chapter 2 - chemical kinetics.pptx

8/19/2019 chapter 2 - Chemical Kinetics.pptx

1/92

Chemical Kinetics

CHAPTER 2

1

NURUL’ AIN BINTI JAMION


2/92

2

CHAPTER 2

Rates of ReactionFactors aecting rates of

reaction

Rates law and order ofreaction

Methods to determine order

of reactionRelation between rate andtemperature

Reaction MechanismCatalytic kinetics


3/92


4/92

Chemical kinetics - is the study of chemicalrxns/processes with respect to reaction rates,efect o various variables

INTRODUCTION


5/92

INTRODUCTION


6/92

RA!s of R!AC"#$

Defnition:

A measure of the change in the concentration ofa reactant or product in a gi%en time&

6

'nit ( mole / liter (mol / )

second s

( Mol )-* s-*

t

xrate

∆

∆=

][


7/92


8/92

Reaction Rates+cont&,

depending on the experimental conditions

average rate rate measured between long timeinter%al

instantaneos rate rate measured between%ery short inter%al

initial rate instantaneous rate at the beginningof an experiment


9/92


10/92


11/92


12/92

RA!s of R!AC"#$

Reactant .roduct

hen the reactants change into products

the concentration of reactants decreases

the concentration of products increases

with time&

12


13/92

13

rate = -∆[A]

∆t

rate =∆[B]

∆t

A !


14/92

Measuring Reaction Rate

o measure a reaction rate0 we usually monitoreither a "ro#ct or a reactant for its change&

1ome of the characteristics to be monitoredare

-changes in "ressre

$changes in color %se s"ectro&otometer '

$tem"eratre &or e(othermic or

en#othermic reaction) an#$"resence o& certain ke* s+stance


15/92

Reaction Rates an#

,toichiometr*2A B

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

formed.

rate = ∆[B]

∆ t rate = - ∆[A]

∆ t 12

a A + bB c C + d D

rate = -∆[A]

∆ t

1

a= -

∆[B]

∆ t

1

b=

∆[C]

∆ t

1

c =

∆[D]

∆t

1

d


16/92

16

Example : Write the rate expression for the following reaction:

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

rate = -∆[CH4]

∆t = -

∆[O2]

∆t

1

2 =

∆[H2O]

∆t

1

2 =

∆[CO2]

∆t

Reaction Rates an#

,toichiometr*


17/92

17

R2

rite the 3ierential rate e4uation forabo%e e4uation&

1olution

+g,#+g,5$#+g,#6$ 6676 +→

+=

+=

−=

dt

]d[O

dt

]d[NO

4

1

dt

]Od[N

2

1Rate 2252


18/92

18

R2

)(3)(2 23 g O g O →

dt

Od

dt

Od rate

][

3

1][

2

1 23 +=−=


19/92

R2


20/92

R2

Consider the reaction for the combustionof methane0 C85 0

C85+g, 9 6#6+g, → C#6+g,9 686#+g,

"f the methane is burning at a rate of :&*; mol)-*s-* 0 at what rates are C#6 and 86# being formed<


21/92

1olution

C85+g, 9 6#6+g, → C#6+g,9 686#+g,

=i%en d>C85? ( - :&*; mol)-*s-*

dt

so@ -d>C85? ( d>C#6? ( d>86#?

dt dt 6 dt


22/92

1olution o calculate rate of $6@

d>C#6? ( - d>C85? dt dt

( -+-:&*; mol)-*s-* ,

( :&*; mol)-*s-*


23/92

1olution

o calculate rate of 86#@

d>86#? ( 6x -d>C85?

dt dt

( 6 x +-:&*; mol)-*s-* ,

( :&B6 mol)-*s-*


24/92

24

Answer :

At t = 0s, [C2H4] = 0.884 M

At t = 40s, [C2H4] = 0.328 M

Rate = (0.884 0.328)M 40s

= 0.013! Ms-1

= 0.014 Ms-1 s", t#e rate "$ rea%t&"'s at 40s &s 0.014 Ms-1

TRY : Calculate the rate of reaction of C2! con"erted to C!#from t$% to t$!%s

2C2H4 () C4H8()

ime+s, : *: 6: 5: ;:

