Chemistry Project

Term paper of ChemistrySubmitted to:-Mr. Kultar Singh

Submitted By:-Aakashdeep singh

Class:-B.Tech.(E.C.E.)Hons.

Roll no.:-R261B49

Reg.No. :-10802459

ACKNOWLEDGEMENT :-

I would to show my gratitude towards the help and guidance rendered tome by my

teacherMr.Kultar Singh. I would also like to acknowledge the support and help I received from my family and friends, for completing this term paper.

Contents:-

Rate of Reaction

Factors affect the rate of reaction

Rate law and reaction order

First order reaction

Example:-

RATES OF REACTIONRate of reaction:-Chemical reactions require varying lengths of time for completion, depending upon the characteristics of the reactants and products and the conditions under which the reaction is taking place. Chemical Kinetics is the study of reaction rates, how reaction rates change under varying conditions and by which mechanism the reaction proceeds.

What factors affect the rate of a reaction?We've already said that the characteristics of the reactants affect the rate of the reaction, what we want to do here is see what physical factors affect the rate. We can list these;

1. The concentration of the reactants. The more concentrated the faster the rate (note in some cases the rate may be unaffected by the concentration of a particular reactant provided it is present at a minimum concentration). Remember for gasses, increasing the pressure simply increases the concentration so that's the same thing.

2. Temperature. Usually reactions speed up with increasing temperature ("100C rise doubles rate").

3. Physical state of reactants. Powders react faster than blocks - greater surface area and since the reaction occurs at the surface we get a faster rate.

4. The presence (and concentration/physical form) of a catalyst (or inhibitor). A catalyst speeds up a reaction, an inhibitor slows it down. Check your class notes - there should be a graph with the effect of various catalysts upon the decomposition of hydrogen peroxide; draw it in below.

5. Light. Light of a particular wavelength may also speed up a reaction - refer to your organic chemistry notes for reaction of alkanes with halogens.

We've used the term reaction rate, so we should define what we mean by this; the reaction rate is the increase in molar concentration of product of a reaction per unit time or the decrease in molar concentration of reactant in unit time. The question is, how does reaction rate vary with time? Consider the decomposition of N2O5;

2 N2O5 > NO2 + O2

The table shows the concentration of N2O5 as a function of time (@450C);

Time / min [N2O5] / moldm-3 0 0.0175620 0.00933

40 0.0053160 0.0029580 0.00167100 0.00094160 0.00014

Plot this data.

To find the rate we divide the change in concentration of N2O5 by the time period over which this change occurred. Because N2O5 is consumed in this reaction, this will be a negative figure, by convention, rate data are given as positive numbers:

If we calculate the rate over the first twenty minutes we get;

Rate = 8.31 x 10-4 moldm-3min-1. That is the change in concentration from 0 to 20 minutes / 20 minutes. Now you try for the other times: (you should get 2.01 x 10 -4, 1.18 x 10-4, 0.64 x 10-4 and 0.37 x 10-4moldm-3min-1).

Why should it slow down? - the rate slows down as the concentration of the reactant decreased.

Rate Law and Reaction Order:-If we examine the effect on the rate of a reaction by changing the initial concentration of reactants we may be able to derive the rate law and hence the reaction order. Consider the following reaction;

2NO2 + F2 > 2NO2F

If the concentration of NO2 is doubled then the rate is doubled, likewise when the concentration of the fluorine is doubled the rate doubles and so we get the following rate law;

Rate = k [NO2][F2]

The rate law is an equation that relates the rate of a reaction to the concentration of reactants raised to various powers.

For the reaction above both reactant concentrations have an exponent of 1.

The rate constant, k, is a proportionality constant in the relationship between rate and concentrations. This has a fixed value at any given temperature but varies with temperature.

For a general reaction;

aA + bB > cC + dD

we would have a rate law;

Rate = k [A]m[B]n

Where m and n are determined experimentally (they are usually, but not always, integers). The values of m and n are the orders of reaction, hence the reaction is mth order with respect to A and nth order with respect to B giving a (m+n)th order overall. So if m = 1 and n = 2 the reaction would be first order with respect to A, second order with respect to B and third order overall. If we doubled the concentration of A for this reaction then we would double the rate but doubling the concentration of B would quadruple the rate. (Since 2nd order wrt to B.... work it out if you take [2[2 what do you get ? now take [4]2)

First Order Reactions:-Consider a first order reaction;

A > Products

The rate law is of the form;

Rate = k[A] / t

If we set a = the initial concentration of A, x = loss of A with time then (a-x) is the concentration of A at any given time.

