Applications of Oxidation and

Reduction Reactions in

Everyday life

Submitted To

Department of Chemistry

Manipal Institute of Technology

Manipal.

Faculty Advisor

Dr. Santhosh.L.Gaonkar

Submitted By

Hemant Hegde

M.Sc.III sem.

Reg.No.103101015

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Contents Page.no.

1. Introduction………………………………………………….3-8

1.1 Types of reactions………………………………………………3

1.2 Oxidation - Reduction reaction…………………………………..5

1.3 Oxidizing and reducing agents…………………………………...7

1.4 Oxidation number………………………………………………....8

2. Application in everyday life………………………………..9-62

2.1 Combustion……………………………………………………...92.2 Bleaching………………………………………………….........12 2.3 Nitrogen fixation……………………………………………….16 2.4 Metabolism………………………………………………..........242.5 Electrochemical cells…………………………………………...312.6 Corrosion……………………………………………………….362.7 Enzymatic browning………………………………………........382.8 Water purification………………………………………………452.9 Aging…………………………………………………………...502.10 Biogas production…………………………………………….522.11 Weathering of rock…………………………………………....56 2.12 Photo-oxidation……………………………………………….592.13 Photography…………………………………………………..60

3. Conclusion………………………………………………….63

4. References………………………………………………….64

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1. Introduction:

A chemical reaction is a process that is usually characterized by a chemical change in which the starting

materials (reactants) are different from the products. Chemical reactions tend to involve the transfer of

electrons, leading to the formation and breaking of chemical bonds. The change produced by a chemical

reaction is quite different from a purely physical change, which does not affect the fundamental properties of

the substance itself.

For example: A piece of copper can be heated, melted, moulded into different shapes, and so forth, yet

throughout all those changes, it remains pure copper, an element of the transition metals family. But suppose a

copper roof is exposed to the elements for many years. Copper is famous for its highly noncorrosive quality,

and this, combined with its beauty, has made it a favored material for use in the roofs of imposing building

However, copper does begin to corrode when exposed to air for long periods of time.

1.1 Types of Chemical Reactions

There are several different types of chemical reactions and more than one way of classifying them. Here is a

list of the different types of reactions, with examples of each type included.

combination reaction : This is a reaction in which two or more elements or compounds combine to form a

single product. This type of reaction follows the general equation[1]

A + B C

where A and B may be either elements or compounds.

Here are some examples:

2Na(s) + Cl2(g) 2NaCl(s)

MgO(s) + H2O(l) Mg(OH)2(aq)

SO2(g) + H2O(l) H2SO3(aq)

Decomposition reaction: In this type of reaction, a single reactant, a compound, breaks into two or more parts.

Often these are the most difficult to predict. Here is the general equation:

AB A + B

where A and B may be either elements or compounds.

Here are some examples of decomposition reactions:

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2H2O(l) 2H2(g) + O2(g)

H2CO3(aq) H2O(l) + CO2(g)

CaCO3(s) CaO(s) + CO2(g)

2KClO3(s) 2KCl(s) + 3O2(g)

Single replacement or displacement reaction: In this type of reaction, a more active element replaces a less

active element in a compound. Among the halogens, F2 is the most active halogen, and the activity of the

halogens decreases as we go down the group. For the metals, it will need to be given an activity series.

General equation:

A + BC AC + B

where A is a metal.

Here is an example of a displacement reaction in which a metal is involved:

Cu(s) + 2AgNO3(aq) 2Ag(s) + Cu(NO3)2(aq)

General equation: A + BC BA + C

where A is a nonmetal.

Here is an example of a displacement reaction where a nonmetal is involved:

Cl2(g) + 2NaI(aq) 2NaCl(aq) + I2(s)

Double replacement or displacement reaction : In this type of reaction, two compounds react to form two new

compounds. The formation of a molecular compound such as water, the formation of a gas, or the formation of

a precipitate usually drives these reactions. Here’s the general equation:

AB + CD AD + CB

And here are a couple of examples:

Pb(NO3)2(aq) + 2KI(aq) 2KNO3(aq) + PbI2(s)

HCl(aq) + NaOH(aq) H2O(l) + NaCl(aq)

Oxidation-Reduction or Redox Reaction

In a redox reaction the oxidation numbers of atoms are changed. Redox reactions may involve the transfer of

electrons between chemical species.The reaction in which I2 is reduced to I- and S2O32- (thiosulfate anion) is

oxidized to S4O62- provides an example of a redox reaction:

2 S2O32−(aq) + I2(aq) → S4O6

2−(aq) + 2 I−(aq)

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Hydrolysis reaction : A reaction that involves water. Here is the general equation for a hydrolysis reaction:

X(aq) + H2O( l ) HX(aq) + OH- (aq)

1.2. Oxidation-Reduction Reactions

Reduction and oxidation (redox) reactions are an important class of chemical reactions since they are the

driving force behind a vast range of process, both desirable (for example breathing in mammals) and

undesirable (for example rusting of iron). A redox reaction is characterized by the fact that electrons are

produced as in oxidation reaction or are used by the reaction as in reduction reaction. An oxidation reaction

must always be paired with a reduction reaction, as the oxidation reaction produces the electrons required by the

reduction reaction. Redox reactions are the energy-producing reactions in living systems as it is play a major

role in electron transport process during cell metabolism. The core of a redox reaction is the passing of one or

more electrons from one species to another. The species that loses electrons is said to be oxidized, and the

species gaining electrons is reduced. Redox reactions can be defined in several ways as follows:

Oxidation and reduction in terms of electron transfer

Oxidation is a chemical change in which electrons are lost by an atom or group of atoms and reduction

is a chemical change in which electrons are gained by an atom or group of atoms. As a simple example of an

oxidation-reduction consider what happens when we a strip of metallic zinc is dipped into a blue solution of

copper (II) sulfate. The strip of zinc becomes coated with a reddish-brown layer of metallic copper.  The

molecular equation for this reaction is         

          Zn(s) + CuSO4(aq) ZnSO4(aq) + Cu(s)

In this simple example we can say that zinc has been oxidized and copper has been reduced.  For convenience,

we can think of this reaction as two separate parts or half reactions one involving the loss of two electrons by an

atom of zinc and the other being the gain of two electrons by the copper(II) ion.  The two half reactions are

Zn(s) Zn2+(aq) + 2 e-        oxidized

Cu2+(aq) + 2 e- Cu(s)        reduced

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Oxidation and reduction in terms of oxygen transfer

In terms of oxygen transfer ,oxidation is the oxygen gaining process whereas, reduction is the electron

losing process.

For example, in the extraction of iron from its ore:

Oxidation and reduction in terms of hydrogen transfer

These are old definitions which are not used very much now a day. This definition most likely

come across in organic chemistry. Where oxidation is a hydrogen losing process and reduction is hydrogen

gaining process.It is clear that these are exactly the opposite of the oxygen definitions.

For example, ethanol can be oxidised to ethanal:

We need to use an oxidising agent to remove the hydrogen from the ethanol. A commonly used oxidising agent

is potassium dichromate(VI) solution acidified with dilute sulphuric acid.

1.3. Oxidizing and reducing agents6

In redox processes the reductant transfers electrons to the oxidant. Thus, in the reaction, the reductant

or reducing agent loses electrons and is oxidized, and the oxidant or oxidizing agent gains electrons and is

reduced. The pair of an oxidizing and reducing agent that are involved in a particular reaction is called a redox

pair[1].

Oxidizing agents or Oxidizers

Substances that have the ability to oxidize other substances are said to be oxidative or oxidizing and

are known as oxidizing agents, oxidants, or oxidizers. In another way, the oxidant (oxidizing agent) removes

electrons from another substance i.e. it oxidizes other substances, and also itself reduced. Because it "accepts"

electrons, so it is also called an electron acceptor. Oxidants are usually chemical elements or substances with

elements in high oxidation numbers. For example , H2O2, MnO−4, CrO3, Cr2O2−7, OsO4 etc. or highly

electronegative substances or elements that can gain one or two extra electrons by oxidizing an element or

substance. For example O, F, Cl, Br etc.

Reducing agents or Reducers

Substances that have the ability to reduce other substances are said to be reductive or reducing and

are known as reducing agents, reductants, or reducers. The reductant (reducing agent) transfers electrons to

another substance i.e. it reduces others, and is thus itself oxidized. And, as it "donates" electrons, it is also called

an electron donor. Electron donors can also form charge transfer complexes with electron acceptors.

Reductants in chemistry are very diverse. Electropositive elemental metals, such as lithium, sodium,

magnesium, iron, zinc, and aluminium, are good reducing agents. These metals donate or give away electrons

readily. Hydride transfer reagents, such as NaBH4 and LiAlH4, are widely used in organic chemistry,. primarily

in the reduction of carbonyl compounds to alcohols. Another method of reduction involves the use of hydrogen

gas (H2) with a palladium, platinum, or nickel catalyst. These catalytic reductions are used primarily in the

reduction of carbon-carbon double or triple bond.

