D[H2O]

=

Dt

D[O2]

= -

1

1

Dt

2

2

13.1


Chemical kinetics

aA + bB cC + dD

The Rate Law

The rate law expresses the relationship of the rate of a reaction to the rate constant and the concentrations of the reactants raised to some powers.

Rate = k [A]x[B]y

reaction is xth order in A

reaction is yth order in B

reaction is (x +y)th order overall

13.2


Chemical kinetics

F2(g) + 2ClO2(g) 2FClO2(g)

rate = k [F2]x[ClO2]y

Double [F2] with [ClO2] constant

Rate doubles

x = 1

rate = k [F2][ClO2]

Quadruple [ClO2] with [F2] constant

Rate quadruples

y = 1

13.2


Chemical kinetics

F2(g) + 2ClO2(g) 2FClO2(g)

1

Rate Laws

  • Rate laws are always determined experimentally.

  • Reaction order is always defined in terms of reactant (not product) concentrations.

  • The order of a reactant is not related to the stoichiometric coefficient of the reactant in the balanced chemical equation.

rate = k [F2][ClO2]

13.2


Chemical kinetics

Determine the rate law and calculate the rate constant for the following reaction from the following data:

S2O82-(aq) + 3I-(aq) 2SO42-(aq) + I3-(aq)

rate

k =

2.2 x 10-4 M/s

=

[S2O82-][I-]

(0.08 M)(0.034 M)

rate = k [S2O82-]x[I-]y

y = 1

x = 1

rate = k [S2O82-][I-]

Double [I-], rate doubles (experiment 1 & 2)

Double [S2O82-], rate doubles (experiment 2 & 3)

= 0.08/M•s

13.2


Chemical kinetics

A product

rate

=

[A]

M/s

D[A]

-

M

= k [A]

Dt

[A] = [A]0exp(-kt)

ln[A] = ln[A]0 - kt

D[A]

rate = -

Dt

First-Order Reactions

rate = k [A]

= 1/s or s-1

k =

[A] is the concentration of A at any time t

[A]0 is the concentration of A at time t=0

13.3


Chemical kinetics

Decomposition of N2O5

13.3


Chemical kinetics

The reaction 2A B is first order in A with a rate constant of 2.8 x 10-2 s-1 at 800C. How long will it take for A to decrease from 0.88 M to 0.14 M ?

0.88 M

ln

0.14 M

=

2.8 x 10-2 s-1

ln

ln[A]0 – ln[A]

=

k

k

[A]0

[A]

[A]0 = 0.88 M

ln[A] = ln[A]0 - kt

[A] = 0.14 M

kt = ln[A]0 – ln[A]

= 66 s

t =

13.3


Chemical kinetics

[A]0

ln

[A]0/2

0.693

=

=

=

=

k

k

What is the half-life of N2O5 if it decomposes with a rate constant of 5.7 x 10-4 s-1?

ln2

ln2

0.693

=

k

k

5.7 x 10-4 s-1

First-Order Reactions

The half-life, t½, is the time required for the concentration of a reactant to decrease to half of its initial concentration.

t½ = t when [A] = [A]0/2

= 1200 s = 20 minutes

How do you know decomposition is first order?

units of k (s-1)

13.3


Chemical kinetics

A product

# of

half-lives

[A] = [A]0/n

First-order reaction

1

2

2

4

3

8

4

16

13.3


Chemical kinetics

13.3


Chemical kinetics

A product

rate

=

[A]2

M/s

D[A]

1

1

-

M2

= k [A]2

=

+ kt

Dt

[A]

[A]0

t½ =

D[A]

rate = -

Dt

1

k[A]0

Second-Order Reactions

rate = k [A]2

= 1/M•s

k =

[A] is the concentration of A at any time t

[A]0 is the concentration of A at time t=0

t½ = t when [A] = [A]0/2

13.3


Chemical kinetics

A product

rate

[A]0

D[A]

-

= k

Dt

[A]0

t½ =

D[A]

2k

rate = -

Dt

Zero-Order Reactions

rate = k [A]0 = k

= M/s

k =

[A] is the concentration of A at any time t

[A] = [A]0 - kt

[A]0 is the concentration of A at time t=0

t½ = t when [A] = [A]0/2

13.3


Chemical kinetics

Concentration-Time Equation

Order

Rate Law

Half-Life

1

1

=

+ kt

[A]

[A]0

=

[A]0

t½ =

t½ =

ln2

2k

k

1

k[A]0

Summary of the Kinetics of Zero-Order, First-Order

and Second-Order Reactions

[A] = [A]0 - kt

rate = k

0

ln[A] = ln[A]0 - kt

1

rate = k [A]

2

rate = k [A]2

13.3


Chemical kinetics

A + B C + D

Endothermic Reaction

Exothermic Reaction

The activation energy (Ea) is the minimum amount of energy required to initiate a chemical reaction.

