In order to understand the reactions and mechanisms of homogeneous as well as heterogeneous reactions, it is important to develop a basic understanding of the terms homogeneous and heterogeneous and how they are different from each other.
A heterogeneous reaction is a class of reaction that happens between two or more reactants that are present in two or more different phases, for instance, the phases of the reactants can be solid-liquid, solid-gas or two immiscible liquids. There are certain reactions that take place on the surface of another substance. To be more specific, when two reactants undergo a chemical change on the interface of a catalyst that can be in solid or liquid form (studies suggest mostly solid) also falls under the category of heterogeneous reactions.
Some examples of heterogeneous reactions are the reaction of solid metals with acids, the corrosion of iron, the electrochemical reaction occurring in batteries and electrolytic cells are all subjected to a heterogeneous reaction. Most of the research and studies regarding heterogeneous reactions are done for heterogeneous catalysts such as the reaction between liquids or gases that happens on the surface of a solid catalyst that helps in initiating the reaction or increasing the reaction rate.
On the other hand, a homogeneous reaction is a class of reaction where the reactants are present in a single phase, that is, either solid, liquid or gas. Homogeneous, unlike heterogeneous reaction, is based on the physical state of the reactants. The important homogeneous reactions are the reactions between gases and reactions happening between liquids or substances dissolved in liquids.
Examples of homogeneous reactions are the combination of natural gas and oxygen to produce flame or a reaction of aqueous solutions of acid and bases. Out of the two homogeneous and heterogeneous reactions, the former is easy to understand as the nature of homogeneous reactions solely depends upon the nature of the interaction of the reactants.
Heterogeneous Chemical Reaction
The chemical interactions that happen between the two reactants that belong to different phases are illustrated with various heterogeneous chemical reactions of reactants of different phases for a clearer understanding of the nature of heterogeneous reactions.
1. Solid-Fluid (Liquid and Gas)
Dissolution of solids, example, MgCO3(S) + 2 HNO3(L) → Mg(NO3)2(aq) + H2O(L) + CO2(g)
Chemical vapour deposition , example, SiH4(g) → Si(S) + 2H2
Sublimation, example, U(s) + 3F2(g) → UF6(g)
Reduction of solid oxides, example, NiO2(s) + H2(g) → Ni(s) + H2O(g)
Metal reduction, example, Zn(s) + O2(g) → ZnO2(S)
2. Liquid - Gas
Dissolution with chemical reaction, examples,
Cl2(S) + 2NaOH(l) → NaOCl(l) + NaCl(l) + H2O(l)
3NO2(g) + H2O(l) → 2HNO3(l) + NO(g)
3. Solid - Solid
Heterogeneous precipitation and calcination reaction, for example,
CoO(S) + Al2O3(S) → Co Al2O4(S)
All the heterogeneous reactions have one thing in common and that is before any heterogeneous chemical reaction occurs the reactant in a bulk of a particular phase is transferred to either the interface of the two bulk phases or is completely transported to the bulk of another phase. Usually, heterogeneous reactions take place in a number of steps among which one or more chemical steps involved are accompanied by the intermediate steps that are purely physical in nature and are just responsible for transporting the bulk phase of certain reactants to the bulk phase of other reactants. The concentration of the reactant in one bulk phase is different from the concentration of reactants at the interface of the two different bulk phases and this difference in the concentration is actually the driving force for this phase transfer process. Thus in order to study the rate of chemical reaction, these are the factors that are of utmost importance.
The change in the rate equation is mostly from the mass transfer from one phase to the other.
The contact pattern of the reacting phase.
Thus in steady-state condition when the heterogeneous reaction takes place in multiple steps, then the overall rate of reaction is equal to the rate of reaction of the individual reactions taking place in a series. Hene the rates of reaction can be expressed as roverall = r1= r2 =.........= rn.
