Phenomenon that changes reaction rate

INTRODUCTION

For a long time, chemistry has been interested in studying the phenomenon of changing the rate of reaction when a very small amount of a substance is present. Perhaps this phenomenon comes from very coincidental things.

Around the beginning of the 18th century, Russian scientist MA Ilinski studied to prepare aromatic sulfuric acid (an intermediate product for synthesizing dyes) from the organic compound anthraquinone C 6 H 4 (CO) 2 C 6 H 4 . According to his calculations, anthraquinone when heated at 100 o C with sulfuric acid H 2 SO 4 will form sulfuric acid with a definite structure. He conducted many experiments but still failed. One day, while he was conducting the experiment, the thermometer broke, a drop of mercury fell into the flask. And like a miracle, sulfuric acid was formed in the flask. This means that the drop of mercury directed the process in the desired direction. It is difficult to say whether this story is reliable or not, but one thing is clear, a small amount of impurity - mercury - has a clear effect on the reaction, which means that Hg catalyzed the reaction.

Also in the early 18th century, the British scientist David conducted an experiment that caught the attention of scientists in many countries. He blew a mixture of CH 4 and air into a heated Pt wire, and saw that the Pt wire became red hot in the mixture and continued to be red hot for a long time. Many times he took the wire out to cool in the air and then put it back into the gas mixture, the Pt wire became red hot and glowed again. The Pd wire also showed the same phenomenon, but Cu, Ag, Fe... did not. It turned out that Pt and Pd accelerated the oxidation reaction of methane with oxygen in the air, which means they were catalysts. CH 4 burned into CO 2 and H 2 O, releasing a large amount of heat, causing the temperature of the metal to increase and the metal to glow.

Nearly 300 years have passed since the invention of the CH 4 combustion reaction on Pt, and until now the catalyst has not lost its value. During World War I and II, Russian scientists applied this reaction by filling a mesh shell with Pt-impregnated asbestos fibers and holding the shell above a small tank containing gasoline. When gasoline vapor penetrates the Pt, it will gradually oxidize into CO 2 and H 2 O. This chemical process releases a lot of heat, causing the asbestos fibers to heat up and radiate heat. Thanks to such equipment, Soviet soldiers were saved from freezing during the harsh winter days of the Great Patriotic War.

And many catalytic experiments were studied, clarifying the nature of the catalytic effect of many substances. In 1836, Swedish scientist Berselius first introduced the term "catalysis" into science.

So what is the phenomenon of catalysis? The phenomenon of catalysis is to increase the rate of reaction under the influence of a substance, that substance is called a catalyst. The catalyst forms an intermediate compound with the reactant. Finally, the catalyst is reduced (ie there is no change in the chemical aspect). That phenomenon is called catalysis and the reaction is called a catalytic reaction.

If the catalyst is not regenerated, it is called " accelerator ". For example, the vulcanization process of rubber (rubber combined with S): when adding Na, the vulcanization speed increases and at the end of the process, Na is in the rubber. So Na is an accelerator for the vulcanization process of rubber.

The catalyst after participating in the process does not change chemically but can change physical properties (such as changing shape: from granular to fine dust...)

The influence of catalysts is very strong and under their influence, the reaction rate can increase hundreds of times, thousands of times and more. Catalysts can stimulate reactions that would not actually occur without them under certain conditions of investigation. Many chemicals react very slowly; for the reaction to occur, it is necessary to proceed at very high temperatures and pressures. If we wait for the reaction under normal conditions, it will take a long time, not hours but days, months. Such processes are not suitable for industry. But if we use catalysts, the reaction becomes completely possible at low temperatures and pressures. That is, catalysts increase the reaction rate and reduce the activation energy. For example, a mixture of pure CO and O 2 does not react even when heated, but if a very small amount of manganese dioxide MnO 2 is added , all the CO is converted very quickly to CO 2 .

In addition to the properties of accelerating reaction rate and reducing activation energy, catalysts also have selectivity, directing the process towards the main reaction, reducing the rate of side reactions, and increasing the yield of the main product.

water

For example, isopropyl alcohol can be converted to acetone and hydrogen, or to propylene and

CH 3 COCH 3 + H 2 (1)

C3H7OH

C 3 H 6 + H 2 O (2)

* If the catalyst is ZnO: the reaction occurs mainly in direction (1)

* If the catalyst is Al 2 O 3 : the reaction occurs mainly in direction (2)

Normally, a catalyst only acts for one reaction; in particular, enzyme catalysts only act for one or several stages in a reaction; but there are also catalysts that are active for several reaction groups; for example, acid catalysts act for cracking, isomerization, hydrolysis, dehydration, alkylation, etc.

