Low reaction rate. This process is slow

Reaction rate

Iron 

rusting

 has a low reaction rate. This process is slow.

Wood 

combustion

 has a high reaction rate. This process is

fast.

The reaction rate or rate of reaction is the

speed at which reactants are converted into

products. For example, the oxidative rusting of

iron under Earth's atmosphere is a slow reaction

that can take many years, but the combustion

of cellulose in a fire is a reaction that takes

place in fractions of a second. For most

reactions, the rate decreases as the reaction

proceeds.

Chemical kinetics is the part of physical

chemistry that studies reaction rates. The

concepts of chemical kinetics are applied in

many disciplines, such as chemical

engineering

[1][2][3][4]

, enzymology and

environmental engineering

[5][6][7]

.

Consider a typical chemical reaction:

The lowercase letters (a, b, p, and q) represent

stoichiometric coefficients, while the capital

letters represent the reactants (A and B) and

the products (P and Q).

According to IUPAC's Gold Book definition

[8]

 the

reaction rate v for a chemical reaction occurring

in a closed system under isochoric conditions,

without a build-up of reaction intermediates, is

defined as:

Formal definition

where [X] denotes the concentration of the

substance X (= A, B, P or Q). Reaction rate thus

defined has the units of mol/L/s.

The rate of a reaction is always positive. A

negative sign is present to indicate that the

reactant concentration is decreasing. The

IUPAC

[8]

 recommends that the unit of time

should always be the second. The rate of

reaction differs from the rate of increase of

concentration of a product P by a constant

factor (the reciprocal of its stoichiometric

number) and for a reactant A by minus the

reciprocal of the stoichiometric number. The

stoichiometric numbers are included so that the

defined rate is independent of which reactant or

product species is chosen for

measurement.

[9]:349

 For example, if a = 1 and b

= 3 then B is consumed three times more

rapidly than A, but v = -d[A]/dt = -(1/3)d[B]/dt is

uniquely defined. An additional advantage of

this definition is that for an elementary and

irreversible reaction, v is equal to the product of

probability of overcoming the transition state

activation energy and the number of times per

second the transition state is approached by

reactant molecules. When so defined, for an

elementary and irreversible reaction, v is the

rate of successful chemical reaction events

leading to the product.

The above definition is only valid for a single

reaction, in a closed system of constant volume,

an assumption which should be stated explicitly

in the definition. If water is added to a pot

containing salty water, the concentration of salt

decreases, although there is no chemical

reaction.

For an open system, the full mass balance must

be taken into account:

in − out + generation − consumption = accumulation

,

where F

A0

 is the inflow rate of A in molecules

per second, F

A

 the outflow, and v is the

instantaneous reaction rate of A (in number

concentration rather than molar) in a given

differential volume, integrated over the entire

system volume V at a given moment. When

applied to the closed system at constant

volume considered previously, this equation

reduces to:

,

where the concentration [A] is related to the

number of molecules N

A

 by [A] = 

N

A

N

0

V

. Here N

0

 is

the Avogadro constant.

For a single reaction in a closed system of

varying volume the so-called rate of conversion

can be used, in order to avoid handling

concentrations. It is defined as the derivative of

the extent of reaction with respect to time.

Here ν

i

 is the stoichiometric coefficient for

substance i, equal to a, b, p, and q in the typical

reaction above. Also V is the volume of reaction

and C

i

 is the concentration of substance i.

When side products or reaction intermediates

are formed, the IUPAC

[8]

 recommends the use

of the terms the rate of increase of

concentration and rate of the decrease of

concentration for products and reactants,

properly.

Reaction rates may also be defined on a basis

that is not the volume of the reactor. When a

catalyst is used the reaction rate may be stated

on a catalyst weight (mol g

−1

 s

−1

) or surface

area (mol m

−2

 s

−1

) basis. If the basis is a

specific catalyst site that may be rigorously

counted by a specified method, the rate is given

in units of s

−1

 and is called a turnover

frequency.

Factors that influence the reaction rate are the

nature of the reaction, concentration, pressure,

reaction order, temperature, solvent,

electromagnetic radiation, catalyst, isotopes,

surface area, stirring, and diffusion limit. Some

reactions are naturally faster than others. The

number of reacting species, their physical state

(the particles that form solids move much more

slowly than those of gases or those in solution),

the complexity of the reaction and other factors

can greatly influence the rate of a reaction.