Thus when the mass transfer takes place the rate of reaction is calculated on the basis of molar reflux which is basically no. of moles per unit time per surface. Therefore, reaction rate based on unit volume will be
-rA = (-1 / V) (dNA / dt) = mole of a reactant / (volume of fluid * time )
Based on the mass of a solid,
-rA। = (-1 / W) (dNA / dt) = mole of a reactant / (mass of * time )
Based on unit interfacial surface/unit surface of solid
-rAII = (-1 / S) (dNA / dt) = mole of a reactant / (interfacial surface * time )
All these rates of reactions are based on
mole of a reactant / time = (-rA * V) (-rA। * W) (-rAII * S)
-rA = (W /V) rA।
-rA। = (S / W) rAII
-rAII = (V / S) rA
Thus for a non-elementary rate of reaction as follows
C6H5CH3(l) + H2(g) → C6H6(l) + CH4(g)
The rate of reaction for this heterogeneous reaction will be
-r⋎। = (kPH2 PT) / (1 + KBPB + KTPT)
where KB and KT are the absorption constants with units kPa-1 (atm-1) and the units of reaction rate is
[k] = mol toluene (C6H5CH3) / kgcat • s • kPa2
Alternatively, the rate of reaction in terms of concentration can be written as
Pi = Ci RT.
A homogeneous reaction can be defined as the reaction where all the reactants and the products formed from the chemical interaction of the reactants are all in one single phase that can be gas, liquid or solid phase and do not possess any phase boundaries. A more clear picture of the nature of the reaction can be drawn by a few examples of homogeneous reactions.
2SO2(g) + O2(g) ⇋ 2SO3(g) whose rate of forward reaction is v1 = k1 • [SO2]2 • [O2] and the rate of backward reaction will be v2 = k2 • [SO3]2 . Now when a homogeneous reaction proceeds, just like a heterogeneous reaction, it takes place in a series of many intermediate chemical reactions and it constitutes the mechanism of the reaction and is known as a non-elementary reaction. This particular series mechanism of non-elementary homogeneous reactions rate law. Thus elaborating the derivation of the rate law with an elementary example of homogeneous reaction as follows:-
2H2 + O2 → 2H2O
This reaction actually follows many elementary steps:
H2 + O2 → HO2 + H
H2 + HO2 → OH + H2O
OH + H2 → HO2 + H
O2 + H → OH + O
H2 + O → OH + H
Now, for most of the reactions, the elementary steps that are the constituents of the mechanism of such reactions are not known and only single reactions are observed. It is because the amount of reactants and the products that are formed in the intermediate reactions formed are very few and often escape detection. Also, the velocity with which they form and dissociate is very fast. Thus it makes it very difficult to detect. Thus based on the experimental observations of the collision theory, the law of mass action states that, for an elementary reaction, its reaction rate at a constant temperature is equal to the product of the concentration of the reactants. Thus for a single elementary reaction, the reaction rate is given by:-
𝑓( [A] [B] [C] T t) = k [A] [B], where k is the rate constant. Thus the series of equations that describes the evaluation of each species is given by-
d[A] / dt = -k a [A] [B]
d[B] / dt = -k b [A] [B]
d[C] / dt = -k c [A] [B]
And this equation is said to be in order 2 and is also very difficult to determine due to the complexity of the mechanism of such a reaction. Thus, it has also been observed that the rate of reaction is approximated to be proportional to the power of the products of concentration of the reactants, the equation is as follows,
𝑓( [A] [B] [C] T t) = k [A]α [B]𝛽
Where α is called the order of reaction of A, 𝛽 is called the order of reaction of B and the sum of the two exponents, n = α + 𝛽 is called the rate of reaction. Also, the rate constant k of the nth order of reaction has dimension (time)-1 (concentration)n-1. Thus the rate of reaction between hydrogen and bromine will be
H2 + Br2 → 2HBr
𝑓( [H2] [Br2] [HBr]) = k1 [H2] [Br2]½ / k2 + [HBr] / [Br2]
Heterogeneous Catalytic Reaction
Heterogeneous catalytic reactions are the reactions that occur in the presence of a catalyst where the phase of the catalyst differs from that of the reactants. Catalysts are generally useful for a reaction to increasing the rate of reaction without themselves getting consumed and thus can be used later. In heterogeneous catalytic reactions, the catalysts are generally present in the solid phase and the reactants are generally in the gaseous phase. These reactions occur on the interface of the catalyst where the molecules of the reactants are adsorbed on the surface and then the reaction occurs resulting in the formation of desired products which are then desorbed from the surface of the catalyst. Thus the thermodynamics, heat transfer or mass transfer generally affects the rate of heterogeneous catalytic reaction. One of the most common examples of heterogeneous catalyst reaction is hydrogenation of ethane on the catalytic solid surface that involves adsorption, reaction and desorption. This is illustrated by the diagram below.
