konsep kinetika reaksi

Reaction rate

Learning objective

Explore the rate of chemical reaction

Topics:Definition of reaction rate

Outline the techniques for its measurement

Variables control

Definition of reaction rate

The reaction rate is the increase in

molar concentration of a product of

a reaction per unit time

It can also be expressed as the decrease

in molar concentration of a reactant

per unit time.

Why do we study the reaction rate?

Video of reaction

Halloween reaction

Nurdrage

HALLOWEN CLOCK.mp4

hydrogen peroxide.mp4

nurdrage.mp4

Why do we study the reaction rate?

To predict how quickly a reaction mixture

approaches equilibrium

To optimise the rate by the appropiate

choice of condition

To an understanding of the mechanisme of

reactions, their analysis into a sequence of

elementary steps

A B

13.1

rate = -D[A]

Dt

rate = D[B]

Dt

time

Reaction rate is the change in the concentration of a

reactant or a product with time (M/s).

A B

rate = -D[A]

Dt

rate = D[B]

Dt

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.

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

time

393 nm

light

Detector

D[Br2] a DAbsorption


average rate = -D[Br2]

Dt= -

[Br2]final – [Br2]initial

tfinal - tinitial

slope of

tangentslope of

tangent slope of

tangent

instantaneous rate = rate for specific instance in time

rate a [Br2]

rate = k [Br2]

k = rate

[Br2]= rate constant

= 3.50 x 10-3 s-1

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

PV = nRT

P = RT = [O2]RTnV

[O2] = PRT

1

rate = D[O2]

Dt RT

1 DP

Dt=

measure DP over time

Video of reaction

Hydrogen peroxide reaction

Nurdrage



nurdrage.mp4

Definition of Reaction Rates

Consider the gas-phase decomposition of

dinitrogen pentoxide.

)g(O)g(NO4)g(ON2 2252

– If we denote molar concentrations using

brackets, then the change in the molarity of

O2 would be represented as

where

the symbol, D (capital Greek delta), means

“the change in.”

]O[ 2D

Figure 1:

The instantaneous

rate of reaction.

Figure 2 shows the increase in concentration

of O2 during the decomposition of N2O5.


• Note that the rate decreases as the

reaction proceeds.

Figure 2:

Calculation

of the

average rate.

Then, in a given time interval, Dt , the molar

concentration of O2 would increase by D[O2].

– This equation gives the average rate over the

time interval, Dt.

– If Dt is short, you obtain an instantaneous

rate, that is, the rate at a particular instant.

(Figure 1)


– The rate of the reaction is given by:

t

][O oxygen of formation of Rate 2

D

D

Because the amounts of products and

reactants are related by stoichiometry, any

substance in the reaction can be used to

express the rate.


t

]O[N ON of iondecomposit of Rate 52

52 -D

D

• Note the negative sign. This results in a

positive rate as reactant concentrations

decrease.

The rate of decomposition of N2O5 and the

formation of O2 are easily related.


)(t

]O[N

21

t

][O 522 -D

D

D

D

• Since two moles of N2O5 decompose for

each mole of O2 formed, the rate of the

decomposition of N2O5 is twice the rate of

the formation of O2.

Experimental Determination of

Reaction Rates To obtain the rate of a reaction you must

determine the concentration of a reactant or

product during the course of the reaction.

– One method for slow reactions is to withdraw

samples from the reaction vessel at various

times and analyze them.

– More convenient are techniques that

continuously monitor the progress of a reaction

based on some physical property of the

system.

Figure 3: An experiment to follow the concentration

of N2O5 as the decomposition proceeds.

Experimental Determination of

Reaction Rates Gas-phase partial pressures.

– Manometer readings provide the concentration

of N2O5 during the course of the reaction based

on partial pressures.

– When dinitrogen pentoxide crystals are sealed

in a vessel equipped with a manometer (see

Figure 3) and heated to 45oC, the crystals

vaporize and the N2O5(g) decomposes.