Integration of the equation gives;

kt = In [a/(a-x)]

Plots of In [a/(a-x)] against t should give a straight line with gradient k

Plots of log10 [ a/ (a-x)] against t should give a straight line with gradient k/2.303

Plot In (a-x) against t give straight line with gradient -k

(Units of k are second-1)

For a first order reaction, the half life is dependant of the initial concentration and is the time for half of the initial concentration to have reacted. Since half has reacted then (a-x) must equal a/2 (i.e. half the initial concentration).

Since (a-x) = a/2, we can rewrite the equation above to give;

kt 1/2 = In 2

Where t 1/2is the half life.

Examples:-

The half - life for a first order reaction is 100 seconds. Calculate the rate constant and determine what fraction will have reacted after 250 seconds.

kt 1/2 = 0.693            k = 0.693 / t        k = 0.693 / 100

= 6.93 x 10-3s-1

Asked then to calculate x - how much has reacted. Set a, the initial concentration to 1.00

kt = In [ a/(a-x)]    In[a/(a-x)] = 6.93 x 10-3 x 250

= 1.7325

[a/(a-x)] = 5.6547     = 0.1768

x = 1 - 0.1768    x = 0.8232    (82.32% has reacted)

1860 years)

type of molecule such as  AB ® A + B or ABA ® BAA Bimolecular reactions involve a collision between two molecules such as A + B ® AB or AB + CD ® AC + BD Termolecular reactions involve a collision between three molecules. But the rate determining step in the reaction usually does not involve a mechanism of a simple uni-, bi-, or termolecular reaction. The order of the reaction, n, which must be evaluated experimentally, is important in determining the mechanism by which the reaction takes place. It is defined by the equation  dc/dt = kcn (1) where n is evaluated from the rate of change of concentration of reactant c with time. If n is 1, the reaction is first order, if it is 2, the reaction is second order, and if it is 3, the reaction is third order. If (as is usually the case) n is found to have other values that are not integers, the reaction is complex and involves more than one uni-, bi-, or tri-molecular reaction. Fortunately the rates of many unimolecular or bimolecular reactions can be estimated from molecular structure or other properties, and often a complex

reaction may be broken up into a series of predictable units molecular and bimolecular reactions. The first-order reaction equation -dc/dt = kc1 (2) is integrated to give- ln c = kt + constant (3) or  

kt t

c

c

1

2 1

1

2

ln(4)

 where c1 and c2 are the concentrations at times t1 and t2. For first-order reactions k is numerically equal to the fraction of the substance which reacts per unit time, usually expressed in reciprocal seconds (or minutes). in such reactions it is not necessary to know the initial concentration of the reactants or the absolute concentrations at various times. The concentrations may be determined directly by experiment using chemical or physical measurements; or any property, e.g., volume, electrical conductance, or light absorption, which is proportional to the concentration may be measured and substituted for c in formulas (3), (4), or (5).   The kinetics of a second-order reaction is described by the equation  -dc/dt = kcA

2 (5) where cA is the concentration of the reactant A, or dcA/dt = kcAcB (6) where cA and cB are the concentrations of two reactants A and B.  The numerical value of the rate constant k for a second-order reaction depends on the units in which the concentrations are expressed, such as moles per liter moles per cubic centimeter, or atmospheres. In a first order reaction these units cancel out, but in a second-order reaction they do not. In a second-order reaction, if one reactant is present in sufficiently large excess, its concentration remains essen tially constant and so the second-order reaction then appears to be of the first order.  

HYDROLYSIS OF METHYL ACETATE:- 

Apparatus:-25oC thermostat (L); 35oC thermostat (L); three 250 ml Erlenmeyers (D); two 125 ml Erlenmeyers (D); 5 ml pipette (D); 50 ml pipette (D); timer (S); Buret (S); Buret Clamp (S). 