>C685?/mol )-* :&5 :&;6* :&5DE :&B6 :&*;E


25/92

The Rate &aw

13.2

a A + bB c C + d D

Rate = k [A]a

[B]b

THE RATE -A. !xpresses the relationship of the rate of a

reaction to the rate constant and the concentration of thereactants raised to some powers&

k ( rate constant

a b ( reaction order withrespect to A G

a 9 b ( o%erall reactionorder


26/92

2 (g ) + 2C*O2 (g ) 2C*O2 (g )

rate = k [2][C*O2]

• Rate *as are alwa's eter&'e e/er&e'ta** '*ess

" #ae ee' t"* a rate-*&&t&' "r s*" ste/.

• Rea%t&"' "rer &s alwa's e$&'e &' ters "$ reactant

('"t /r"%t) %"'%e'trat&"'s.

• #e "rer "$ a rea%ta't is not re*ate t" t#est"&%#&"etr&% %"e$$&%&e't "$ t#e rea%ta't &' t#e a*a'%e

%#e&%a* e5at&"'.

1

13.2


27/92

27

he number of concentration factors inthe rate e4uation&

Can only be determined e("erimentall*/

"t shows the exact dependence of rate onconcentration&

not from the number of moles as written

in the stoichiometric e4uation& he nits o& the rate constant #e"en#

on the overall reaction or#er&

THE ORDER O0 REACTION


28/92

28

First order reaction- A reaction with a single reactant A- Rate is directly proportional to >A?

rate= k [A]

1econd order reaction- Rate is directly proportional to the s4uare of >A?

rate= k[A]2

Hero order reaction-

the rate is not dependent on >A?

rate= k [A]0 = k (1) = k


29/92

2!

8! #I!RA)) R!AC"#$#R3!R

Rate ( k > A ?m > G ?n

he o%erall order of reaction ( sum ofthe indi%idual orders ( m9n&

with respect to each reactant in the ratelaw&


30/92


31/92


32/92

Rates as Functions of Reactant

Concentrations "n a general reaction0

a A 1 b ! 1 products

Rate = k[A] x [B]y

"f concentrations of ! is ke"t

constant0 you can measure the reaction

rate of A at %arious concentrations&


33/92

33

DETERININ3 REACTIONORDER

"nitialrate

"ntegrated Rate)aw

8alflife


34/92

34

*& "nitial Rate Method

o Jnd the reaction orders by initialrate method0

Run a series of experiments0

!ach of the experiments ha%e adierent set of reactant

concentration0From each experiments0 obtain an

initial rate&


35/92

3

!xample *

he results shown below are obtained forthe reaction between A and G& 3eterminethe rate e4uation for the reaction betweenA and G&

!xperiment Concentration+mol )-*,

"nitial rate +mol )-* s-*,

>A? >G?

* :&:7D7 :&:6*; &6*x*:-B

6 :&:D*B :&:6*; *&6;x*:-6

B :&:7D7 :&:BBB *&6;x*:

-6


36/92

36

Calculating m in >A?m@ we take the ratio ofthe rate laws for two experiment * and 60

in which >A? %aries but >G? is constant

hat is0 :&;76 ( :&:;m

log :&;76 ( m x log :&:;

m ( *&E≈ 6

hus0 the reaction is

second order with respect to A

nm

nm

>G?>A?krate6

>G?>A?krate*

=

= m6

B

:&:D*B

:&:7D7

*:*&6;

*:&6*

=

×

×−

−


37/92


38/92

38

E(am"le 2

he reaction between C and 3 is represented asfollows

he experiment results are shown in the tablebelow& 3etermine the rate e4uation for thereaction between C and 3&

!xperiment

"nitial concentration+M,

imeinter%al

+minutes,

he change inconcentration

of C +M,

>C? >3?