Catalysts are used in many different forms, they can be a complex mixture of many oxides such as zeolites, clays, aluminosilicates...; or a pure substance such as metal catalysts Ag, Cu, Pt...; or a simple compound such as oxides, sulfur...; or in the form of a complex compound such as yeast catalyst.

Because of its many advantages, most reactions in chemical engineering, especially in the field of petrochemical refining and organic synthesis, use catalysts. Currently, all modern petrochemical refineries use catalytic cracking, catalytic reforming, etc. instead of the previous thermal cracking and thermal reforming processes.

To effectively support the search for new catalysts, researchers have combined physical methods with kinetic methods. So the task of kinetics is to study the rate of chemical reactions, factors that affect the reaction rate such as concentration of reactants, temperature, pressure ... and the reaction mechanism when there is the participation of catalysts.

To explain the phenomena of catalysis, there are still many things that are not clearly understood, but people have determined the main features of the phenomenon. Catalysis plays a great role in chemistry. Deep penetration into the nature of catalysis, the creation of theoretical bases, allowing to predict the effect of this or that catalyst on given chemical processes, will give people the tools to enrich the material basis for humanity better and better.

CHAPTER I: HOMOGENEOUS CATALYTIC REACTIONS

I. Concept

 Homogeneous catalyst is a catalyst in the same phase as the reactants.

 Homogeneous catalytic reactions occur only in the gas and liquid phases; there is no homogeneous catalysis in the solid phase.

For example:

1) Gas phase : SO 2 oxidation reaction with NO catalyst to form SO 2 to produce sulfuric acid

industrial

SO 2 + O 2

NOSO  H SO

3 2 4

The reaction occurs through the following stages:

O 2 + 2 NO  2 NO 2

2 SO 2 + 2 NO 2  2 SO 3 + 2 NO

2 SO 2 + O 2  2 SO 3

In which: NO 2 is an intermediate compound.

2) Liquid phase : Homogeneous catalytic reactions in the liquid phase are mostly acid-base catalytic reactions. For example, the oxidation reaction of thiosulfate ion with H 2 O 2 with I - ion as catalyst.

2 S 2 O 3 2- + H 2 O 2 + 2 H + I

The reaction occurs in three stages as follows:

S 4 O 6 2- + 2 H 2 O

H 2 O 2 + I -  IO - + H 2 O I - + IO - + 2 H +  I 2 + H 2 O

I 2 + 2 S 2 O 3 2-  S 4 O 6 2- + 2 I -

2 S 2 O 3 2- + H 2 O 2 + 2H +  S 4 O 6 2- + 2 H 2 O

In which: IO - and I 2 are intermediate compounds

 Homogeneous autocatalytic reactions: usually reactions that occur in an H + environment

For example:

CH 3 COOC 2 H 5 + H 2 O

CH 3 COOH + C 2 H 5 OH

1) Reaction with catalyst production : Ester hydrolysis reaction in acidic environment

The first stage requires more acid as a catalyst, but later on, acetic acid is produced as a catalyst.

2) Autocatalytic reaction with reactant acting as catalyst : esterification reaction

C6H5COOH + C2H5OH

H C 6 H 5 COOC 2 H 5 + H 2 O

The catalyst for this reaction is H + ion but here the starting medium is acid so it also acts as a catalyst.

II. Spitalski - Kodozeb's homogeneous catalysis theory

In 1926, Spitalski proposed the theory of homogeneous catalysis as follows:

1) There is a HCTG formation stage between the catalyst and the reactant.

For example: Oxidation reaction of H 3 PO 3 to a. H 3 PO 4 with oxidizing agent K 2 S 2 O 8 on catalyst HI.

H 3 PO 3 + K 2 S 2 O 8 + H 2 O

HI H 3 PO 4 + K 2 SO 4 + H 2 SO 4

Watch the process to see purple color appear and when the process ends, the purple color disappears. This purple color is due to the formation of HCTG I 2 .