Example: Formation of NaCl. The reaction can be written as follows

Na  +  Cl2 2NaCl

The Na starts out with an oxidation number of zero (0) and ends up having an oxidation number of +1. It has

been oxidized from a sodium atom to a positive sodium ion.

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The Cl2 also starts out with an oxidation number of zero (0), but it ends up with an oxidation number of -1.

It therefore, has been reduced from chlorine atoms to negative chloride ions. The substance bringing about the

oxidation of the sodium atoms is the chlorine, thus the chlorine is called an oxidizing agent. In other words, the

oxidizing agent is being reduced (undergoing reduction). The substance bringing about the reduction of the

chlorine is the sodium, thus the sodium is a reducing agent.

1.4. Oxidation Numbers  

The oxidation number of an atom in a substance is defined as the charge of the atom if it existed as

a monoatomic ion, or a hypothetical charge assigned to the atom in the substance by a set of rules. Sometimes

they are also called as oxidation state. A comparison of the oxidation states of atoms in reactants enable us to

keep a track of the transfer of electrons in a chemical reaction.So this will make us to know that there was a

transfer of electrons in other words there was a redox reaction.In case of organic molecules oxidation state is

equal to the charge an atom would have if all the electrons were assigned to the more electronegative atom.[6]

Iron is a good reductant and it becomes Fe2+ or Fe3+ depending on the reaction conditions.

Fe Fe2+ + 2e-

Fe Fe3+ + 3e-

Thus, it is necessary to specify the number of electrons to be donated and to be accepted. For this

purpose, a parameter, oxidation number, was defined. The oxidation number for monatomic elements is the

number of charges possessed by that atom. The oxidation numbers of Fe are 0, +2 and +3, respectively. In

order to extend the concept of oxidation number to polyatomic molecules, it is necessary to know the accurate

distribution of electrons in the molecule. Since this is a difficult procedure, it was decided that a formal charge

is to be assigned to each atom under some rule, and the oxidation number is defined based on the formal charge.

2. Applications in everyday life

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It has been found that there are wide range applications of redox- reaction in our daily life. In other

words, the world is full of examples of this highly significant form of chemical reaction. One such example is

combustion, or an even more rapid form of combustion, which is otherwise known as explosion. Likewise the

metabolism of food, as well as other biological processes, involves oxidation and reduction reactions. Also

found to ,do a number of processes that take place on the surfaces of metals: when iron rusts; when copper turns

green; or when aluminum forms a coating of aluminum oxide that prevents it from rusting. Oxidation-reduction

reactions also play a major role in electrochemistry, which has a highly useful application to daily life in the

form of batteries.

Here are the some familiar examples which are found to occur through the redox mechanism in our daily life.

2.1. Combustion

Combustion means burning. Any time a material burns, an oxidation-reduction reaction occurs. The two

equations below show what happens when coal (which is nearly pure carbon) and gasoline (C8 H18 ) burn. One

can see that the fuel is oxidized in each case:

C + O2 CO2

2 C8 H18 + 25 O2 16 CO2 + 18 H2O

In reactions such as these, oxidation occurs very rapidly and energy is released. That energy is utilised

in homes and buildings; to drive automobiles, trucks, ships, airplanes, and trains; to operate industrial processes;

and for numerous other purposes

A simple combustion reaction is given for methane. The combustion of methane means that it is possible to

burn it. Chemically, this combustion process consists of a reaction between methane and oxygen in the air.

When this reaction takes place, the result is carbon dioxide (CO2), water (H2O), and a great deal of energy. The

following reaction represents the combustion of methane[20]:

CH4(g) + 2O2(g) CO2(g) + 2H2O(g) + energy

  One molecule of methane combined with two oxygen molecules, react to form a carbondioxide molecule

and two water molecules usually given off as steam or water vapour during the reaction and energy.

Natural gas is the cleanest burning fossil fuel. Coal and oil, the other fossil fuels, are more chemically

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complicated than natural gas, and when combusted, release a variety of potentially harmful air pollutants.

Burning methane releases only carbon dioxide and water. Since natural gas is mostly methane, the combustion

of natural gas releases fewer byproducts than other fossil fuels.

Fuel:

Solid:- Coal, wood - consists mainly of C, H, & O + impurities

Liquid:- Large hydrocarbon molecules of varying boiling point mainly C & H : Petrol, Diesel, Fuel oil etc;

Gas:- Small hydrocarbon molecules - methane, ethane, propane, butane, etc plus a range of manufactured gases

e.g.: H2 , Acetylene etc..

Oxidant:

Usually air (the oxygen in the air), where air is unavailable the oxidant has to be carried as well as the fuel -

space vehicles, rockets etc

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Mixing

Solids :- pulverizing to powder or small lumps

Liquids :- spray nozzles, atomizers, vaporisers,carburettors, burners.

Gases :- mixing valves, chambers, burners; usually need precautions to avoid explosion and flash back .

Ignition

Simply mixing methane and air will not cause it to burn.The molecules need to reach a certain threshold energy

level before the combustion process will proceed. This may be provided initially by another flame, a spark or

hot surface. The combustion process itself ( if sustained )then continues the ignition process.

Combustion

The chemical dissociation of the fuel and it's recombination with oxygen. Energy and Mass are conserved.

Combustion reactions always involve molecular oxygen (O2). Anytime anything burns , it is a combustion

reaction. Combustion reactions are almost always exothermic (i.e., they give off heat). For example when wood

burns, it must do so in the presence of O2 and a lot of heat is produced.

Wood as well as many common items that combust are organic (i.e., they are made up of carbon, hydrogen and

oxygen). When organic molecules combust the reaction products are carbon dioxide and water (as well as heat)

. organic molecules + O2 CO2 + H2O + Heat

For example consider the combustion of methanol (rubbing alcohol):

CH3OH + O2 CO2 + H2O + Heat

Of course, not all combustion reactions release CO2 and water, e.g., the combustion of magnesium metal:

2Mg + O2 2MgO + Heat

As propane burns in air, its carbon atoms are oxidized when they combine with oxygen to form carbon dioxide.

In turn, molecular oxygen is reduced by the hydrogen atoms, forming water. The heat produced can be used

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directly such as in the cooking of foods or to cause the expansion of the gaseous products produced to perform

mechanical work such as in an internal combustion or steam engine.

C3H8 + 5 O2 3 H2O + CO2 + Heat

Many other substances besides hydrocarbons can be used as fuels. For example, the alcohols, such as methanol

(CH3OH) and ethanol (CH3CH2OH) are often used in racing cars. Ethanol mixed with gasoline, called gasohol ,

is currently being explored as a substitute for gasoline. Among the simplest fuels is molecular hydrogen (H 2)

which readily reacts with oxygen forming water as shown:

2 H2 + O2 2 H2O + Energy

The simplicity and "nonpolluting" aspect of this oxidation-reduction reaction, the amount of energy produced,

and the relative abundance of both hydrogen and oxygen in our environment, makes hydrogen a very attractive

alternative fuel source. Research efforts are currently focused on further developing the technology to broaden

its use as a source of energy.

2.2.Bleaching

A bleaching is a process that can whiten or decolorize the materials. Substance that decolorizes or whiten the

materials are called “bleaching agents”. Coloured substances generally contain groups of atoms, called

chromophores , that can absorb visible light having specific, characteristic wavelengths, and reflect or transmit

the part of light that is not absorbed. For example, if a chromophore absorbs blue light, it will reflect light of the

complementary color, and the chromophore-containing substance will appear yellow. Bleaching agents

essentially destroy chromophores thereby removing the color, via the oxidation or reduction of these absorbing

groups.[4]

Action of bleaching agents

Bleaching agents are compounds which are used to remove color from substances such as textiles. In

earlier times textiles were bleached by exposure to the sun and air. Today most commercial bleaches are

oxidizing agents, such as sodium hypochlorite (NaOCl) or hydrogen peroxide (H2O2) which are quite effective

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in "decolorizing" substances via oxidation. The action of these bleaches can be illustrated in the following

simplified way:

We know that, an oxidizing agent is any substance which causes another substance to lose one or more

electrons. The decolorizing action of bleaches is due in part to their ability to remove those electrons which

are activated by visible light to produce the various colors. The hypochlorite ion (OCl -) , found in many

commercial preparations, is reduced to chloride ions and hydroxide ions forming a basic solution as it accepts

electrons from the colored material as shown below.

OCl- + 2e- + HOH --------> Cl- + 2 OH-

Bleaches are often combined with "optical brighteners". These compounds are quite different from bleaches.