* :&:7 *&: B: 6&7 x *:-B

6 :&:7 6&: B: *&: x *:-6

-B

!3C →+


39/92

3!

hat is0 :&67 ( +:&7,n

log :&67 ( m x log :&7 n ( 6

hus0 the reaction is second order with respect to 3

nm

nm

>3?>C?k6rate

>3?>C?krate*

== n

2

3

2.0

1.0

101.0

102.5

=

××

−

−


40/92

E(am"le :

Derive the rate la4 &or the reaction

A 1 ! 1 C "ro#cts&rom the &ollo4ing #ata :

Expriment 1 2 3 4

[A]0 0.100 0.200 0.200 0.100

[B]0 0.100 0.100 0.300 0.100

[C]0 0.100 0.100 0.100 0.400

rate 0.100 0.800 7.200 0.400


41/92

solution

Assume rate = k >A? x >G? y >C? z

From experiment 1 and 2, we have

!"#!! k $!"2% x $!"1% y $ !"1% z

&&&&& = &&&&&&&&&&&&&&&&&&&&&&&&

!"1!! k $!"1% x $!"1% y $!"1% z

'hus, # = 2 x ( and x = ln# ) ln2 = *

+y similar procedures, we et y = 2 and z = 1"

'hus, the rate law is rate = k $-%* $+%2 $.%

overall order = */2/1 = 0


42/92

42

hat is0 :&7 ( +:&7,m

log :&7 ( m x log :&7 m ( *

hus0 the reaction is rst order with respect to C

he rate e4uation is

nm

nm

>3?>C?kBrate

>3?>C?krate*

=

= m

B

B

:&*

:&:7

*:7&:

*:6&7

=×

×−

−

6>C?>3?krate =


43/92

TR5 he following data refers to the following

reactionC8BC#C8B 9 "6 9 89 → C8BC#C86" 9 8"

Expt. (C)C*C)+, - (

+,- (/2+,

- /nitial rate,-s01

1 3.0 0.2 0.02 6.0 10-6

2 3.0 0.4 0.02 12.0 10-6

3 4.0 0.4 0.02 16.0 10-6

4 4.0 0.2 0.04 8.0 10-6


44/92

Gased on the abo%e information0

Calculate the order with respect toC8BC#C8B 0 "6 and 89&

rite the rate law for the reaction&

hat is the o%erall order of the reaction<

3etermine the rate constant0k with thecorrect unit&

Calculate the rate of reaction if theconcentration of propanone0 iodine and

hydrogen ion are each :&7 mol dm-B&


45/92

45

Lero order reaction

you will get a hori6ontalline0

because rate ( k>A?:

rate = k +a horiLontal line,

rate

k

independent to >reactant?

>A?


46/92

46

Jrst order reaction

For Jrst order0 rate ( k>A?

rate

the plot is a straight line %linear'

[A]


47/92

47

second order reaction

the plot is a branch of a prabola0 because

rate = k >A?6

Rate

[A]

The Rate -a4 7


48/92

Summary of the Kinetics of Zero-Orer! "irst-Oreran Secon-Orer Reactions

The Rate -a4 7Or#er O& Reaction

*rder Rate &aw

/ntegrated rate

law alf0&ife

0

1

2

rate = k

rate = k [A]

rate = k [A]2

*'[A] = *'[A]0 - kt

1

[A]=

1

[A]0+ kt

[A] = [A]0 - kt

t *' 2

k =

t ½ =[A]02k

t ½ =1

k [A]0

nit

Ms-1

s-1

M-1s-1


49/92

#e rea%t&"' "$ '&tr&% "&e &t# #r"e' at 1280" C &s

2O() + 2H2() 2() + 2H2O()

r" t#e $"**"&' ata %"**e%te at t#&s te/eratre, eter&'e 9

a) #e rate *a

) #e rate %"'sta't

%) #e rate "$ t#e rea%t&"' #e' [O] = 1.2 10-3M a' [H2] = 6.0 10-3 M

experiment >$#? +M, >86? +M, "nitialrate+M/s,

* 7&: x *:-B 6&: x *:-B *&B x *:-7

6 *:&: x *:-B 6&: x *:-B 7&: x *:-7

B *:&: x *:-B 5&: x *:-B *:&: x *:-7

E(ercise 8 $ Initial rate

metho#


50/92

.:#at &s t#e #a*$-*&$e "$ 2O &$ &t e%"/"ses &t# a rate

%"'sta't "$ .7 10-4 s-1;

t *' 2

k =

0.6!3

.7 10-4 s-1= = 1200 s = 20 &'tes

H" " "


