Chemical Kinetics, Chemistry tutorial


Define: It is a stream of chemistry which mainly deals by the rate of chemical reactions, with factors affecting such rates, and by applications of rate studies to explain the method of reactions.

Chemical kinetics is mainly concerned by the rates of chemical reactions, that is, by the quantitative explanation of how fast chemical reactions take place and the factors influencing these rates. The chemist employs kinetics as a tool to comprehend the basic features of reaction pathways, a subject which continues to evolve by ongoing research. The applied chemist employs this understanding to work out new and/or better ways of accomplishing desired chemical reactions. This might comprise enhancing the outcome of desired products or making a better catalyst. The chemical engineer employs kinetics for reactor design in chemical reaction or process engineering.

Chemical kinetics is basically the study and conversation of chemical reactions with respect to reaction rates, effect of different variables, re-arrangement of atoms and preparation of intermediates and so on. There are numerous topics to be discussed and each of such topics as a tool for the study of chemical reactions. By the way, the study of motion is termed as kinetics, from the Greek kinesis, meaning movement.

At the macroscopic level, we are concerned in amounts reacted, formed, and the rates of their formation. At molecular or microscopic level, the given considerations should as well be made in the conversation of chemical reaction methods.

  • Molecules or atoms of reactants should collide by one other in chemical reactions.
  • The molecules should encompass adequate energy to initiate the reaction.
  • In certain cases, the orientation of the molecules throughout the collision should as well be considered.

Factors that affect Reaction rates:

1) Nature of the reactants:

Some of the reactant molecules react in a hurry, others react very slowly. Why? The state of matter comes into play; solids will react much more slowly as compare to liquids or aqueous solutions.

a) Physical state matters a big deal. Let assume gasoline (l) versus gasoline (g).

K2SO4(s) + Ba(NO3)2(s) → no reaction; whereas the reaction is fast in the aqueous solution.

b) Chemical identity: What precisely is reacting? Generally ions of opposite charge react very rapidly. As well, the more bonds between reacting atoms in a molecule, the slower the reaction rate. Why? More energy is needed to apart the molecule into its 'bits'. Substances having strong bonds (that is, larger bond energies) will react much more gradually. Illustrations: metallic sodium reacts much faster with water as compare to metallic calcium. The oxidation of methane can be raised by an increase in temperature; photosynthesis is extremely slow and changes very little by an increase in temperature simply as there are numerous intermediate steps frequently needing specific catalysts.

2) Concentration of reactants:

The more molecules present, the more collisions take place, the faster reaction carries on, the greater its rate. Simple adequate. It is much similar to the crowded halls throughout the passing period!

3) Temperature:

"Heat 'em up and speed 'em up!"; the faster molecules move, the more possibly they are to collide and the more energetic the collisions become. In another words, the more possibly 'molecular damage' will take place. The bonds will be broken and latest bonds will form.

a) The increase in temperature generates more successful collisions which are capable to overcome the required activation energy, thus, a general increase in reaction rate having increasing temperature.

b) However, a general rule of thumb is that 10oC increases in temperature will twice the reaction rate.

c) This in reality based on the magnitude of the activation energy, Ea* and the temperature range. 

4) Surface area of reactants: exposed surfaces affect speed:

a) Apart from substances in the gaseous state or solution, the reactions take place at the boundary, or interface, among the two phases. In other words, at the exposed surface of the substance.

b) Thus, the greater surface area exposed, the greater chance of collisions between the particles, therefore, the reaction must carry on at a much faster rate. Example: coal dust is extremely explosive as opposed to a piece of charcoal. The aqueous solutions are eventual exposure!

5) Adding an inert gas has No Effect on the rate [or equilibrium] of the reaction as it is NOT in the reaction procedure! This is a standard rule with regard to the test questions.