Chemicals:- Two liters 0.2 N NaOH (P); phenolphthalein indicator (S); 500 ml 1 N HCl (L); Distilled H2O (L); ice (L); 25 ml methyl acetate (S). NOTE: It is important to initiate the first reaction no later than 30-45 minutes after the laboratory period has begun (the sooner the better). When preparing the two liters of 0.2 N NaOH, weigh out the necessary amount of NaOH pellets into a glass container (NO WEIGHING PAPER IS TO BE USED). CAUTION: NaOH pellets are CAUSTIC, use a spatula for transferring. NO HANDS. 

Procedures:-Two runs are made at 25oC during the first laboratory period (your instructor may ask that you do these runs at room temperature) and two runs are made at 35oC the second laboratory period The concentration of methyl acetate at a given time is determined through titration of samples with a standard sodium hydroxide solution; the experimental accuracy depends chiefly on the care used in pipetting and titrating. The sodium hydroxide solution used could be prepared by dilution of a saturated stock solution to minimize the amount of carbonate present and hence to reduce the fading of the phenolphthalein end point. It is not necessary, however, to use CO2-free distilled water, because the amount of carbonate introduced in air-saturated water is negligible when titrating with 0.2 N sodium hydroxide.  A test tube containing about 12 ml methyl acetate is set into a thermostat at 25° C. Approximately 250 ml of standardized 1 N hydrochloric acid is placed in a flask clamped in the thermostat. After thermal equilibrium has been reached (10 or 15 min should suffice), two or three 5-ml aliquots of the acid are titrated with the standard sodium hydroxide solution to determine the exact normality of the sodium hydroxide in terms of the standardized hydrochloric acid. Then 100 ml of acid is transferred to each of two 250-ml flasks clamped in the thermostat and 5 min allowed for the reestablishment of thermal equilibrium. Precisely 5 ml of methyl acetate is next transferred to one of the

flasks with a clean, dry pipette; the timing watch is started when the pipette is half emptied. The reaction mixture is shaken to provide thorough mixing.  A 5-ml aliquot is withdrawn from the flask as soon as possible and run into 50 ml of distilled water. This dilution slows down the reaction considerably, but the solution should be titrated at once; the error can be further reduced by chilling the water in an ice bath. The time at which the pipette has been half emptied into the water in the titration flask is recorded, together with the titrant volume. Additional samples are taken at 10-min intervals for an hour; then at 20-min intervals for the next hour and a half. A second determination is started about a quarters of an hour after the first one to provide a check experiment.  In similar fashion, two runs are made at a temperature of 35°. Because of the higher rate of reaction, three samples are first taken at 5-min intervals, then several at 10-min intervals, and a few at 20-min intervals. It is convenient to start the check determination about a half hour after the first experiment is begun. The 5 ml aliquots are best quenched by diluting in 50 ml of ice water! Use an ice-H2O slurry! Titrate them immediately with 0.2 N NaOH using two drops phenolphthalein as indicater. 

THEORY.:- The hydrolysis of methyl acetate presents several interesting aspects. The reaction, which is extremely slow in pure water, is catalyzed by hydrogen ion:

k1’CH3COOCH3 + H2O + H+ Û CH3COOH + CH3OH + H+ (7)

k2

 The reaction is reversible, so that the net rate of hydrolysis at any time is the difference between the rates of the forward and reverse reactions, each of which follows the simple rate law given by Eq. (7). Thus  

OHCHCOOHCHCOOCHCHOH

COOCHCHcckcck

dt

dc33332

33

2'1

(8) where k1’ is the rate constant for the forward reaction and k2 for the reverse reaction. For dilute solutions, water is present in such large excess that its concentration undergoes a negligible proportional change while that of the methyl acetate is changed considerably. For this case Eq. (8) may be written  

dc

dtk c k c c

CH COOCH

CH COOCH CH COOH CH OH3 3

3 3 3 31 2(9)

In the early stages of the hydrolysis, the concentrations of acetic acid and methanol remain small enough for the term involving them to be negligible, and the reaction appears to be of first order:  