51/92

E(ercise 9 Integrate# rate $la4

#e rea%t&"' 2A B &s $&rst "rer &' A &t# a rate%"'sta't "$ 2.8 10-2 s-1 at 800C. H" *"' &** &t ta


52/92

*& he rate constant for the decomposition of$6#7 to $#6 and #6 at D:°C is ;&6 x *:-B s-*&

1uppose we start with :&B:: mol of $6#7+g, in

a :&7:: ) container&8ow many moles of $6#7 will remain after *&7

min<

8ow many minutes will it take for the 4uantity

of $6#7 to drop to :&:B: mol<hat is the half-life of $6#7 at D:°C <

R2


53/92


54/92

The Collision Theor* "n a chemical reaction0 bonds are

broken and new bonds are formed&

Molecules can only react if theycollide with each other&

Furthermore0 molecules must collidewith the correct orientation and withenough energy to cause bondbreakage and formation&


55/92

The Collision Theor*

E;ective collision

E;ective collision

Ine;ective collision

Ine;ective collision

..3ut how do we define it as an effecti"e44


56/92

The Collision Theor* E00ECTI


57/92

The Collision Theor* : -ctivation nery

• "n other words0 there is a minimm amonto& energ* re=ire# &or reaction theactivation energ*) Ea&

• Oust as a ball cannot get o%er a hill if it does notroll up the hill with enough energy0 a reactioncannot occur unless the molecules possesssucient energy to get o%er the acti%ationenergy barrier& he activation energy (Ea ) is the

minimum amount of energy re4uired to

initiate a chemical reaction&


58/92

Acti%ation energy +cont&,

The activation energyis represented on areaction profile(reaction energydiagram) as theenergy barrierbetween reactants

and products.


59/92

Acti%ation energy+cont&,

he height of the acti%ation barrier inPuences the

rate

"fEa is large

0 then few molecules possessenough energy to get o%er the barrier@ fewproducts are formed and the rate is slo4&

"f Ea is small0 then many molecules possess

enough energy to get o%er the barrier@ moreproducts are formed and the rate is &aster&


60/92

.hat &actors a;ect the rate o&reactions>

&'%rease temperature

&'%rease concentration "$

&ss"*e rea%ta'ts

&'%rease surface area "$ s"*&

rea%ta'ts

se "$ a catal'st

&'%rease pressure "$ ase"s

rea%ta'ts

Eff t f t t t


61/92

Effect of temperature on rate#e #&#er t#e te/eratre, t#e $aster t#e rate "$ a rea%t&"'.

' a' rea%t&"'s, a r&se &' te/eratre "$ 10

>C %ases t#erate "$ rea%t&"' t" a//r"&ate* "*e.

:# "es &'%rease te/eratre

&'%rease t#e rate "$ rea%t&"';

At a #&#er te/eratre, /art&%*es#ae "re e'er. #&s ea's

t#e "e $aster a' are "re

*&


62/92

reaction

#e #&#er t#e %"'%e'trat&"' "$ a &ss"*e rea%ta't, t#e

$aster t#e rate "$ a rea%t&"'.

:# "es &'%rease %"'%e'trat&"' &'%rease t#e rate "$

rea%t&"';

At a #&#er %"'%e'trat&"', t#ere are "re /art&%*es &' t#e

sae a"'t "$ s/a%e. #&s ea's t#at t#e /art&%*es are

"re *&


63/92

1lower and slowerQ

Rea%t&"'s " '"t /r"%ee at a stea rate. #e start "$$ at

a %erta&' s/ee, t#e' et s*"er a' s*"er 't&* t#e st"/.

As t#e rea%t&"' /r"resses, t#e %"'%e'trat&"' "$ rea%ta'ts

e%reases.

#&s re%es t#e $re5e'% "$ %"**&s&"'s etee' /art&%*es

a' s" t#e rea%t&"' s*"s "'.