6) Catalysis is the raise in rate of a chemical reaction because of the participation of a substance known as a catalyst. (Sounds similar to a circular argument, huh?) Unlike other reagents in the chemical reaction, a catalyst is not consumed; thus it can be utilized again and again. A catalyst might participate in multiple chemical conversions. The effect of a catalyst might differ because of the presence of other substances termed as inhibitors (that decrease the catalytic activity) or promoters (that raise the activity).

a) The biological catalysts are proteins known as enzymes. They are extremely 'shape' particular to their substrate and encompass an active site which fascinates the substrate by employing intermolecular forces. We are familiar with the 'lock and key' enzyme analogy. As their shape is generally determined through IMFs (as opposed to real chemical bonds), which are extremely temperature sensitive. Modifying the pH or heating the enzyme simply interrupts the IMFs and causes the enzyme to denature (that is, lose its 3-D shape).

b) The catalyst is a substance which changes the rate of reaction via modifying the reaction pathway. Most of the catalysts work via lowering the activation energy required for the reaction to proceed, thus, more collisions is successful and the reaction rate is raised.

c) Keep in mind! The catalyst is not a part of the chemical reaction and is not utilized throughout the reaction. Generally, the catalyst contributes in the rate-determining or slowest step.

d) Catalysts might be homogeneous or heterogeneous catalysts. Homogeneous catalyst in the similar phase as the reactants. Therefore, go ahead; guess what heterogeneous signifies in this context. Yep. A heterogeneous catalyst is in a different state as compare to the reactants. Heterogeneous catalysis offers the benefit that products are readily separated from the catalyst, and heterogeneous catalysts are frequently more stable and degrade much slower as compare to the homogeneous catalysts.

Illustration: H2O2 decomposes comparatively slowly into H2O and O2; though; exposure to light accelerates this procedure AND by the assistance of MnO2, it goes very FAST!!

Note: As a catalyst lowers the activation energy barrier, the reverse and forward reactions are both accelerated to the similar degree.

e) Homogeneous catalysts in reality appear in the rate law as their concentration influences the reaction.

Reaction Rates:

The Chemical reaction rates are the rates of change in concentrations or amounts of either reactants or products. For changes in amounts, the units can be one of mol/s, g/s, lb/s, kg/day and so on. For the changes in concentrations, the units can be one of mol/(L s), g/(L s), %/s and so on.

With respect to the reaction rates, we might deal by average rates, instantaneous rates and initial rates based on the experimental conditions.

Thermodynamics and kinetics are the two factors which influence reaction rates. The study of energy gained or liberated in chemical reactions is known as thermodynamics, and such energy data are known as thermodynamic data. Though, thermodynamic data encompass no direct correlation with the reaction rates, for which the kinetic factor is possibly more significant. For illustration, at room temperature (that is, a broad range of temperatures), thermodynamic data points out that diamond shall transform to graphite, however in reality, the conversion rate is extremely slow that most of the people assume that diamond is forever.

Rate Laws:

The rate of a chemical reaction is, possibly, it's most significant property as it dictates whether a reaction can take place throughout a lifetime. Recognizing the rate law, an expression associating to the rate to the concentrations of reactants, can assists a chemist in adjusting the reaction conditions to obtain a more appropriate rate. Whenever there are two competing reactions for a single reagent, one can, recognizing the rate law, favor the exclusive preparation of a single product.

To get this type of knowledge regarding reactions, we will initially define what rate signifies. We will then derive the rate law expression. By employing the process of initial rates, we will discuss how to find out the form and order of the rate law. Subsequently, we will probe rate laws in depth and introduce the integrated rate law as the alternative form of the simple rate law which allows us the other, simpler, experimental process to find out the order of the rate law. The integrated rate law will as well let us to find out the half-lives of chemical reactions.

Units of rate constant:

A point that frequently seems to cause endless confusion is the fact that the units of rate constant based on the form of the rate law in which it appears that is, a rate constant appearing in the first order rate law will encompass different units from a rate constant appearing in the second order or third order rate law. This follows instantly from the fact that the reaction rate for all time consists of the similar units of concentration per unit time that should match the overall units of a rate law in which concentrations increased to varying powers might appear. The good news is that it is extremely straightforward to find out the units of a rate constant in any particular rate law. Below are some illustrations.

i) Let suppose the rate law ν = k[H2][I2]. If we replace with units into the equation, we get

(mol dm-3 s-1) = [k] (mol dm-3) (mol dm-3)

Here, the notation [k] signifies 'the units of k'. We can reorganize this expression to determine the units of the rate constant, 'k'.

[k] = (mol dm-3 s-1)/(mol dm-3) (mol dm-3) = mol-1 dm3 s-1

ii) We can apply the similar treatment to a first order rate law, for illustration

ν = k [CH3N2CH3].