33

33

1 COOCHCH

COOCHCHck

dt

dc

(10) The value of k1 can then be determined by one of the methods conventional for first order reactions.  Evaluation of kl at two different temperatures permits the calculation of the Arrhenius heat of activation DHa for the forward reaction:  d k

dT

H

RTaln 12

D

(11)  

ln,

,

k

k

H

R

T T

T TT

T

a1

1

2 1

2 1

2

1

D

(12) In obtaining the integrated form, it is assumed that DHa is a constant. The heat of activation is usually expressed in calories per mole and is interpreted as the amount of energy the molecules must have in order to be able to react.  A more accurate calculation of the influence of temperature may be made on the basis of the Eyring equation,  

kRT

N he e

o

S R H RT D D / /

(13) where No is Avogadro’s number, h is Planck’s constant, and DS‡ and DH‡ are the standard entropy and enthalpy changes for formation of the activated complex from the reactants   CH3COOCH3 + H2O + H+--------> [activated complex]  and is a constant, of the order of 1/2, defined as the probability that an activated complex will decompose to form product species (rather than regenerating reactant species). Thus DH‡ may be determined from measurements of k at two or more temperatures, on the assumption DS‡, DH‡, and are independent of temperature.  k

k

T

T

H

R T TT

T

1

1

2

1 2 1

2

1

1 1,

,

exp

D

(13) 

D HRT T

T T

k T

k TT

T

1 2

2 1

1 1

1 2

2

1

ln,

, (14) Although DS‡ cannot be determined from these data, for lack of knowledge of the value of , it is sometimes possible to gain some information about the magnitude of DS‡ by making a guess as to the value of . In ordinary cases, a value of 1/2 to 1 is considered a reasonable estimate, but under certain circumstances may be very small. The value of DH‡ can be used, of course, to calculate the value of kiT at any temperature (over the range in which DH‡ and DS‡ remain constant) from a knowledge of k1 at one temperature.  An explicit solution to the kinetic equation may also be written for the case where the reverse reaction cannot be ignored. If the concentration of methyl acetate is a moles per liter initially, and a - x moles per liter at time t, then Eq. (8) can be written as - d(a - x)/dt = dx/dt = kl(a - x) - k2x2 (15) since for each mole of methyl acetate hydrolyzed a mole of acetic acid and a mole of methanol are produced. Integration of this relation gives  

t

k ak k

a x ak k

a x ak k

1

4 1

2 4 1 1

2 4 1 11 2 1

1 2

2 1

1 2

2 1

1 2/

ln/

//

/

/

(16) Making use of the relation that the equilibrium constant Kh for the hydrolysis reaction is given by the expression  

Kc c

c c

k

k

k

k ch

CH COOH CH OH

CH COOCH H O H O

3 3

3 3 2 2

1

2

1

2

'

(17) one obtains 

tk a K c

a x a K c

a x a K ch H Oo

h H Oo

h H Oo

1

4 1

2 4 1 1

2 4 1 117

11 2

1 2

1 22

2

2

( / )ln

( / )

/( )/

/

/

 Here co

H2O represents the concentration of water present, which is treated as a constant in accordance with the assumption made in obtaining Eq. (8) from Eq. (7).   