Influencing factors

Reaction rate increases with concentration, as

described by the rate law and explained by

collision theory. As reactant concentration

increases, the frequency of collision increases.

The rate of gaseous reactions increases with

pressure, which is, in fact, equivalent to an

increase in concentration of the gas. The

reaction rate increases in the direction where

there are fewer moles of gas and decreases in

the reverse direction. For condensed-phase

reactions, the pressure dependence is weak.

The order of the reaction controls how the

reactant concentration (or pressure) affects

reaction rate.

Usually conducting a reaction at a higher

temperature delivers more energy into the

system and increases the reaction rate by

causing more collisions between particles, as

explained by collision theory. However, the main

reason that temperature increases the rate of

reaction is that more of the colliding particles

will have the necessary activation energy

resulting in more successful collisions (when

bonds are formed between reactants). The

influence of temperature is described by the

Arrhenius equation. For example, coal burns in

a fireplace in the presence of oxygen, but it

does not when it is stored at room temperature.

The reaction is spontaneous at low and high

temperatures but at room temperature its rate

is so slow that it is negligible. The increase in

temperature, as created by a match, allows the

reaction to start and then it heats itself,

because it is exothermic. That is valid for many

other fuels, such as methane, butane, and

hydrogen.

Reaction rates can be independent of

temperature (non-Arrhenius) or decrease with

increasing temperature (anti-Arrhenius).

Reactions without an activation barrier (e.g.,

some radical reactions), tend to have anti

Arrhenius temperature dependence: the rate

constant decreases with increasing

temperature.

Many reactions take place in solution and the

properties of the solvent affect the reaction

rate. The ionic strength also has an effect on

reaction rate.

Electromagnetic radiation is a form of energy.

As such, it may speed up the rate or even make

a reaction spontaneous as it provides the

particles of the reactants with more energy.

This energy is in one way or another stored in

the reacting particles (it may break bonds,

promote molecules to electronically or

vibrationally excited states...) creating

intermediate species that react easily. As the

intensity of light increases, the particles absorb

more energy and hence the rate of reaction

increases. For example, when methane reacts

with chlorine in the dark, the reaction rate is

slow. It can be sped up when the mixture is put

under diffused light. In bright sunlight, the

reaction is explosive.

The presence of a catalyst increases the

reaction rate (in both the forward and reverse

reactions) by providing an alternative pathway

with a lower activation energy. For example,

platinum catalyzes the combustion of hydrogen

with oxygen at room temperature.

The kinetic isotope effect consists in a different

reaction rate for the same molecule if it has

different isotopes, usually hydrogen isotopes,

because of the relative mass difference

between hydrogen and deuterium. In reactions

on surfaces, which take place for example

during heterogeneous catalysis, the rate of

reaction increases as the surface area does.

That is because more particles of the solid are

exposed and can be hit by reactant molecules.

Stirring can have a strong effect on the rate of

reaction for heterogeneous reactions.

Some reactions are limited by diffusion. All the

factors that affect a reaction rate, except for

concentration and reaction order, are taken into

account in the reaction rate coefficient (the

coefficient in the rate equation of the reaction).

For a chemical reaction a A + b B 

→ p P + q Q,

the rate equation or rate law is a mathematical

expression used in chemical kinetics to link the

rate of a reaction to the concentration of each

reactant. It is often of the type:

Rate equation

For gas phase reaction the rate is often

alternatively expressed by partial pressures.

In these equations k(T) is the reaction rate

coefficient or rate constant, although it is not

really a constant, because it includes all the

parameters that affect reaction rate, except for

concentration, which is explicitly taken into

account. Of all the parameters influencing

reaction rates, temperature is normally the most

important one and is accounted for by the

Arrhenius equation.

The exponents n and m are called reaction

orders and depend on the reaction mechanism.

For elementary (single-step) reactions the order

with respect to each reactant is equal to its

stoichiometric coefficient. For complex

(multistep) reactions, however, this is often not

true and the rate equation is determined by the

detailed mechanism, as illustrated below for the

reaction of H

2

 and NO.

For elementary reactions or reaction steps, the

order and stoichiometric coefficient are both

equal to the molecularity or number of

molecules participating. For a unimolecular

reaction or step the rate is proportional to the

concentration of molecules of reactant, so that

the rate law is first order. For a bimolecular

reaction or step, the number of collisions is

proportional to the product of the two reactant

concentrations, or second order. A termolecular

step is predicted to be third order, but also very

slow as simultaneous collisions of three

molecules are rare.