K 2 S 2 O 8 + 2 HI I 2 + K 2 SO 4 + H 2 SO 4

H 3 PO 3 + I 2 + H 2 O H 3 PO 4 + 2 HI

2) The process of forming intermediate products is reversible and occurs at a fairly fast rate because of the effect of the catalyst, and this rate does not depend on the nature of the HCTG.

3) The unstable HCTG will decompose relatively slowly to the reaction products and release the catalyst.

The overall rate of the process depends mainly on the rate of HCTG decomposition:

v c = f (v HCTG decomposition )

4) The formation of HCTG is due to the combination of reactant molecules or active groups of reactant molecules with active groups of catalyst molecules.

5) The catalytic reaction creates many HCTGs with different activities and the decomposition of HCTGs occurs differently.

Example: Decomposition reaction of H 2 O 2

HCTG
Active compound		Less active compound
MoO42-	MoO 8 2-	MoO 6 2-	MoO52-
		 average activity
WO 4 2-	WO 8 2-		2- WO 5

Maybe you are interested!

Catalysis

6) Homogeneous catalytic reactions in the presence of a catalyst will reduce the activation energy of the reaction, thereby increasing the value of the rate constant k and leading to an increase in the reaction rate under the same conditions (compared to when the catalyst is not present).

Reaction: E

A + B X C

k 1

k 2

A + X [AX] B

[ABX] C + X

E o : energy of mixture A + BE 1 : energy of product C Sugar (1): non-catalyzed reaction

(2): catalytic reaction

 E I : activation energy for non-catalyzed reaction

E O

 E 1

 E I1

 E 2

 E II : activation energy for catalytic reaction

 E II =  E 1 (if  E 1 >  E 2 )

 E 2 (if  E 2 >  E 1 )

reaction direction

For uncatalyzed reactions, the Arrhenius equation takes the form:

k kxt = z 1 . e -  E I /RT

For a catalyzed reaction, the Arrhenius equation takes the form:

k xt = z 2 . e -  E II /RT

in which: k kxt , k xt : rate constant of uncatalyzed and catalyzed reactions. If z 1  z 2 we have:

k xt k k xt

= e  E/RT with  E =  E I

-  E II

The activation energy of a catalyzed reaction is reduced by about 10,000 cal/mol or more compared to a non-catalyzed reaction.

If the reaction occurs at 300K, when replacing the value with numbers we get:

e  E/RT = e 10 000/ 1.987. 300  2.0. 10 8

That is, the catalytic reaction occurs hundreds of millions of times faster than the non-catalytic reaction.

7) Some examples to calculate the kinetic equation of a reaction 1/ The reaction has the form:

n AXC

k 1

Reaction process:

n A + X

k 2

Z : HCTG k 3

C + X

Calculate the overall velocity of the reaction v c : Because v c mainly depends on the decomposition reaction of HCTG, we have:

v c = k 3 . C z (1)

where C z : concentration of HCTG is calculated through the equilibrium constant of the reaction to form HCTG is K

K = n C Z

= n C Z

(2)

C A . C X balance

C A . (C X o - C Z )

From (2):

(C X o : initial concentration of catalyst)

K. C A n . C Xo - K. C A n . C Z = C Z

C Z =

K. C A n . C Xo

K. C A n + 1

From (1):

v c = k 3

K. C A n . C Xo

K. C A n + 1

Consider 2 cases:

*1. K is very large: that is, the reaction quickly reaches equilibrium. Then: K. C A n >> 1  K. C A n + 1  K. C A n

v c

k 3

K . C n . C

AX o k . CK . C n 3 X o

 v c = f(C Xo )

Note: the overall reaction rate does not depend on the concentration of the reactant but only depends on the concentration of the catalyst in the case where the reaction occurs in the direction of forming a large amount of HCTG.

*2. K is very small: meaning the reaction takes a long time to reach equilibrium. Then: K. C A n << 1  K. C A n + 1  1

 v c = k 3 . K. C A n . C Xo

 v c = f(C A , C Xo )

Note: the overall reaction rate depends not only on the concentration of the reactant but also on the concentration of the catalyst in the case where the reaction occurs in the direction of forming less HCTG.

k 1k 2

2/ Reaction in H + environment :

n A + X + H + Z : HCTG

k 3 C + X + H +

Kinetic equation of the reaction:

v c = k 3 . C Z

Equilibrium constant of HCTG formation stage:

K = n

C Z

C A . C X . C H+

C Z

= n