CALCULATIONS:- The titrant volume at time t, Vt, measures the number of equivalents of hydrochloric acid and acetic acid then present in the 5ml reation mixture aliquot. Let V¥ represent what the titrant volume per 5-ml aliquot would be if the hydrolysis were complete. Then V¥ -Vt measures the number of equivalents of methyl acetate remaining per 5-ml aliquot at time t, because one molecule of acetic acid is produced for each molecule of methyl acetate hydrolyzed. The corresponding concentration of methyl acetate in moles per liter is N(V¥ - Vt)/5, where N is the normality of the sodium hydroxide solution.  If the reaction actually proceeded to completion, V¥ could be measured directly by titration of an aliquot from the equilibrium mixture. An appreciable amount of unhydrolyzed methyl acetate is present at equilibrium, however, so V¥ must be calculated. The volume of the solution initially formed on mixing the 100 ml of 1 N hydrochloric acid with 5 ml of methyl acetate is designated by Vs. At 25°C, Vs is 104.6 ml rather than 105 ml because the solution is not ideal. Let the number of milliliters of sodium hydroxide solution required to neutralize a 5-ml aliquot of the original 1 N hydrochloric acid be Vx. The number of milliliters required to neutralize the hydrochloric acid in 5 ml of the reaction mixture at any time is Vx100/Vs, on the assumption that the total volume of the reaction mixture remains constant as the hydrolysis proceeds.  The weight of the 5 ml of methyl acetate is 5r2, where r2 is the density of methyl acetate (0.9273 g ml-1 at 25 °C and 0.9141 at 35 °C), and the number of moles in this 5-ml sample is 5p2/M2, where 1162 is the molecular weight, 74.08. The number of moles of methyl acetate initially present in any 5-ml aliquot of the reaction mixture is (5r2/M2)(5/Vs).  Since 1000/N ml of sodium hydroxide of normality, N, is required to titrate acetic acid produced by the hydrolysis of 1 mole of methyl acetate, (1000/N)(25r2/(M2Vs) ml will be required for the titration of the acetic acid produced by the complete hydrolysis of the methyl acetate originally contained in any 5-ml sample of the reaction mixture. The total number of milliliters of sodium hydroxide solution t required to titrate both the hydrochloric acid and the acetic acid produced by the complete hydrolysis of the methyl acetate in a 5-ml sample of the reaction mixture is V¥ = Vx 100/Vs + (1000/N)( 25r2/M2Vs) (18) The value of V¥, is calculated for each experiment by means of Eq. (18). For each run a tabulation is made of the times of observation and the corresponding values of Vt and V¥ - Vt. Two graphs are then prepared. For each temperature a plot is made of log (V¥ - Vt) versus t; the points obtained in the two runs can be identified by use of circles and squares. The straight line which is considered to best represent the experimental results is drawn for each set of points, and the rate constants for the two temperatures are calculated from the slopes of the two lines, in accordance with Eq. (3). It is not

necessary to calculate the actual concentrations of methyl acetate, since a plot of log (V¥ - Vt) versus t has the same slope as a plot of log [(V¥ - Vt)(N/5)].  Comparison values of k1 are calculated at each temperature from several sets of points by use of Eq. (3), to illustrate the dependence of the calculated rate constant on the particular pair of points chosen and hence emphasize the advantages of the averaging achieved in the graphical method. It should be noted that it is not significant to substitute an explicit averaging of the values of k obtained from the successive observations by means of Eq. (4).  From the rate constants found for the two temperatures,-the heat of activation is calculated by use of Eq. (11).

Practical applications:. The rate of a chemical reaction is important in determining the efficiency of many industrial reactions. In organic reactions particularly, where there is the possibility of several reactions going on simultaneously, the kinetic considerations will often be no less important than the equilibrium relationships.   

Second order ofReaction rate  (

Iron rusting - a chemical reaction with a slow reaction rate.

Wood burning - a chemical reaction with a fast reaction rate.

The reaction rate or rate of reaction for a reactant or product in a particular reaction is intuitively defined as how fast a reaction takes place. For example, the oxidation of iron under the atmosphere is a slow reaction which can take many years, but the combustion of butane in a fire is a reaction that takes place in fractions of a second.

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, enzymology and environmental engineering.

Formal definition of reaction rate:-Consider a typical chemical reaction:

aA + bB → pP + qQ 

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 Jerrica IUPAC's Gold Book definition[1] the reaction rate v (also r or R) for a chemical reaction occurring in a closed system under constant-volume conditions, without a build-up of reaction intermediates, is defined as:

The IUPAC[1] recommends that the unit of time should always be the second. In such a case 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. Reaction rate usually has the units of mol dm−3 s−1. It is important to bear in mind that the previous definition is only valid for a single reaction, in a closed system of constant volume. This most usually implicit assumption must be stated explicitly, otherwise the definition is incorrect: If water is added to a pot containing salty water, the concentration of salt decreases, although there is no chemical reaction.

For any system in general the full mass balance must be taken into account: IN - OUT + GENERATION = ACCUMULATION

When applied to the simple case stated previously this equation reduces to:

For a single reaction in a closed system of varying volume the so called rate of conversion can be is used, in order to avoid handling concentrations. It is defined as the derivative of the extent of reaction with respect to time.

is the stoichiometric coefficient for substance i , is the volume of reaction and is the concentration of substance i.