They are capable of absorbing wavelengths of ultraviolet light invisible to the human eye, and converting these

wavelengths to blue or blue-green light. The blue or blue-green light is then reflected by the substance making

the fabric appear much "whiter and brighter" as more visible light is seen by the eye.

Thus, bleaches can be classified as either oxidizing agents or reducing agents .

Oxidizing Bleaches

The oxidizing bleaches (and bleaching agents) in common use today are: chlorine, chlorine dioxide, alkaline

hypochlorites, hydrogen peroxide, peroxygen compounds, and sunlight and artificial light[10].

Chlorine Dioxide (ClO2): Chlorine dioxide has been used as a bleaching agent both in its gaseous phase and in

aqueous solution. Because of its explosive nature, chlorine dioxide in the gaseous phase is often diluted with

nitrogen or carbon dioxide. If stored or shipped, chlorine dioxide is passed through cold water and kept under

refrigeration.

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In acidic solution, chlorine dioxide behaves as an oxidizing agent. The complete reduction of ClO 2 is shown in

equation .

ClO2 + 4H + + 5 e − → Cl − + 2H2O          

The individual steps of this overall reduction reaction produce HClO 2 , HOCl, and Cl 2 , which all behave as

oxidizing agents. An acidic medium is required, as ClO2 disproportionates in alkaline solution, as shown in

equation .

2ClO2 + 2OH − → ClO3 − + ClO2

− + H 2 O          

Chlorine dioxide is mainly used for pulp bleaching.

Hypochlorites (OCl − ). Hypochlorite bleach solutions are made from NaOCl and, to a lesser extent, Ca(OCl) 2 .

Hypochlorites are used in laundering, as disinfectants, in the bleaching of pulp and textiles, and in the removal

of ink from recycled paper. Commercial bleaching solutions are obtained by passing chlorine gas through cold,

dilute, aqueous sodium hydroxide, as shown in equation .

Cl2 + 2OH − → OCl − + Cl − + H2O          

To be an effective bleach, the hypochlorite solution should be kept alkaline (pH > 9.0), in order to suppress the

hydrolysis of OCl − and prevent the formation of unstable HOCl.

OCl − + H 2 O → HOCl + OH −          

The active ingredients in hypochlorite bleaches vary with pH. At pH < 2, Cl 2 is the main component

in solution; at pH 4 to 6, HOCl is the dominant species; at pH > 9, OCl − is the only component present. It is the

hypochlorite ion in basic solution that is the active ingredient in household bleach, which is typically about 5 to

6 percent NaOCl. The OCl − ion oxidizes chromophores in colored materials, and is itself reduced to chloride

and hydroxide ions.

OCl− + H2O + 2e− Cl− + 2OH−          

The whitening process effected by commercial hypochlorite bleach is often enhanced by the use of

optical brighteners, compounds that absorb incident ultraviolet light and emit visible light, making the fabric

appear brighter and whiter.

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Hydrogen Peroxide (H2O2)

Hydrogen peroxide, as a bleaching agent used in the pulp and paper industry, has the advantage that it

is nonpolluting. Because of the instability of pure hydrogen peroxide, aqueous solutions are employed in

bleaching. At room temperature, hydrogen peroxide very slowly decomposes to water and oxygen.

2H2O2 H2O + O2          

However, the presence of transition metal cations (particularly Fe 3+ , Mn 2+ , and Cu 2+ ) and other catalysts

dramatically accelerates this reaction. As a result, aqueous hydrogen peroxide must be stabilized with

complexing agents that sequester transition metal cations.

The active bleaching species in hydrogen peroxide is the perhydroxyl anion , OOH − , formed through the

ionization of H 2 O 2 .

H2O2 + H2O → H3O + + OOH −          

The acid ionization constant of hydrogen peroxide is very low ( K a = 2 × 10 −12 ) with the result that solutions of

H 2O2 must be made alkaline in order

Peroxygen Compounds. A number of solid peroxygen compounds that release hydrogen peroxide when

dissolved in water exist. These include sodium perborate (NaBO3 . 4H 2 O or NaBO 2 z H 2 O 2 z 3H 2 O) and

sodium carbonate peroxyhydrate (2Na 2 CO 3 z 3H 2 O 2 ). The structure of sodium perborate contains the

peroxoanion B2(O2)2(OH)42− , which contains two O–O linkages that join two tetrahedral BO 2 (OH) 2− groups.

These peroxygen compounds are used in detergents, denture cleaners, and tooth powders.

Reducing Bleaches

Reducing agents used in bleaching include sulfites, bisulfites, dithionites, and sodium borohydride, all of

which are used in pulp and textile bleaching.

Sulfites (SO32− ) and Bisulfites (HSO 3

− ). The oxidation state of sulfur in both SO32− and HSO3

− is +4, and

oxidation to +6 occurs readily, with the formation of SO42− and HSO4

− , respectively, making sulfites and

bisulfites good reducing agents.

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Dithionites (S2O42− ) . Both sodium and zinc dithionite have found use in the bleaching of mechanical pulps and

textiles. The preparation of the dithionite ion is accomplished via the reduction of the bisulfite ion and sulfur

dioxide with Zn dust.

2HSO3 − + SO2 + Zn Zn2+ + S2O4

2− + SO32− + H2O          

The dithionite ion, S 2 O 4 2− , which has sulfur in the +3 oxidation state, behaves as a strong reducing agent in

alkaline solution.

S2O42− + 4OH − 2SO3

2− + 2 H2O + 2 e−          

As the pH is lowered, the reducing power of the dithionite ion drops off, as predicted by LeChatelier's principle.

Dithionites are useful in removing rust stains, and neutral citrate solutions of Na2S2O4 were used to remove iron

corrosion products from objects recovered from the Titanic.

2.3. Nitrogen Fixation

Nitrogen is the most abundant element in our atmosphere. It is a vital element as many classes of

compounds essential to living systems are nitrogen-containing compounds. Nitrogen is primarily present in the

atmosphere freely as dinitrogen or nitrogen gas. Molecular nitrogen or diatomic nitrogen (N 2) is highly stable as

it is triple bonded (N≡N). Because of this stability, molecular nitrogen as such is not very reactive in the

atmosphere under normal conditions. In the atmosphere molecular nitrogen is 78.03% by volume and it has a

very low boiling point (-195.8oC) which is even lower than oxygen. Proteins present in living organisms contain

about 16% nitrogen.

Nitrogen is a primary nutrient for all green plants, but it must be modified before it can be readily

utilized by most living systems. Nitrogen fixation is one process by which molecular nitrogen is reduced to

form ammonia. In other words the conversion of dinitrogen to ammonia is called nitrogen fixation. Because

ammonia is necessary for the formation of biologically essential, nitrogen containing compounds such as

amino acids and nucleic acids. a fixed nitrogen i.e. ammonia is necessary to sustain life on earth .

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In nature, nitrogen fixation takes place takes place mainly by two important processes such as

biological and abiological process of nitrogen fixation. Biological nitrogen fixation process is carried out by

microorganisms where as abiological nitrogen fixation is done by the abiological sources like lightning.

Nitrogen fixation process is carried out by nitrogen-fixing bacteria present in the soil. This complex

process is a distinctive property possessed by a select group of organisms, because of the presence of the

enzyme nitrogenase in them. The process of biological nitrogen fixation is primarily confined to microbial cells

like bacteria and cyanobacteria. These microorganisms may be independent and free living. Some free living

microbes which fix nitrogen are given below

Organisms Status

Clostridium Anaerobic bacteria (Non photosynthetic)

Klebsiella Facultative bacteria (Non photosynthetic)

Azotobacter Aerobic bacteria (Non photosynthetic)

Rhodospirillum Purple, non-sulphur bacteria (Photosynthetic)

Anabaena Cyanobacteria (Photosynthetic)

Some microbes may become associated with other oragnisms and fix nitrogen. The host organism may be a

lower plant or higher plant. The host organism and the nitrogen fixing microbes establish a special relationship

called symbiosis and this result in symbiotic nitrogen fixation.

Some symbiotic nitrogen fixing organisms are given below

System Symbionts

Lichens Cyanobacteria and Fungus.

Bryophyte Cyanobacteria and Anthoceros.

Pteridophyte Cyanobacteria and Azolla.

Gymnosperm Cyanobacteria and Cycas.

Angiosperms Legumes and Rhizobium.

Angiosperms Non leguminous and actinomycete

Mechanism of Biological Nitrogen Fixation

Nitrogen fixation requires

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(i) The molecular nitrogen

(ii) A strong reducing power to reduce nitrogen like FAD (Flavin adenine dinucleotide)

(iii) A source of energy (ATP) to transfer hydrogen atoms to dinitrogen

(iv) Enzyme nitrogenase

(v) Compound for trapping the ammonia formed since it is toxic to cells.