)g(O)g(NO4)g(ON2 2252

Reaction Rates and Stoichiometry

2A B

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

rate = D[B]

Dtrate = -

D[A]

Dt

1

2

aA + bB cC + dD

rate = -D[A]

Dt

1

a= -

D[B]

Dt

1

b=

D[C]

Dt

1

c=

D[D]

Dt

1

d

Write the rate expression for the following reaction:

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

rate = -D[CH4]

Dt= -

D[O2]

Dt

1

2=

D[H2O]

Dt

1

2=

D[CO2]

Dt

The Rate Law

13.2

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.

aA + bB cC + dD

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

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

rate = k [F2]x[ClO2]

y

Double [F2] with [ClO2] constant

Rate doubles

x = 1

Quadruple [ClO2] with [F2] constant

Rate quadruples

y = 1

rate = k [F2][ClO2]

13.2

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

rate = k [F2][ClO2]

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.

1

13.2

Determine the rate law and calculate the rate constant for

the following reaction from the following data:

S2O82- (aq) + 3I- (aq) 2SO4

2- (aq) + I3- (aq)

Experiment [S2O82-] [I-]

Initial Rate

(M/s)

1 0.08 0.034 2.2 x 10-4

2 0.08 0.017 1.1 x 10-4

3 0.16 0.017 2.2 x 10-4

13.2

Determine the rate law and calculate the rate constant for

the following reaction from the following data:

S2O82- (aq) + 3I- (aq) 2SO4

2- (aq) + I3- (aq)

Experiment [S2O82-] [I-]

Initial Rate

(M/s)

1 0.08 0.034 2.2 x 10-4

2 0.08 0.017 1.1 x 10-4

3 0.16 0.017 2.2 x 10-4

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

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

y = 1

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

x = 1

k = rate

[S2O82-][I-]

=2.2 x 10-4 M/s

(0.08 M)(0.034 M)= 0.08/M•s

13.2

rate = k [S2O82-][I-]

First-Order Reactions

A product

rate = -D[A]

Dt

rate = k [A]

D[A]

Dt= k [A]-

(1)

(2)

(3)

kdtdAA

1

A

A

t

o

dtkdAA

0

1

ktA

A

ktAA

0

0

ln

lnln

kteAA 0

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

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


A product rate = -D[A]

Dt

rate = k [A]

k = rate

[A]= 1/s or s-1

M/sM

=

D[A]

Dt= k [A]-



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

2N2O5 4NO2 (g) + O2 (g)

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 ?

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

kt = ln[A]0 – ln[A]

t =ln[A]0 – ln[A]

k= 66 s

[A]0 = 0.88 M

[A] = 0.14 M

ln[A]0

[A]

k=

ln0.88 M

0.14 M

2.8 x 10-2 s-1=


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

ln[A]0

[A]0/2

k=t½

ln2

k=

0.693

k=

What is the half-life of N2O5 if it decomposes with a rate

constant of 5.7 x 10-4 s-1?

t½ln2

k=

0.693

5.7 x 10-4 s-1= = 1200 s = 20 minutes

How do you know decomposition is first order?

units of k (s-1)

A product

First-order reaction

# of

half-lives [A] = [A]0/n

1

2

3

4

2

4

8

16

Second-Order Reactions

A product

rate = -D[A]

Dt

rate = k [A]2

D[A]

Dt= k [A]2-



1

[A]=

1

[A]0+ kt

kdtdAA

2

1

A

A

t

dtkdAA

0 0

2

1

ktAA o

11

Second-Order Reactions

13.3


Dtrate = k [A]2

k = rate

[A]2= 1/M•s

M/sM2=

D[A]

Dt= k [A]2-



1

[A]=

1

[A]0+ kt

t½ = t when [A] = [A]0/2

t½ =1

k[A]0

Zero-Order Reactions

13.3


Dtrate = k [A]0 = k

k = rate

[A]0= M/s

D[A]

Dt= k-



t½ = t when [A] = [A]0/2

t½ =[A]02k

[A] = [A]0 - kt

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

and Second-Order Reactions

Order Rate Law

Concentration-Time

Equation Half-Life

0

1

2

rate = k

rate = k [A]

rate = k [A]2

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

1

[A]=

1

[A]0+ kt

[A] = [A]0 - kt

t½ln2

k=

t½ =[A]02k

t½ =1

k[A]0

13.3

Exothermic Reaction Endothermic Reaction

The activation energy (Ea ) is the minimum amount of

energy required to initiate a chemical reaction.