13.4


Chemical kinetics

lnk = -

+ lnA

Ea

1

T

R

Temperature Dependence of the Rate Constant

k = A •exp( -Ea/RT )

(Arrhenius equation)

Eais the activation energy (J/mol)

R is the gas constant (8.314 J/K•mol)

T is the absolute temperature

A is the frequency factor

13.4


Chemical kinetics

lnk = -

+ lnA

Ea

1

T

R

13.4


Chemical kinetics

13.4


Chemical kinetics

2NO (g) + O2 (g) 2NO2 (g)

Elementary step:

NO + NO N2O2

+

Elementary step:

N2O2 + O2 2NO2

Overall reaction:

2NO + O2 2NO2

Reaction Mechanisms

The overall progress of a chemical reaction can be represented at the molecular level by a series of simple elementary steps or elementary reactions.

The sequence of elementary steps that leads to product formation is the reaction mechanism.

N2O2 is detected during the reaction!

13.5


Chemical kinetics

Elementary step:

NO + NO N2O2

+

Elementary step:

N2O2 + O2 2NO2

Overall reaction:

2NO + O2 2NO2

Intermediates are species that appear in a reaction mechanism but not in the overall balanced equation.

An intermediate is always formed in an early elementary step and consumed in a later elementary step.

  • The molecularity of a reaction is the number of molecules reacting in an elementary step.

  • Unimolecular reaction – elementary step with 1 molecule

  • Bimolecular reaction – elementary step with 2 molecules

  • Termolecular reaction – elementary step with 3 molecules

13.5


Chemical kinetics

Unimolecular reaction

Bimolecular reaction

Bimolecular reaction

A + B products

A + A products

A products

The rate-determining step is the sloweststep in the sequence of steps leading to product formation.

Rate Laws and Elementary Steps

rate = k [A]

rate = k [A][B]

rate = k [A]2

  • Writing plausible reaction mechanisms:

  • The sum of the elementary steps must give the overall balanced equation for the reaction.

  • The rate-determining step should predict the same rate law that is determined experimentally.

13.5


Chemical kinetics

The experimental rate law for the reaction between NO2 and CO to produce NO and CO2 is rate = k[NO2]2. The reaction is believed to occur via two steps:

Step 1:

Step 2:

NO2 + NO2 NO + NO3

NO2+ CO NO + CO2

NO3 + CO NO2 + CO2

What is the equation for the overall reaction?

What is the intermediate?

NO3

What can you say about the relative rates of steps 1 and 2?

rate = k[NO2]2 is the rate law for step 1 so

step 1 must be slower than step 2

13.5


Chemical kinetics

CH2

CH2

CH2

CH2

CH2

CH2

2

2

CH2

CH2

Chemistry In Action: Femtochemistry

CH2

CH2

CH2

CH2

CH2

CH2

CH2

CH2

13.5


Chemical kinetics

uncatalyzed

catalyzed

Ea

k

Ea< Ea

A catalyst is a substance that increases the rate of a chemical reaction without itself being consumed.

k = A •exp( -Ea/RT )

ratecatalyzed > rateuncatalyzed

13.6


Chemical kinetics

In heterogeneous catalysis, the reactants and the catalysts are in different phases.

  • Haber synthesis of ammonia

  • Ostwald process for the production of nitric acid

  • Catalytic converters

In homogeneous catalysis, the reactants and the catalysts are dispersed in a single phase, usually liquid.

  • Acid catalysis

  • Base catalysis

13.6


Chemical kinetics

Fe/Al2O3/K2O

N2 (g) + 3H2 (g) 2NH3 (g)

catalyst

Haber Process

13.6


Chemical kinetics

4NH3(g) + 5O2(g) 4NO (g) + 6H2O (g)

2NO (g) + O2(g) 2NO2(g)

2NO2(g) + H2O (l) HNO2(aq) + HNO3(aq)

Hot Pt wire

over NH3 solution

Pt-Rh catalysts used

in Ostwald process

Ostwald Process

Pt catalyst

13.6


Chemical kinetics

catalytic

CO + Unburned Hydrocarbons + O2

CO2 + H2O

converter

catalytic

2NO + 2NO2

2N2 + 3O2

converter

Catalytic Converters

13.6


Chemical kinetics

Enzyme Catalysis

13.6


Chemical kinetics

enzyme

catalyzed

uncatalyzed

D[P]

rate =

Dt

rate = k [ES]

13.6


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