The reducing agent and ATP are provided by photosynthesis and respiration.The overall biochemical

process involves stepwise reduction of nitrogen to ammonia. The enzyme nitrogenase is a Mo-Fe containing

protein and binds with molecule of nitrogen (N2) at its binding site. This molecule of nitrogen is then acted

upon by hydrogen (from the reduced coenzymes) and reduced in a stepwise manner. It first produces diimide

(N2H2) then hydrazine (N2H4) and finally ammonia (2NH3). NH3 is not liberated by the nitrogen fixers. It is

toxic to the cells and therefore these fixers combine NH3 with organic acids in the cell and form amino acids.[13]

The general equation for nitrogen fixation may be described as follows:

     N2 + 8e- + 8H+  + 16ATP + 16H2O 2NH3 + H2 + 16ADP + 16Pi + 8H+

Some bacterial species, the symbiotic eubacteria Rhizobium (in plant root nodules) and the archaea

cyanobacteria (formerly blue-green algae) contain an enzyme complex for this process. This is the nitrogenase

complex and contains Fe-S and Mo-Fe cofactors for the transfer of electrons from ferredoxin to N2. The process

of nitrogen reduction is extremely energy dependent. The triple bond energy in molecular nitrogen is

225kcal/mol and the industrial production of ammonia requires temperatures of 500 degrees Celsius and a

pressure of 300 atmospheres. Rhizobium uses 8 reducing equivalents and 16 ATPs as shown in the equation

above

This reaction is catalyzed by the hetero-oligomeric protein complex composed of a reductase and a

nitrogenase part. The reductase is a homodimer containing a 4Fe-4S cluster and an ATP binding site at the

subunit interface, which is used to oxidize ferredoxin, which is supplied either by photosynthetic membranes

(PSI) or from oxidative catabolism. The reductase donates 8 electrons in succession to the nitrogenase cofactor,

a molybdenum-iron containing active center, where one molecule of N2 is reduced in the presence of protons to

2 NH3, and H2 as a byproduct. The reduction catalysis is powered by sixteen ATP molecules hydrolyzed by the

reductase subunit. Molecular oxygen is a strong inhibitor of the nitrogenase Mo-Fe cofactor and is removed by

the plant oxygen binding protein leghemoglobin in the root nodules.

Two metalloproteins i.e. larger Mo-Fe-protein and smaller, a Non-Symbiotic N 2 fixation Fe-protein

components are involved in N2 fixation. Fe-protein interacts with ATP and Mg++, and receives an electron from

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ferredoxin or flavodoxin when it is oxidized. Mo-Fe-protein of nitrogenase complex combines with the ecology

of diazotrophs reducible substrates i.e. N2 and yields two molecules NH3. It appears that N2 is reduced step-wise

without breaking N-N bond until the final reduction and production a sites of N2 fixation of ammonia is

accomplished. Finally two molecules of NH3 are released from the enzyme.[2]

Fe-protein (oxidized form) gets electrons from ferredoxin (when it combines with 2H + and yields H 2 ) and

energy from ATP. Mg ++ activates this reaction.

Finally, electron is transferred to oxidize Mo-Fe-protein which becomes reduced and Fe-protein is

oxidized. The reduced form of Mo-Fe-protein combines with N2 and other substrates to result in NH3 and other

various products with respect to substrate. H2 produced during this reaction is further utilized by some

microorganisms which possess hydrogenase. Reutilization of H increases nitrogenase 2 activity by protecting it

from inhibition of H2 .

Nitrogenase Action -

(i) In most organisms, the physiologically functional reductant of nitrogenase is ferredoxin. Other

natural reductants include flavodoxin and NADPH. Artificial reductants include Na2S2O4 and reduced methyl

viologen. Reduced ferredoxin, the electron donor, reduces the Fe protein of nitrogenase.

Fe Protein (Oxidized ) + e- Fe Protein (reduced)

(ii) The reduction of N 2 to NH4 is exothermic. Yet nitrogen fixation requires energy in the form of

adenosine 5'-triphosphate (ATP), because of a high activation energy. The Fe protein of nitrogenase specifically

binds to Mg ATP and lowers its redox potential. A complex containing both Fe and Mo-Fe proteins and Mg

ATP is assembled. In the absence of Mg ATP the midpoint potential (Em) of Fe protein is about -250 to -295

mV. After binding with Mg ATP the Em is about -400 mV.

Fe protein + 2Mg ATP + Mo-Fe protein Fe protein. 2MgATP. Mo-Fe protein

(iii) The Fe protein transfers an electron to the Mo-Fe protein. This results in the oxidation of the Fe

protein and the reduction of the Mo-Fe protein. This reaction is coupled to ATP--->ADP hydrolysis. ATP is

not hydrolyzed to A TP until the Fe protein transfer an electron to the Mo-Fe protein. There are 12 or more

ATP hydrolyzed for each N2 reduced, or 4 A TP per pair of electrons transferred to the Mo- Fe protein or to 19

the substrate.

Thus, there appear to be 2 ATPs hydrolyzed for each electron transferred. In vivo and in vitro ATP

requirements are not necessarily the same. Growth yield experiments indicate that in Azotobacter only 4 or 5

ATPs are required for each N2 fixed. On the other hand 29 ATPs are required in K. pneumoniae and 20 ATPs

in C. pasteurianum.

(iv) The reduced Mo-Fe protein can in turn reduce the substrate. A number of substrates other than N 2can be

reduced by nitrogenase. Both Fe and Mo-Fe proteins are required for all these reductions, which are coupled to

Mg ATP hydrolysis.

Although nitrogen-fixation involves a number of oxidation-reduction reactions that occur sequentially, that

reaction which describes its reduction can be written in a simplified way as:

N2 + 6 e- + 8H+ 2 NH4+ (ammonium ion)

The ammonium ion (the conjugate acid of ammonia, NH3 ) that is produced by this reaction is the form of

nitrogen that is used by living systems in the synthesis of many bio-organic compounds.[15]

It is glutamine synthetase (GS) catalyses the reaction of glutamic acid plus NH 3 and converts into

glutamine, which in turn combines with 2-oxoglutarate and results in two molecules of glutamic acid in the

presence of an enzyme, glutamine oxoglutarate amino transferase (GOGAT). Glutamic acid is the source of

several metabolic products such as amino acids, nucleotides, proteins, etc. The 2-oxoglutaric acid is produced

by combining maltose with CO2 . Also maltose gives rise to glucose-6-phosphate. Its further conversion in

different microorganisms differs. In cyanobacteria, glucose-6-phosphate is converted to ribose 5-phosphate with

an intermediate product 6-phosphogluconate, and produces H+. In bacteria it differs from genus to genus. In

Clostridium pasteurianum, pyruvic acid is produced from glucose-6-phosphate. However, in R. rubrum pyruvic

acid supports nitrogenase activity (Ludden and Burris, 1981) releasing ATP. Thus, nitrogen fixed by

microorganisms is released into their surrounding

20

The presence of MgATP is an absolute requirement for electron transfer from the Fe protein to the MoFe

protein. For MgATP hydrolysis by nitrogenase, on the other hand, it is not necessary that electron transfer take

place: nitrogenase hydrolyses MgATP even when no reductant is present and the Fe protein is oxidized.The

precise mechanism of action of MgATP hydrolysis in nitrogenase catalysis is not yet known.

The transfer of the first electron from the Fe protein to the MoFe protein is accompanied by a fast

increase of the absorbance at 430 nm, due to oxidation of the Fe protein. Following this absorbance increase,

smaller absorbance changes are observed , probably caused by subsequent redox changes of the MoFe protein.

Lowe et al. were able to simulate the absorbance changes that occur during the first 0.6 s of the reaction of the

nitrogenase of Klebsiella pneumoniae.

Another way by which ammonia may be formed is by the process called nitrification. In this process

compounds called nitrates and nitrites, released by decaying organic matter are converted to ammonium ions by

nitrifying bacteria present in the soil. The process carried out by these bacteria is also a complex series of

oxidation-reduction reactions. The reduction reactions involving nitrate and nitrite ions can be simplified as

follows:

NO3- + 2e- + 2H+ NO2

- + H2O

(nitrate ion) (nitrite ion)

 

  NO2- + 6e- + 2H+ NH4

+ + 2 H2O

21

Another way in which molecular nitrogen is modified is via the discharge of lightning. The tremendous energy

released by the electrical discharges in our atmosphere breaks the rather strong bonds between nitrogen atoms,

causing them to react with oxygen. In this process, nitrogen is oxidized and oxygen is reduced.

lightning

N2 + O2 2 NO (nitrous oxide)

The nitrous oxide formed combines with oxygen to form nitrogen dioxide.