(mol dm-3 s-1) = [k] (mol dm-3)

[k] = (mol dm-3 s-1)/(mol dm-3) = s-1

iii) As a final example, consider the rate law ν = k [CH3CHO]3/2.

(mol dm-3 s-1) = [k] (mol dm-3)3/2

[k] = (mol dm-3 s-1)/(mol dm-3)3/2 = mol-1/2 dm3/2 s-1

Determining the rate law from experimental data:

A kinetics experiment comprises of measuring the concentrations of one or more reactants or products at a number of various times throughout the reaction. We will evaluation some of the experimental methods used to make such measurements.

i) Isolation method:

The isolation process is a method for simplifying the rate law in order to find out its dependence on the concentration of a single reactant. Once the rate law has been simplified, the differential or integral procedures illustrated in the given subsections might be employed to find out the reaction orders.

The dependence of reaction rate on the selected reactant concentration is isolated by having all other reactants present in a big excess, in such a way that their concentration remains essentially constant all through the course of the reaction. As an illustration, consider a reaction A + B → P, in which B is present at a concentration 1000 times more than A. Whenever all of species A has been employed, the concentration of B will merely have changed via 1/1000, or 0.1%, and therefore 99.9% of the original B will still be present. It is thus a good approximation to treat its concentration as constant all through the reaction.

This very much simplifies the rate law as the (constant) concentrations of all the reactants present in large excess might be combined by the rate constant to result a single effective rate constant. For illustration, the rate law for the reaction considered above will become:

ν = k [A]a[B]b ≈ k [A]a[B]0b = keff[A]a with keff = k[B]0b

Whenever the rate law consists of contributions from a number of reactants, a sequence of experiments might be carried out in which each and every reactant is isolated in turn.

ii) Differential methods:

Whenever we encompass a rate law which based only on the concentration of one species, either as there is merely a single species reacting, or as we have utilized the isolation process to manipulate the rate law, then the rate law might be written:

ν = k[A]a

log ν = log k + a log[A]

The plot of logν against log[A] will then be a straight line having a slope equivalent to the reaction order, 'a', and an intercept equivalent to log k. There are mainly two ways in which to get data to plot in this manner.

a) We can evaluate the concentration of the reactant [A] as the function of time and use this data to compute the rate, ν = -d[A]/dt, as a function of [A]. A plot of logν vs log[A] then results the reaction order with respect to A.

b) We can prepare a sequence of measurements of the initial rate ν0 of the reaction by various initial concentrations [A]0. These might then be plotted as above to find out the order, a.

This is a generally used method termed as the initial rates procedure.

iii) Integral methods:

If we have evaluated the concentrations as a function of time, we might compare their time dependence by the suitable integrated rate laws. Again, this is the most straightforward if we have simplified the rate law in such a way that it depends on just one reactant concentration. The differential rate law given in equation ν = k[A]a will give mount to various integrated rate laws based on the value of a. The most generally encountered ones are:

Zeroth order integrated rate law: [A] = [A]0 - kt

                    A plot of [A] vs. t will be linear, having a slope of -k.

First order integrated rate law: ln[A] = ln[A]0 - kt

                    A plot of ln[A] vs. t will be linear having a slope of -k.

Second order integrated rate law: 1/[A] = 1/[A]0 + 2kt

                    A plot of 1/[A] vs. t will be linear having a slope of 2k.

If none of such plots yield in a straight line, then more complex integrated rate laws should be tried. 

iv) Half lives:

The other way of finding out the reaction order is to investigate the behavior of the half life as the reaction carries on. Particularly, we can assess a sequence of successive half lives. t = 0 is employed as the beginning time from which to assess the first half life, t1/2(1). Then t1/2(1) is employed as the start time from which to assess the second half life, t1/2(2) and so forth.

=> Zeroth order t1/2 = [A]0/2k

As at t1/2(1), the new starting concentration is 1/2[A]0, successive half lives will reduce via a factor of two for a Zeroth order reaction.

=> First order t1/2 = ln2/k 

There is no dependence of the half life on concentration; therefore t1/2 is constant for a first order reaction.

=> Second order t1/2 = 1/k[A]0

The inverse dependence on concentration signifies that successive half lives will two times for the second order reaction.

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