A + B AB C + D++

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

Temperature Dependence of the Rate Constant

k = A • exp( -Ea / RT )

Ea is 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

lnk = -Ea

R

1

T+ lnA

(Arrhenius equation)

13.4

lnk = -Ea

R1T

+ lnA

c2) A + B products

ktBA

BA

BA

baxabax

xABA

dx

xBBA

dx

xBxA

dx

xBxAkdt

dx

BBAAx

BAKdt

AdR

t

t

tt

][][

][][ln

][][

1

)ln(11

][][][][][][][][

][][

][][][][

][][][

0

0

00

00000000

00

00

11

H2 + O OH + H

b

a

b

a

d) Third-Order reaction:

d1) A + A + A products

KtAA

dtKA

AdKt

AA

xnx

dx

KdtA

AdAK

dt

AdR

t

tA

At

n

b

a

n

t

6][

1

][

1

3][

][3

][

1

][

1

2

1

1

)1(

1

3][

][][

][

3

1

20

2

0

][

][

320

2

1

3

3

0

t

[A]t-2

[A]0-2

m = 6K

b

a

I + I + M I2 + M

d2) A + A + B products

Kt

BA

BA

BAAABA

KdtyByA

dy

yByAKdt

dy

BA

yBB

yAA

BAKdt

AdR

t

t

t

t

t

][][

][][ln

][2][

1

][

1

][

1

][2][

1

][2][

][2][

][,][

][][

2][][

][][][

2

1

0

0

200000

02

0

02

0

00

0

0

2

Partial fractions!

tKAA

orderpseudoBKK

AKdt

AdR

t

nd

'2][

1

][

1

2],['

]['][

2

1

0

2

[B]>>[A]

[B] = const.

O + O2 + M O3 + M

d3) A + B + C products

]][][[][

CBAKdt

AdR

Reaction rate for n order respect with only one reactant

KtnnAA

nKtAAn

dtnkA

Ad

AKdt

Ad

nR

nnt

nnt

A

A

t

n

n

t

)1(][

1

][

1

][

1

][

1

)1(

1

][

][

][)(

1

)1(0

)1()1(0

)1(

][

][ 00

e) Reaction half-lives: Alternative method to

determine the reaction order

2

][][ 0

2

1

2

1

AAt

2ln1

][

2/][ln

2/][][

][

][ln

2/1

2/1

0

0

0

0

Kt

KtA

A

AA

KtA

A

t

t

First-order reaction

For a reaction of order n>1 in a

single reactant

Independent of [A]0

1

0

1

2

1])[1(

12

n

n

AnKt

This is function of [A]0

1st order reaction half-lives:

0 10 20 30 40 500

0.2

0.4

0.6

0.8

11

6.738 103

A t( )

500 t

T dependence of the rate constant, K

The Arhenius equation:

K(T):

•[i]

•t

•pH (only in solution)

T! (Strongly)

kT

E

T

Act

AeK )(

T

ln[K(T)]

lnA

m = -EAct/k

Chemical coordinates

E

EAct(F)EAct(R)

Reac

Prod.

DH0Rxn

A :Frequency factor

0

Re

0

Pr

0

0 )()(

acodRxn

ActActRxn

HHH

REFEH

DDD

D

Ineffective Effective

Determinant Parameters

Catalysis

Task for next week:

Video animation or experiment

on factors affect the reaction rate

Work in group

Each group must collect at least

four video

konsep kinetika reaksi

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