2 NO + O2 2NO2

Nitrogen dioxide readily dissolves in water to product nitric and nitrous acids;

2 NO2 + H2O HNO3 + HNO2

These acids readily release the hydrogen forming nitrate and nitrite ions which can be readily utilized by plants

and micro-organisms.

HNO3 H+ + NO3- (nitrate ions)

HNO2 H+ + NO2- (nitrite ions)

Denitrifying bacteria, act on ammonia as well as nitrates produced by death and decay, recycling these

compounds as free nitrogen (N2). The nitrogen that is fixed by the processes described above is eventually

returned to the atmosphere by this denitrification process, to complete what is commonly referred to as the

"nitrogen cycle".

Ammonia is further synthesized into a number of metabolic products in microbial cells. However, ammonia

is not accumulated in the cell, although a few species may create it. Rather it is incorporated into organic forms

by combining with an organic acid (a -keto-glutaric acid) to give rise to amino acid e.g. glutamic acid. The

ammonia may also combine with organic molecules to yield alanin or glutamine.

It has been observed that ammonia formation is achieved by plants either by nitrogen fixation or by

reduction of nitrate to nitrite. Ammonium (NH4+) is the most reduced form of inorganic combined nitrogen.

This ammonium now becomes the major source for the production of amino acids, which are the building

22

blocks of enzymes and proteins. Amino acids have two important chemical groups. (i)amino group (NH) and

(ii) carboxy1 group.

H

|

R — C — COOH

|

NH2

A typical amino acid with functional groups is shown above. R represents alkyl group.Ammonium so produced

is the major source of amino group. However, the carboxyl group has to be provided by other organic molecule

synthesized by the plants. There are two major reactions for amino acid biosynthesis in plants:

Reductive amination reaction:

In this reaction, ammonia combines with a keto acid. The most important keto acid is the alpha

ketoglutaric acid produced during the operation of Krebs cycle. The keto acid then undergoes enzymatic

reductive amination to produce an amino acid.

glutamate dehydrogenase

a-ketoglutaric acid + NH3 Glutamic acid

(keto acid) (amino acid)

Similarly another amino acid called aspartic acid is produced by reductive amination of oxaloacetic acid.

It has been noted that reductive amination represents the major ‘port of entry’ for ammonia into the metabolic

stream in plants. This initiates synthesis of glutamic acid followed by other amino acids.

23

2.4. Metabolism

Metabolism is a general term used to refer to all of the chemical reactions which occur in a living system.

Metabolism can be divided into two parts i.e anabolism and catabolism. Anabolism is reactions involving the

synthesis of compounds and catabolism is reactions involving the breakdown of compounds. In terms of

oxidation-reduction principles, anabolic reactions are primarily characterized by reduction reactions, such as the

dark reaction in photosynthesis where carbon dioxide is reduced to form glucose. Catabolic reactions are

primarily oxidation reactions. Although catabolism involves many separate reactions, an example of such as

process can be described by the oxidation of glucose which is known as respiration. Photosynthesis and

respiration are the important metabolic processes without which we can`t expect life on the earth. Both of these

involves number of oxidation and reduction reactions and are dealt in detail as follows.

2.4.1.Photosynthesis

Plants represent one of the most basic examples of biological oxidation and reduction, i.e. the process

of photosynthesis. It is a very complex process carried out by green plants, blue-green algae, and certain

bacteria. These organisms are able to harness the energy contained in sunlight, and via a series of oxidation-

reduction reactions, produce oxygen and sugar, as well as other compounds which may be utilized for energy as

well as the synthesis of other compounds.

For plants, the upper and lower ends of the visible spectrum are the wavelengths that help drive the

process of splitting water (H2O) during photosynthesis, to release its electrons for the biological reduction of

carbon dioxide (CO2) and the release of diatomic oxygen (O2) to the atmosphere. It is through the process of

photosynthesis that plants are able to use the energy from light to convert carbon dioxide and water into the

chemical energy storage form called glucose.So it is an anabolic process where chemical conversion of carbon

dioxide and water into sugar (glucose) and oxygen takes place, which is a light-driven reduction process:

The overall equation for the light-driven reduction of CO2 may be expressed as:

6 CO2 + 6H2O C6H12O6 + 6 O2

(glucose)

24

The equation is the net result of two processes. One process involves the splitting of water. This process

is really an oxidative process that requires light, and is often referred to as the "light reaction". This reaction

may be written as:

12 H2O 6 O2 + 24 H+ + 24e-

light or radiant energy

The oxidation of water is accompanied by a reduction reaction resulting in the formation of a compound, called

nicotinamide adenine dinucleotide phosphate (NADPH). This reaction is illustrated below:

NADP+ + H20 NADPH + H+ + O

(oxidized form) (reduced form) (oxygen)

This reaction is linked or coupled to yet another reaction resulting in the formation of a highly energetic

compound, called adenosine triphosphate, (ATP). As this reaction involves the addition of a phosphate group

(labeled, as Pi) to a compound called, adenosine diphosphate (ADP) during the light reaction, it is called

photophosphorylation.

ADP + Pi ATP

Think of the light reaction, as a process by which organisms "capture and store" radiant energy as they produce

oxygen gas. This energy is stored in the form of chemical bonds of compounds such as NADPH and ATP.

The energy contained in both NADPH and ATP is then used to reduce carbon dioxide to glucose, a type of

sugar (C6H12O6). This reaction, shown below, does not require light, and it is often referred to as the "dark

reaction".

6CO2 + 24 H+ + 24 e- C6H12O6 + 6 H2O

The process by which non photosynthetic organisms and cells obtain energy, is through the consumption of the

energy rich products of photosynthesis. By oxidizing these products(especially glucose), electrons are passed

along to make the products carbon dioxide and water, in an environmental recycling process. The chemical

bonds present in glucose also contain a considerable amount of potential energy. This stored energy is released

whenever glucose is catabolized (broken down) to drive cellular processes. The carbon skeleton in glucose also

serves as a source of carbon for the synthesis of other important biochemical compounds such as, lipids, amino

acids, and nucleic acids. The process  of oxidation of glucose in presence of atmospheric oxygen produces high

25

energy which is released in plants and is utilized for some metabolic processes . The following reaction

represents this process:

C6H12O6 + O2 6CO2 + 6H2O +Energy                                                

~Energy-yielding oxidation of glucose reaction~

It is therefore through this process that heterotrophs ( plants which consume other organisms obtain to

energy) and autotrophs (able to produce their own energy) participate in an  environmental cycle of

exchanging carbon dioxide and water to produce energy containing glucose for organismal  oxidation and

energy production, and subsequently allowing the regeneration of the byproducts carbon dioxide and water, to

begin the cycle again.  Therefore, these two groups of organisms have been allowed to diverge interdependently

through this natural life cycle.

In simplest terms, the process of photosynthesis can be viewed as one-half of the carbon cycle. In this half,

energy from the sun is captured and transformed into nutrients which can be utilized by higher organisms in the

food chain. The release of this energy during the metabolic re-conversion of glucose to water and carbon

dioxide represents the second half of the carbon cycle and it may be referred to as catabolism or "oxidative

processes".

26

2.4.2. Respiration

Respiration is an oxidative process in which chemically bound energy from complex organic fuel molecules

such as carbohydrates, proteins and fats is captured in the form of ATP. It involves the two processes, one is the

releasing energy stored in the “fuel molecules” by breaking them in series of steps and another is the capturing

the energy released during some of these steps in ATP. Primarily it involves the intake of oxygen from the

atmosphere by the respiratory organs like lungs. Then it is moves into the blood stream where it is attached to

the haemoglobin and transported into the cell.Each molecule of haemoglobin contains four iron metal atoms.

oxygen binds to the Fe of the haemoglobin,where the Fe will remain in Fe2+ state. So it shows that the reaction

is oxygenation and not oxidation. So this process doesn`t involve the any chemical reaction.

Oxygenated blood then transported into the cell where actual catabolic processes taking place. Oxygen is

utilized for the breakdown of sugar molecules which intern produces energy required by the cell. The process is

known as cellular respiration. Cellular respiration is the major energy producing process in living organisms.

This energy is stored in the bonds of ATP, and a maximum of 38 moles of ATP are produced for every mole of

glucose that catabolized. Cellular respiration is a redox process; the carbon atoms in glucose are oxidized while

oxygen atoms in oxygen gas are reduced to the oxygen in water. The net reaction that takes place during cellular

respiration is shown below. This equation is the reverse of the photosynthetic equation.