/er%e'tae %"/*et&"' "$ rea%t&"'

1%%5%5 265 6%5 765

reactants

product

!ect of pressure on rate of reaction


64/92

!ect of pressure on rate of reaction

#e as /art&%*es e%"e %*"ser t"et#er, &'%reas&' t#e

$re5e'% "$ %"**&s&"'s. #&s ea's t#at t#e /art&%*es are "re*&


65/92

!ect of surface area on rate of reaction A' rea%t&"' &'"*&' a s"*& %a' "'* ta


66/92

reaction 8time9

e n

e r g '

8 : ; 9

8ow catalysts increase rate of rxn<S#ees u# a reaction $ut is chemica%%y unchan&e at theen of the reaction'Ain& a cata%yst has e(act%y effect on )A'#ro*ies an a%ternati*e route for the reaction +ith%o+er acti*ation ener&y'

D&$$ere't %ata*sts "r< &'&$$ere't as, t "st

*"er t#e rea%t&"'?s

a%t&at&"' e'er (@a).

Ea with

catal'st

Ea without

catal'st


67/92

67

ype of catalysis

a, 8omogeneous

he catalyst is in the same phase as the reactants&

b, 8eterogeneous

he catalyst is in a dierent phase

than the reactants often a solidcatalyst and li4uid or gaseousreactants&


68/92

,o+ tem#erature an ata%ysts Affect Reaction Rates

The Arrhenius’ ).uation

/0 A e- Ea/RT

"re.uency

of co%%isions

Acti*ation

ener&y

Tem#erature in K

1as constant

2'345 J6mo% 7 K

3etermination of the %alues of


69/92

3etermination of the %alues of - and a

lineariLing the Arrhenius e4uation@


70/92

70

3etermination of the %alues of - and a

A plot of ln k verss 8?T will ha%e a slope of-a/ and an intercept of ln -&


71/92

71

!xample A plot of ln+k, %s& */+,

Rate "s. Temperature

= -7330.4 + 26.417

:

:&5

:&C

*&6

*&;

6

B&*!-:B B&6!-:B B&B!-:B B&5!-:B B&7!-:B B&;!-:B

1


72/92

!xample +cont&,

S*Calculate the acti%ation energy for this

reaction<

S6 Calculate rate would be expected for this

reaction at B:°C<SB 3etermine temperature is needed to

double this rate<


73/92

7

1olution for S*

0irst) correlate the actal e=ation o& theline 4ith the Arrhenis E=ation:

ln k @ $ Ea

?R%8?T' 1 ln A* @ m( 1 c

* @ $99B/( 1 2/8


74/92

74

1olution for S*

1lope (

!a ( - +1lope, x +R,

!a ( - +-DBB:&5 , x +&B*5 O/mol,Ea @ B/F kG?mol

1o the acti%ation energy for this reaction

is B/F kG?mol

a!

R−


75/92

75

1olution for S6

he rate at B:TC is U

+ B: 9 6DB ( B:B,

ln+k, ( -DBB:&5+*/, 9 67&5*D +gi%en,

ln+k, ( -DBB:&5+*/B:B , 9 67&5*D

ln+k, ( *&66

k ( e *&66

k @ 9/B s$8


76/92

76

1olution for SB

he temperature needed to double this

rateU&+ B&5 x 6 ( ;&,

ln+rate, ( -DBB:&5+*/, 9 67&5*D +gi%en,ln+;&, ( -DBB:&5+*/, 9 67&5*D

*&E* ( -DBB:&5+*/, 9 67&5*D

-6B&7: ( -DBB:&5+*/,

( B*6 ( B*6 6DB ( BETC


77/92


78/92

If +e ha*e t+o ifferent / *a%ues at ifferent

tem#eratures! you can com$ine the t+o

reactions to &i*e8

%n / 9 0 )a 4 4

/ 4 R T4 T9

.here T8 T2

Example


79/92

Example

#e rate %"'sta't $"r t#e $"rat&"' "$ H at 600 a' 60

are easre a' $"' t" e 2.710-4 M-1s-1 a' 3.10-3

M-1s-1, res/e%t&e*. Deter&'e t#e a%t&at&"' e'er a't#e rate %"'sta't at 800 .