By using the mass balance for the system in

which the reaction occurs, an expression for the

rate of change in concentration can be derived.

For a closed system with constant volume,

such an expression can look like

Example of a complex reaction:

hydrogen and nitric oxide

For the reaction

the observed rate equation (or rate expression)

is: 

…

As for many reactions, the experimental rate

equation does not simply reflect the

stoichiometric coefficients in the overall

reaction: It is third order overall: first order in H

2

and second order in NO, even though the

stoichiometric coefficients of both reactants

are equal to 2.

[10]

In chemical kinetics, the overall reaction rate is

often explained using a mechanism consisting

of a number of elementary steps. Not all of

these steps affect the rate of reaction; normally

the slowest elementary step controls the

reaction rate. For this example, a possible

mechanism is:

1. 

2. 

3. 

Reactions 1 and 3 are very rapid compared to

the second, so the slow reaction 2 is the rate

determining step. This is a bimolecular

elementary reaction whose rate is given by the

second order equation:

,

where k

2

 is the rate constant for the second

step.

However N

2

O

2

 is an unstable intermediate

whose concentration is determined by the fact

that the first step is in equilibrium, so that

[N

2

O

2

] = K

1

[NO]

2

, where K

1

 is the equilibrium

constant of the first step. Substitution of this

equation in the previous equation leads to a

rate equation expressed in terms of the original

reactants

This agrees with the form of the observed rate

equation if it is assumed that k = k

2

K

1

. In

practice the rate equation is used to suggest

possible mechanisms which predict a rate

equation in agreement with experiment.

The second molecule of H

2

 does not appear in

the rate equation because it reacts in the third

step, which is a rapid step after the rate-

determining step, so that it does not affect the

overall reaction rate.

Temperature dependence

Each reaction rate coefficient k has a

temperature dependency, which is usually given

by the Arrhenius equation:

E

a

 is the activation energy and R is the gas

constant. Since at temperature T the molecules

have energies given by a Boltzmann

distribution, one can expect the number of

collisions with energy greater than E

a

 to be

proportional to e

 

−E

a

⁄

RT

. A is the pre-exponential

factor or frequency factor.

The values for A and E

a

 are dependent on the

reaction. There are also more complex

equations possible, which describe temperature

dependence of other rate constants that do not

follow this pattern.

Temperature Is a measure of the average

kinetic energy of the reactants. As temperature

increases this makes the kinetic energy of the

reactants increase too meaning they move

faster. With the reactants moving faster this

allows more collisions to take place at a greater

speed so the chance of reactants forming into

products increases which in turn results in the

rate of reaction increasing. A rise in 10 degrees

Celsius results in around double the reaction

rate. The minimum kinetic energy required for a

reaction to occur is called the activation energy

and this is denoted by Ea. The activated

complex shown on the diagram below is the

energy barrier that must be overcome when

changing reactants into products. On a energy

distribution graph we can see which molecules

have enough energy to react. The molecules

past the dotted line have a greater or equal

energy level to the activation energy and so can

react For a successful collision to take place

the collision geometry must be right meaning

the reactant molecules have to be facing the

right way so that the activated complex can be

formed.

A chemical reaction takes place only when the

reacting particles collide. However, not all

collisions are effective in causing the reaction.

Products are formed only when the colliding

particles possess a certain minimum energy

called threshold energy. As a rule of thumb,

reaction rates for many reactions double for

every 10 degrees Celsius increase in

temperature,

[11]

 For a given reaction, the ratio of

its rate constant at a higher temperature to its

rate constant at a lower temperature is known

as its temperature coefficient (Q). Q

10

 is

commonly used as the ratio of rate constants

that are 10 °C apart.

The pressure dependence of the rate constant

for condensed-phase reactions (i.e., when

reactants and products are solids or liquid) is

usually sufficiently weak in the range of

pressures normally encountered in industry that

it is neglected in practice.

The pressure dependence of the rate constant

is associated with the activation volume. For

Pressure dependence

the reaction proceeding through an activation-

state complex:

A + B 

⇌ |A⋯B|

‡

 

→ P

the activation volume, ΔV

‡

, is:

where V̄ denotes the partial molar volume of a

species and ‡ indicates the activation-state

complex.

For the above reaction, one can expect the

change of the reaction rate constant (based

either on mole fraction or on molar

concentration) with pressure at constant

temperature to be:

[9]:390

In practice, the matter can be complicated

because the partial molar volumes and the

activation volume can themselves be a function

of pressure.