C6H12O6 + 6 O2 6 CO2 + 6 H2O + Energy

In this reaction, the carbon atoms in glucose are oxidized, undergoing an increase in oxidation state

where each carbon loses 2 electrons, as they are converted to carbon dioxide. At the same time, each oxygen

atom is reduced by gaining 2 electrons when it is converted to water. Part of the energy is released as heat and

the remainder is stored in the chemical bonds of "energetic" compounds such as adenosine triphosphate (ATP)

and nicotinamide adenine dinucleotide (NADH).

Catabolic reactions can be divided into many different groups of reactions called, catabolic pathways. In

these pathways (referred to as Glycolysis, the Citric Acid Cycle, and Electron Transport) the carbon atoms are

slowly oxidized by a series of reactions which gradually modify the carbon skeleton of the compound as well as

the oxidation state of carbon. Coupled to these reactions are other reversible oxidation-reduction reactions

designed to capture the energy released and temporarily store it within the chemical bonds of compounds called

adenosine triphosphate (ATP) and nicotamide dinucleotide (NADH) . These compounds are then utilized to

provide energy for driving the cellular machinery.[21]

27

Electron transfer reactions are fundamental to many metabolic processes necessary for the survival of

all organisms. Copper and iron containing proteins play an important role in electron transfer reaction. The

polypeptide or protein component appears to tune the metal centre that can readily accept or donate electron in

required redox reaction. It has been believed that the protein enables electron to move over appreciable

distances from one redox centre to another. There are number of enzymes which are found carryout redox

reactions in living system. Important of them are Cytochrome P-450, Cytochrome c oxidase, Peroxidases and

super oxide dismutases.[2]

In oxidizing atmosphere, one of the major requirements for life is the maintenance of molecules in

reduced state. Another requirement is the consumption of oxygen for the generation of energy by respiration in

which the essential step is the reduction of oxygen to water.

O2 + 4H+ + 4e- 2H2O

The above reaction is catalyzed by a single enzyme cytochrome oxidase. The importance of this oxygen

consumption process is that it constitutes the terminal reaction of the respiratory chain. This chain provides the

energy needed for the life process of aerobic organisms. The electron transport is coupled with the synthesis of

energy rich molecule i.e. adenosine triphosphate (ATP). The overall result of this process is that electrons are

electrons are transferred in stages from the reduced pyridine nucleotide to oxygen ,through a potential of 1.1V

28

O2 + e- O2-

O2- + e- O22-

The reduction of O2 can occur in stages to give peroxide initially, which is radical anion, followed by

the peroxide dianion. These are highly reactive and toxic species and they can be removed by enzymes,

superoxide dismutase, catalase and peroxidase. Copper, zinc superoxide dismutase (Cu, Zn – SOD) protein

functions as catalyst of superoxide (O2-) disproportionation ,i.e. a superoxide dismutase

Cytochrome P-450 is a member of group of enzymes which catalyze the addition of molecular oxygen

to the substrate via oxygen activation. It has attracted considerable attention because it catalyzes the

hydroxylation of the substrate RH at the consumption of molecular oxygen by the reductive cleavage of O-O

bond.

R-H + O2 2e,2H+ R-OH + H2O

Cytochrome P-450 is found in plants, animals and bacteria and participates in numerous metabolic

pathways. In humans, different forms of microsomal P-450 are believed to catalyse the hydroxylation of drugs,

steroid precursors, pesticides and other foreign substances. Cytochrome is a part of body`s detoxification

system.

FERMENTATION

During cellular respiration, glucose is completely oxidized and oxygen gas is required to act as the

oxidizing agent. Cells can extract energy from glucose in the absence of oxygen but not nearly as efficiently.

Without oxygen only a fraction of chemical energy of glucose can be released. Whereas cellular respiration

produces 38 moles of ATP for every mole of glucose catabolized in the absence of oxygen. This provides

enough energy for the oxygen deprived cells so that they don’t die. The process in which glucose is broken

down in the absence of oxygen is known as fermentation. There are two common kinds of fermentation. In one,

ethanol and carbon dioxide are produced. In other lactic acid is produced.

Alcoholic fermentation

Yeast and some bacteria can ferment glucose to produce the alcohol i.e. ethanol. Here first glucose is

converted to pyruvic acid by the glycolysis. The pyruvic acid undergo decarboxylation to form acetaldehyde

which further undergo reduction to give ethanol and carbon dioxide with liberation of energy.

29

C6H12O6 2CH3COCOOH 2CH3CH2OH + 2CO2 +Energy

Glucose Pyruvic acid Ethanol carbon dioxide

This reaction, called alcoholic fermentation is important to certain segments of the food industry.

Alcoholic fermentations is needed to make bread dough rise, form tofu from soybeans and produce the ethanol

in alcoholic beverages. Another use of ethanol that is produced by yeast is as an additive to gasoline, called

gasohol.

Lactic acid fermentation

During strenuous activity muscle cells often use oxygen faster than it can be supplied by the blood. When the

supply of oxygen is depleted, cellular respiration stops. Although animal cells can’t undergo alcoholic

fermentation they can produce lactic acid and a small amount of energy from glucose through lactic acid

fermentation. This reaction also follows the same pathway as that of alcohol fermentation only difference is that

it doesn`t involves the formation of acetaldehyde,so that pyruvic acid directly reduced to lactic acid.

C6H12O6 2CH3COCOOH 2CH3CH(OH)COOH + Energy

Glucose Pyruvic acid Lactic acid

The lactic acid is produced is moved from the muscle through the blood to the liver. There it is

converted back into glucose that can be used in catabolic processes to yield more energy once oxygen becomes

available. However if lactic acid builds up in muscle cells at a faster rate than the blood can remove it, muscle

fatigue results. Build up of lactic acid is what causes a burning pain in the muscle during strenuous exercise.

2.5. Electrochemical Cells

30

Many oxidation-reduction reactions occur spontaneously, giving off energy. An example involves the

spontaneous reaction that occurs when zinc metal is dipped in a solution of copper ions as described by the net

ionic equation shown below.

Cu+2 (aq) + Zn (s) Cu(s) + Zn+2 (aq)

The zinc metal slowly dissolves as its oxidation produces zinc ions which enter into solution. At the same time,

the copper ions gain electrons and are converted into copper atoms which coats the zinc metal or sediments to

the bottom of the container. The energy produced in this reaction is quickly dissipated as heat, but it can be

made to do useful work by a device called, an electrochemical cell. This is done in the following way.

An electrochemical cell is composed to two compartments or half-cells, each composed of an electrode dipped

in a solution called electrolyte. These half-cells are designed to contain the oxidation half-reaction and

reduction half-reaction separately as shown below.

The half-cell, called the anode, is the site at which the oxidation of zinc occurs as shown below.

Zn (s) Zn+2 (aq) + 2e-

During the oxidation of zinc, the zinc electrode will slowly dissolve to produce zinc ions (Zn+2), which enter

into the solution containing Zn+2 (aq) and SO4-2 (aq) ions.

The half-cell, called the cathode, is the site at which reduction of copper occurs as shown below.

Cu+2 (aq) + 2e- Cu (s)

When the reduction of copper ions (Cu+2) occurs, copper atoms accumulate on the surface of the solid copper

electrode.The reaction in each half-cell does not occur unless the two half cells are connected to each other.

31

In order for oxidation to occur, there must be a corresponding reduction reaction that is linked or

coupled with it. Moreover, in an isolated oxidation or reduction half-cell, an imbalance of electrical charge

would occur, the anode would become more positive as zinc cations are produced, and the cathode would

become more negative as copper cations are removed from solution. This problem can be solved by using a

"salt bridge" connecting the two cells as shown in the diagram below. A "salt bridge" is a porous barrier which

prevents the spontaneous mixing of the aqueous solutions in each compartment, but allows the migration of ions

in both directions to maintain electrical neutrality. As the oxidation-reduction reaction occurs, cations ( Zn+2)

from the anode migrate via the salt bridge to the cathode, while the anion, (SO4)-2, migrates in the opposite

direction to maintain electrical neutrality.

The two half-cells are also connected externally. In this arrangement, electrons provided by the oxidation

reaction are forced to travel via an external circuit to the site of the reduction reaction. The fact that the reaction

occurs spontaneously once these half cells are connected indicates that there is a difference in potential energy.

This difference in potential energy is called an electomotive force (emf) and is measured in terms of volts. The

zinc/copper cell has an emf of about 1.1 volts under standard conditions.