H2+2 2H

1olution

ln +B&7x*:-B , ( !a * - * +6&Dx*:-5 , &B*5 Omol-* ;:: ;7:

ln +*B&: , ( !a +*&6x*:-5, &B*5 Omol-*

!a ( 6&7; x &B*5 Omol-* ( *&;;x*:7 Omol-*

+*&6x*:-5, ( 8 kGmol$8

%n / 9 0 )a 4 4

/ 4 R T4 T9

E(am"le


80/92

>olution

To determine rate %"'sta't at 800.

*'


81/92

R2

The rate constant o& a reaction is 8/2 ( 8B$

9 sec$8 at 9BB C an# 2/8 ( 8B$9 sec$8 at BB C/Calclate the energ* o& activation o& thereaction/

Answer !a ( 55&*; kO mol-*

%n / 9 0 )a 4 4

/ 4 R T4 T9

Reaction echanism


82/92

Elementar* reactions

Reaction echanismElementar* reactions

1Elementar* reactions

REACTIONECHANI,

Uni$

moleclar

!i$

moleclar

Trimoleclar

0ast,lo4

Reactionintermediate

Rate –determining step

Reaction -echanisms


83/92

#e "era** /r"ress "$ a %#e&%a* rea%t&"' %a' e re/rese'te

at t#e "*e%*ar *ee* a ser&es "$ s&/*e elementary steps

"r elementary reactions.

#e se5e'%e "$ elementar' steps t#at *eas t" /r"%t

$"rat&"' &s t#e reaction mechanism.

2O (g ) + O2 (g ) 2O2 (g )

@*ee'tar ste/9 O + O 2O2

@*ee'tar ste/9 2O2 + O2 2O2Oera** rea%t&"'9 2O + O2 2O2

+

84Intermediates are s/e%&es t#at a//ear &' a rea%t&"'

# & 3 t t & t# ** * t&


84/92

@*ee'tar ste/9 O + O 2O2

@*ee'tar ste/9 2O2 + O2 2O2

Oera** rea%t&"'9 2O + O2 2O2

+

e%#a'&s 3ut not &' t#e "era** a*a'%e e5at&"'.

A' intermediate &s a*as $"re &' a' ear* e*ee'tar ste/

a' %"'se &' a *ater e*ee'tar ste/.

#e molecularity of a reaction &s t#e 'er "$ "*e%*es

rea%t&' &' a' e*ee'tar ste/.

• Unimolecular reaction e*ee'tar ste/ &t# 1 "*e%*e• Bimolecular reaction e*ee'tar ste/ &t# 2 "*e%*es

• Termolecular reaction e*ee'tar ste/ &t# 3 "*e%*es

85

Rate as a' @*ee'tar Ete/s


85/92

:r&t&' rea%t&"' e%#a'&ss9

• #e s "$ t#e e*ee'tar ste/s must &e t#e "era**

a*a'%e e5at&"' $"r t#e rea%t&"'.

•

#e rate-eter&'&' ste/ s#"* /re&%t t#e sae rate*a t#at &s eter&'e e/er&e'ta**.

#e rate-determining step &s t#e slowest ste/ &' t#e

se5e'%e "$ ste/s *ea&' t" /r"%t $"rat&"'.

#e e/er&e'ta* rate *a $"r t#e rea%t&"' etee' O


86/92

#e e/er&e'ta* rate *a $"r t#e rea%t&"' etee' O2

a' CO t" /r"%e O a' CO2 &s rate = k [O2]2. #e

rea%t&"' &s e*&ee t" "%%r &a t" ste/s9

Ete/ 19 O2 + O2 O + O3

Ete/ 29 O3 + CO O2 + CO2

:#at &s t#e e5at&"' $"r t#e "era** rea%t&"';

O2+ CO O + CO2

:#at &s t#e &'tere&ate;

O3

:#at %a' " sa a"t t#e re*at&e rates "$ ste/s 1 a' 2;

rate = k [O2]2 &s t#e rate *a $"r ste/ 1 s"

ste/ 1 st e s*"er t#a' ste/ 2

13.


87/92


88/92


89/92


90/92


91/92


92/92

Thank *oJ

chapter 2 - chemical kinetics.pptx

Documents