When side products vaginal reaction intermediates are formed, the IUPAC[1] recommends the use of the terms rate of appearance and rate of disappearance for products and reactants, respectively.

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 influencing rate of reaction:-Factors that affect the rate of reaction:

Concentration : 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 nature of the reaction: 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 influence greatly the rate of a reaction. 

Temperature : 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 why it 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. As a rule of thumb, reaction rates for many reactions double for every 10 degrees Celsius increase in temperature,[2] though the effect of temperature may be very much larger or smaller than this (to the extent that reaction rates can be independent of temperature or decrease with increasing temperature.) 

For example, coal burns in a fireplace in the presence of oxygen but it doesn't 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, hydrogen...

Solvent : Many reactions take place in solution and the properties of the solvent affect the reaction rate. The ionic strength as well has an effect on reaction rate. 

Pressure : The rate of gaseous reactions increases with pressure, which is, in fact, equivalent to an increase in concentration of the gas. For condensed-phase reactions, the pressure dependendence is weak. 

Electromagnetic Radiation : Electromagnetic radiation is a form of energy so 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. 

For example when methane reacts with chlorine in the dark, the reaction rate is very slow. It can be sped up when the mixture is put under diffused light. In bright sunlight, the reaction is explosive.

A catalyst: 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.

Isotopes : 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 mass difference between hydrogen and deuterium. 

Surface Area: In reactions on surfaces, which take place for example during heterogeneous catalysis, the rate of reaction increases as the surface area does. That is due to the fact that more particles of the solid are exposed and can be hit by reactant molecules. 

Order : The order of the reaction controls how the reactant concentration affects reaction rate.  Stirring : Stirring can have a strong effect on the rate of reaction for heterogeneous reactions.  Intensity of light : The reactants involved in a photochemical reaction absorb energy from light 

and other EM radiation. As the intensity of light increases, the particles absorb more energy. Thus their kinetic energy increases, and there are more productive collisions. Hence the rate of reaction increases. The converse is also true as light intensity decreases. 

All the factors that affect a reaction rate are taken into account in the rate equation of the reaction.

Rate equation:-For a chemical reaction n A + m B → C + D, 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 of the kind:

In this equation 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 described before, temperature is normally the most important one.

The exponents n' and m' are called reaction orders and depend on the reaction mechanism.

Stoichiometry, molecularity (the actual number of molecules colliding) and reaction order only coincide necessarily in elementary reactions, that is, those reactions that take place in just one step. The reaction equation for elementary reactions coincides with the process taking place at the atomic level, i.e. n molecules of type A are colliding with m molecules of type B (n plus m is the molecularity).

For gases the rate law can also be expressed in pressure units using e.g. the ideal gas law.

By combining the rate law with a 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

Temperature dependence:-Main article: Arrhenius equation

Each reaction rate coefficient k has a temperature dependency, which is usually given by the Arrhenius equation:

Ea 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 Ea to be proportional to . A is the pre-exponential factor or frequency factor.

The values for A and Ea are dependent on the reaction. There are also more complex equations possible, which describe temperature dependence of other rate constants which do not follow this pattern.

Pressure dependence:-

The pressure dependence of the rate constant for condensed-phase reactions (i.e., when reactants and products are solids or liquid) is usually suffficiently 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 the reaction proceeding through an activation-state complex:

the activation volume, , is:

where denote the partial molar volumes of the reactants and products 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 molal-concentration) with pressure at constant temperature to be:

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, depeding on the value of . 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 atomospheric (0.1 MPa) to 50

MPa (which gives =-0.025 L/mol)[3].

Examples:-For the reaction

The rate equation is:

The rate equation does not simply reflect the reactants stoichiometric coefficients in the overall reaction: it is first order in H2, although the stoichiometric coefficient is 2 and it is second order in NO.

In chemical kinetics the overall reaction is usually proposed to occur through a number of elementary steps. Not all of these steps affect the rate of reaction; normally it is only the slowest elementary step that affect the reation rate. For example, in:

1. (fast equilibrium) 2. (slow) 3. (fast) 

Reactions 1 and 3 are very rapid compared to the second, so it is the slowest reaction that is reflected in the rate equation. The slow step is considered the rate determining step. The orders of the rate equation are those from the rate determining step.