Any electrical device can be "spliced" into the external circuit to utilize this potential energy produced by

the cell for useful work. Although the energy available from a single cell is relatively small, electrochemical

cells can be linked in series to boost their energy output. A common and useful application of this characteristic

is the "battery". An example is the lead-acid battery used in automobiles. In the lead-acid battery, each cell has a

lead metal anode and lead (IV) oxide (lead dioxide) cathode both of which are immersed in a solution of

sulfuric acid. This single electrochemical cell produces about 2 volts. Linking 6 of these cells in series produces

the 12-volt battery found in most cars today. One disadvantage of these "wet cells" such as the lead-acid battery

32

is that it is very heavy and bulky. However, like many other "wet cells", the oxidation-reduction reaction which

occurs can be readily reversed via an external current such as that provided by an automobile's alternator. This

prolongs the lifetime and usefulness of such devices as an energy source.

The "Dry-Cell" Battery

The most common type of battery used today is the "dry cell" battery. There are many different types

of batteries ranging from the relatively large "flashlight" batteries to the miniaturized versions used for wrist

watches or calculators. Although they vary widely in composition and form, they all work on the sample

principle. A "dry-cell" battery is essentially comprised of a metal electrode or graphite rod (elemental carbon)

surrounded by a moist electrolyte paste enclosed in a metal cylinder as shown below.[7]

In the most common type of dry cell battery, the cathode is composed of a form of elemental carbon

called graphite, which serves as a solid support for the reduction half-reaction. In an acidic dry cell, the

reduction reaction occurs within the moist paste comprised of ammonium chloride (NH4Cl) and manganese

dioxide (MnO2):

2 NH4+ + 2 MnO2 + 2e- Mn2O3 + 2 NH3 + H2O

A thin zinc cylinder serves as the anode and it undergoes oxidation:

Zn (s) Zn+2 + 2e-

33

This dry cell "couple" produces about 1.5 volts. ( These "dry cells" can also be linked in series to boost

the voltage produced). In the alkaline version or "alkaline battery", the ammonium chloride is replaced by KOH

or NaOH and the half-cell reactions are:

Zn + 2 OH- ZnO + H2O + 2e-

2 MnO2 + 2e- + H2O Mn2O3 + 2 OH-

The alkaline dry cell lasts much longer as the zinc anode corrodes less rapidly under basic conditions than

under acidic conditions.

Other types of dry cell batteries are the silver battery in which silver metal serves as an inert cathode to

support the reduction of silver oxide (Ag2O) and the oxidation of zinc (anode) in a basic medium. The type of

battery commonly used for calculators is the mercury cell. In this type of battery, HgO serves as the oxidizing

agent (cathode) in a basic medium, while zinc metal serves as the anode. Another type of battery is the

nickel/cadmium battery, in which cadmium metal serves as the anode and nickel oxide serves as the cathode in

an alkaline medium. Unlike the other types of dry cells described above, the nickel/cadmium cell can be

recharged like the lead-acid battery.

Wet cell :

A wet cell is composed of a copper and zinc strip, called electrodes. The dilute sulfuric acid found in a

wet cell is the electrolyte. Electrolyte is a liquid that has the ability to conduct electricity. A wet cell battery has

a liquid electrolyte. Other names are flooded cell, since the liquid covers all internal parts, or vented cell, since

gases produced during operation can escape to the air. Wet cells were a precursor to dry cells and are commonly

used as a learning tool for electrochemistry. It is often built with common laboratory supplies, such as beakers,

for demonstrations of how electrochemical cells work. A particular type of wet cell known as a concentration

cell is important in understanding corrosion. Wet cells may be primary cells (non-rechargeable) or secondary

cells (rechargeable). Originally, all practical primary batteries such as the Daniell cell were built as open-topped

glass jar wet cells. Other primary wet cells are the Leclanche cell, Grove cell, Bunsen cell, Chromic acid cell,

Clark cell, and Weston cell. The Leclanche cell chemistry was adapted to the first dry cells. Wet cells are still

used in automobile batteries and in industry for standby power for switchgear, telecommunication or large

uninterruptible power supplies, but in many places batteries with gel cells have been used instead. These

applications commonly use lead-acid or nickel-cadmium cells.

34

The chemical process which produces electricity in a Leclanché cell begins when zinc atoms on the surface

of the anode oxidize, ie they give up both their electrons to become positively-charged ions. As the zinc ions

move away from the anode, leaving their electrons on its surface, the anode becomes more negatively charged

than the cathode.When the cell is connected in an external electrical circuit, the excess electrons on the zinc

anode flow through the circuit to the carbon rod, the movement of electrons forming an electrical current.

When the electrons enter the rod, they combine with manganese dioxide and water, which react with each

other to produce manganese oxide and negatively charged hydroxide ions. This is accompanied by a secondary

reaction in which the negative hydroxide ions react with positive ammonium ions in the ammonium chloride

electrolyte to produce molecules of ammonia and water.

Zn(s) + 2 MnO2(s) + 2 NH4Cl(aq) ZnCl2 + Mn2O3(s) + 2 NH3(aq) + H2O

Alternately, the reaction proceeds further, the hydroxide ions reacting also with the manganese oxide to form

manganese hydroxide.

Zn(s) + 2 MnO2(s) + 2 NH4Cl(aq) + 2H2O(l) ZnCl2 + 2Mn(OH)3(s) + 2 NH3(aq)

2.6.Corrosion35

Millions of dollars are lost each year because of corrosion. Much of this loss is due to the corrosion of iron and

steel, although many other metals may corrode as well. The problem with iron as well as many other metals is

that the oxide formed by oxidation does not firmly adhere to the surface of the metal and flakes off easily

causing "pitting". Extensive pitting eventually causes structural weakness and disintegration of the metal.But in

certain metals such as aluminium form a very tough oxide coating which strongly bonds to the surface of the

metal preventing the surface from further exposure to oxygen and corrosion.[7]

Corrosion occurs in the presence of moisture. For example when iron is exposed to moist air, it reacts with

oxygen to form rust,

Fe2O3.xH2O

The amount of water complexed with the iron (III) oxide (ferric oxide) varies as indicated by the letter "X". The

amount of water present also determines the color of rust, which may vary from black to yellow to orange

brown. The formation of rust is a very complex process which is thought to begin with the oxidation of iron to

ferrous (iron "+2") ions.

Fe -------> Fe+2 + 2 e-

Both water and oxygen are required for the next sequence of reactions. The iron (+2) ions are further oxidized

to form ferric ions (iron "+3") ions.

Fe+2 ------------> Fe+3 + 1 e-

The electrons provided from both oxidation steps are used to reduce oxygen as shown.

O2 (g) + 2 H2O + 4e- ------> 4 OH-

The ferric ions then combine with oxygen to form ferric oxide [iron (III) oxide] which is then hydrated with

varying amounts of water. The overall equation for the rust formation may be written as :

The formation of rust can occur at some distance away from the actual pitting or erosion of iron as

illustrated below. This is possible because the electrons produced via the initial oxidation of iron can be

conducted through the metal and the iron ions can diffuse through the water layer to another point on the metal

36

surface where oxygen is available. This process results in an electrochemical cell in which iron serves as the

anode, oxygen gas as the cathode, and the aqueous solution of ions serving as a "salt bridge" as shown below.

The involvement of water accounts for the fact that rusting occurs much more rapidly in moist conditions

as compared to a dry environment such as a desert. Many other factors affect the rate of corrosion. For example

the presence of salt greatly enhances the rusting of metals. This is due to the fact that the dissolved salt

increases the conductivity of the aqueous solution formed at the surface of the metal and enhances the rate of

electrochemical corrosion. This is one reason why iron or steel tend to corrode much more quickly when

exposed to salt (such as that used to melt snow or ice on roads) or moist salty air near the ocean.[1]

2.7.Enzymatic browning

37

Enzymatic browning is a chemical process which occurs in fruits and vegetables by the enzyme

tyrosinase or polyphenoloxidase, which results in brown pigments. Enzymatic browning can be observed in

fruits like apricots, pears, bananas, grapes and also in vegetables like potatoes, mushrooms, lettuce and also in

seafood i.e.shrimps, spiny lobsters and crabs.For example in apple enzyme reacts with oxygen and iron-

containing phenols that are also found in the apple. The oxidation reaction basically forms a sort of rust on the

surface of the fruit. You see the browning when the fruit is cut or bruised because these actions damage the cells

in the fruit, allowing oxygen in the air to react with the enzyme and other chemicals.

Enzymatic browning is detrimental to quality, particularly in post-harvest storage of fresh fruits, juices and

some shellfish. Enzymatic browning may be responsible for up to 50% of all losses during fruit and vegetables

production.

On the other hand enzymatic browning is essential for the colour and taste of tea, coffee and chocolate.

Polyphenols – main components in enzymatic browning

Polyphenols , also called phenolic compounds, are group of chemical substances present in plants

(fruits, vegetables) which play an important role during enzymatic browning, because they are substrates for the

browning-enzymes. Phenolic compounds are responsible for the colour of many plants, such as apples, they are

part of the taste and flavour of beverages (apple juice, tea), and are important anti-oxidants in plants.