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This kind of reaction is commonly known as heterogeneous catalytic hydrogenation and is one of the most widely used methods in industries these days to produce chemicals. Heterogeneous catalytic hydrogenation reactions have found their place mostly in industries like pharmaceuticals, flavour and fragrance industries, fine chemicals, agrochemicals and dietary supplements. Most of the chemicals that are produced in these industries are using 10-20% of heterogeneous catalytic hydrogenation reactions for their production. It is because the catalytic reactions are generally highly selective in nature and are therefore easy to work upon. The reactions are atomic efficient and the catalyst does not get consumed and can be recovered and recycled.
In industrial heterogeneous catalytic hydrogenation reaction, the most precious metal catalysts are deposited from the solution in form of a powdery substance carefully on a support that is heavy, porous, bulky, cheap and usually granular like active carbon or calcium carbonate or alumina. For example, platinum on carbon is produced by the reduction of chloroplatinic acid that is catalysed by 5% ruthenium on activated carbon or 1% platinum on alumina. One of the most essential steps in the catalytic reaction of any kind is adsorption. Now adsorption is classified broadly into two parts, namely, physisorption and chemisorption. Physisorption is weak adsorption bonding and the molecules of the reactants are bound to the surface of the catalyst by weak Van Der Waal forces like dipole-dipole moment, induced dipole interactions and London dispersion forces. whereas chemisorption is a strong adsorption bonding where the molecules of the reactant approach closely to the atoms on the surface of the catalyst through electron cloud overlapping. Thus in this adsorption process, the adsorbent and the adsorbent share a chemical bonding by sharing their electrons. Usually, the heterogeneous catalytic reaction falls between these two processes. Thus the mechanism of the catalytic reaction in a heterogeneous system will be in the following steps:
Diffusion of reactants to the surface and the rate of diffusion of the reactants on the surface is directly proportional to the concentration of the reactants and the thickness of the boundary layer.
Adsorption of reactants on the surface of the catalyst happens when the adsorbate forms bonding with the adsorbent. Thus the ability of the atoms or the molecules to form bonds with the atoms on the surface of the catalyst can be efficiently calculated as a sticking coefficient. This represents the percentage of the ratio of the no. of molecules or atoms that have stuck to the surface of the catalyst.
The reaction is indicated by the bonds that are formed by the atoms of reactants and the catalyst.
Desorption of the products. This happens when the bonds between the product that is formed and the catalyst breaks.
The product formed is diffused from the surface of the catalyst. This happens when the bond is cleaved and the product diffuses from the surface of the catalyst without changing the characteristics of the catalyst.
An example of this is the contact process which is used in the industries to produce sulphuric acid in high concentrations. Earlier platinum was used as a catalyst but since it has a tendency to react with the arsenic impurities that are present in the sulfur feedstock, thus vanadium oxide (V2O5) is now used as the active catalyst. In this process sulfur oxide and oxygen are gases and vanadium oxide is solid.
2SO2(g) + O2(g) ⇋ 2SO3(g) in presence of V2O5(S)
2V2O5(S) + 2SO2(g) ⇋ 2SO3(g) +2V2O4(S)
2V2O4(S) + O2(g) ⇋ 2V2O5(S)
2SO2(g) + O2(g) ⇋ 2SO3(g)