Reactions can increase or decrease their rates

with pressure, depending on the value of ΔV

‡

.

As an example of the possible magnitude of the

pressure effect, some organic reactions were

shown to double the reaction rate when the

pressure was increased from atmospheric

(0.1 MPa) to 50 MPa (which gives

ΔV

‡

 = −0.025 L/mol).

[12]

See also

Rate of solution

Dilution (equation)

Diffusion-controlled reaction

Steady state approximation

Collision theory and transition state are

chemical theories that attempt to predict and

explain reaction rates.

Isothermal microcalorimetry

1. Silva, Camylla K. S.; Baston, Eduardo P.;

Melgar, Lisbeth Z.; Bellido, Jorge D. A.

(2019-10-01). "Ni/Al2O3-La2O3 catalysts

synthesized by a one-step polymerization

method applied to the dry reforming of

methane: effect of precursor structures of

nickel, perovskite and spinel". Reaction

Notes

Kinetics, Mechanisms and Catalysis. 128

(1): 251–269. doi:10.1007/s11144-019-

01644-3 . ISSN 1878-5204 .

2. Kinetic studies of propane oxidation on Mo

and V based mixed oxide catalysts . 2011.

3. Naumann d'Alnoncourt, Raoul; Csepei,

Lénárd-István; Hävecker, Michael; Girgsdies,

Frank; Schuster, Manfred E.; Schlögl, Robert;

Trunschke, Annette (2014). "The reaction

network in propane oxidation over phase-

pure MoVTeNb M1 oxide catalysts" . J.

Catal. 311: 369–385.

doi:10.1016/j.jcat.2013.12.008 .

hdl:11858/00-001M-0000-0014-F434-5 .

4. Elizalde, Ignacio; Mederos, Fabián S.; del

Carmen Monterrubio, Ma.; Casillas, Ninfa;

Díaz, Hugo; Trejo, Fernando (2019-02-01).

"Mathematical modeling and simulation of

an industrial adiabatic trickle-bed reactor

for upgrading heavy crude oil by

hydrotreatment process". Reaction Kinetics,

Mechanisms and Catalysis. 126 (1): 31–48.

doi:10.1007/s11144-018-1489-7 .

ISSN 1878-5204 .

5. Liu, Jiaqi; Shen, Meiqing; Li, Chenxu; Wang,

Jianqiang; Wang, Jun (2019-10-01).

"Enhanced hydrothermal stability of a

manganese metavanadate catalyst based

on WO3–TiO2 for the selective catalytic

reduction of NOx with NH3". Reaction

Kinetics, Mechanisms and Catalysis. 128

(1): 175–191. doi:10.1007/s11144-019-

01624-7 . ISSN 1878-5204 .

6. Li, Xiaoliang; Feng, Jiangjiang; Xu, Zhigang;

Wang, Junqiang; Wang, Yujie; Zhao, Wei

(2019-10-01). "Cerium modification for

improving the performance of Cu-SSZ-13 in

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ISSN 1878-5204 .

7. Vedyagin, Aleksey A.; Stoyanovskii, Vladimir

O.; Kenzhin, Roman M.; Slavinskaya, Elena

M.; Plyusnin, Pavel E.; Shubin, Yury V. (2019-

06-01). "Purification of gasoline exhaust

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8. IUPAC, Compendium of Chemical

Terminology, 2nd ed. (the "Gold Book")

(1997). Online corrected version:  (2006–)

"Rate of reaction ".

doi:10.1351/goldbook.R05156

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Chemistry. Benjamin/Cummings. ISBN 0-

8053-5682-7.

Chemical kinetics, reaction rate, and order

(needs flash player)

Reaction kinetics, examples of important rate

laws  (lecture with audio).

10. Laidler, K. J. (1987). Chemical Kinetics (3rd

ed.). Harper & Row. p. 277.

ISBN 0060438622.

11. Connors, Kenneth (1990). Chemical

Kinetics:The Study of Reaction Rates in

Solution. VCH Publishers. p. 14. ISBN 978-0-

471-72020-1.

12. Isaacs, Neil S. (1995). "Section 2.8.3" .

Physical Organic Chemistry (2nd ed.).

Harlow: Addison Wesley Longman.

ISBN 9780582218635.

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otherwise noted.

Rates of reaction

Overview of Bimolecular Reactions

(Reactions involving two reactants)

pressure dependence  Can. J. Chem.

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