Polyphenols are normally complex organic substances, which contain more than one phenol group (carbolic

acid):

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Structure of Theaflavin, a polyphenol in tea

Polyphenols can be divided into many different sub categories, such as anthocyans (colours in fruits),

flavonoids (catechins, tannins in tea and wine) and non-flavonoids components (gallic acid in tea leaves).

Flavonoids are formed in plants from the aromatic amino acids phenylalanine and tyrosine.

The colour of apples is due to polyphenols

During food processing and storage many polyphenols are unstable due to the fact that they undergo

chemical and biochemical reactions. The most important is enzymatic oxidation causing browning of

vegetables, fruits. This reaction mostly occurs after cutting or other mechanical treatment of product due to

breaking cells.

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Table 1 : An overview of known polyphenols involved in browning  

Source Phenolic substrates

Apple chlorogenic acid (flesh), catechol, catechin (peel), caffeic acid, 3,4-dihydroxyphenylalanine

(DOPA), 3,4-dihydroxy benzoic acid, p-cresol, 4-methyl catechol, leucocyanidin, p-coumaric

acid, flavonol glycosides

Apricot isochlorogenic acid, caffeic acid, 4-methyl catechol, chlorogenic acid, catechin, epicatechin,

pyrogallol, catechol, flavonols, p-coumaric acid derivatives

Avocado 4-methyl catechol, dopamine, pyrogallol, catechol, chlorogenic acid, caffeic acid, DOPA

Banana 3,4-dihydroxyphenylethylamine (Dopamine), leucodelphinidin, leucocyanidin

Cacao catechins, leucoanthocyanidins, anthocyanins, complex tannins

Coffee beans chlorogenic acid, caffeic acid

Eggplant chlorogenic acid, caffeic acid, coumaric acid, cinnamic acid derivatives

Grape catechin, chlorogenic acid, catechol, caffeic acid, DOPA, tannins, flavonols, protocatechuic

acid, resorcinol, hydroquinone, phenol

Lettuce tyrosine, caffeic acid, chlorogenic acid derivatives

Lobster tyrosine

Mango dopamine-HCl, 4-methyl catechol, caffeic acid, catechol, catechin, chlorogenic acid, tyrosine,

DOPA, p-cresol

Mushroom tyrosine, catechol, DOPA, dopamine, adrenaline, noradrenaline

Peach chlorogenic acid, pyrogallol, 4-methyl catechol, catechol, caffeic acid, gallic acid, catechin,

dopamine

Pear chlorogenic acid, catechol, catechin, caffeic acid, DOPA, 3,4-dihydroxy benzoic acid, p-cresol

Plum chlorogenic acid, catechin, caffeic acid, catechol, DOPA

Potato chlorogenic acid, caffeic acid, catechol, DOPA, p-cresol, p-hydroxyphenyl propionic acid, p-

hydroxyphenyl pyruvic acid, m-cresol

Shrimp tyrosine

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Sweet potato chlorogenic acid, caffeic acid, caffeylamide

Tea flavanols, catechins, tannins, cinnamic acid derivatives

Polyphenoloxidase (PPO, phenolase)

Polyphenoloxidases are a class of enzymes that were first discovered in mushrooms and are widely distributed

in nature. They appear to reside in the plastids and chloroplasts of plants, although freely existing in the

cytoplasm of senescing or ripening plants. Polyphenoloxidase is thought to play an important role in the

resistance of plants to microbial and viral infections and to adverse climatic conditions. Polyphenoloxidase also

occurs in animals and is thought to increase disease resistance in insects and crustaceans. In the presence of

oxygen from air, the enzyme catalyzes the first steps in the biochemical conversion of phenolics to produce

quinones, which undergo further polymerization to yield dark, insoluble polymers referred to as melanins.

These melanins form barriers and have antimicrobial properties which prevent the spread of infection or

bruising in plant tissues. Plants, which exhibit comparably high resistance to climatic stress, have been shown to

possess relatively higher polyphenoloxidase levels than susceptible varieties.

Polyphenoloxidase catalyses two basic reactions: hydroxylation and oxidation. Both reactions utilize molecular

oxygen (air) as a co-substrate. The reaction is not only dependent on the presence of air, but also on the pH

(acidity). The reaction does not occur at acid (pH <5) or alkaline (pH >8) conditions

An example of the formation of melanins from a simple polyphenol, tyrosine, is shown in the figure below:41

2.8.Water Purification

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Chlorination is the most common method for purification of water. Chlorine is one of the most

versatile chemicals used in water and wastewater treatment. Chlorine action on the microbes present in the

water is based on oxidation -reduction mechanism. This powerful oxidizing agent is used for several purposes

such as Disinfection, control of microorganism, control of taste and odor, hydrogen sulphide oxidation, iron

and manganese oxidation etc.

In chemically pure water, molecular chlorine reacts with water and rapidly hydrolyzes to hypochlorous

acid (HOCl) and hydrochloric acid (HCl):

Cl2 + H2O HOCl + HCl

Chlorine water hypochlorous acid hydrochloric acid

Here chlorine is oxidized i.e.oxidaytion number of chlorine changes from 0 to +1, so it can acts as an

oxidizing agent this property of chlorine make this to act as disinfectant.

Both of the acids formed by hydrolysis react with alkalinity to reduce buffering capacity of water and

lower pH. Every pound of chlorine gas added to water removes about 1.4 lb of alkalinity. In cooling water

systems, this alkalinity reduction can have a major impact on corrosion rates.

Hypochlorous acid is a weak acid and dissociates to form a hydrogen ion and a hypochlorite ion.

HOCl H+ + OCl-

Hypochlorous acid Hydrogen ion hypochlorite ion

The primary oxidizing agents in water are hypochlorous acid and the hypochlorite ion, although

hypochlorite has a lower oxidizing potential.i.e. a measure of the tendency of chlorine to react with other

materials. The speed at which these reactions occur is determined by pH, temperature, and oxidation or

reduction potential. The oxidation reactions of chlorine with inorganic reducing agents such as sulfides,

sulfites, and nitrites are generally very rapid. Some dissolved organic materials also react rapidly with chlorine,

but the completion of many organic-chlorine reactions can take hours. The antimicrobial efficacy of

hypochlorous acid (HOCl) is much greater than any of the chloramines.

Besides these beneficial aspects of chlorine as disinfectant,they are known to have some toxic effects on living

beings. In order to overcome these problems chlorine dioxide was discovered in 1814 by Sir Humphrey Davy.

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Chlorine dioxide is very different from elementary chlorine, both in its chemical structure as in its

behavior. Chlorine dioxide is a small, volatile and very strong molecule. In diluted, watery solutions chlorine

dioxide is a free radical. At high concentrations it reacts strongly with reducing agents. Chlorine dioxide is an

unstable gas that dissociates into chlorine gas (Cl2), oxygen gas (O2) and heat. When chlorine dioxide is photo-

oxidized by sunlight, it falls apart. The end-products of chlorine dioxide reactions are chloride (Cl -), chlorite

(ClO-) and chlorate (ClO3-).

Chlorine dioxide gas is used to sterilize medical and laboratory equipment, surfaces, rooms and tools.

Chlorine dioxide can be used as oxidizer or disinfectant. It is a very strong oxidizer and it effectively kills

pathogenic microorganisms such as fungi, bacteria and viruses. It also prevents and removes bio film. As a

disinfectant and pesticide it is mainly used in liquid form. Chlorine dioxide can also be used against anthrax,

because it is effective against spore-forming bacteria.

Chlorine dioxide as an oxidizer

As an oxidizer chlorine dioxide is very selective. It has this ability due to unique one-electron exchange

mechanisms. Chlorine dioxide attacks the electron-rich centers of organic molecules. One electron is transferred

and chlorine dioxide is reduced to chlorite (ClO2- ).

Figure 2: chlorine dioxide is more selective as an oxidizer than chlorine. While dosing the same concentrations,

the residual concentration of chlorine dioxide is much higher with heavy pollution than the residual

concentration of chlorine.

By comparing the oxidation strength and oxidation capacity of different disinfectants, one can conclude

that chlorine dioxide is effective at low concentrations. Chlorine dioxide is not as reactive as ozone or chlorine

and it only reacts with sulphuric substances, amines and some other reactive organic substances. In comparison

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to chlorine and ozone, less chlorine dioxide is required to obtain an active residual disinfectant. It can also be

used when a large amount of organic matter is present.

The oxidation strength describes how strongly an oxidizer reacts with an oxidizable substance. Ozone has

the highest oxidation strength and reacts with every substance that can be oxidized. Chlorine dioxide is weak, it

has a lower potential than hypochlorous acid or hypobromous acid.

The oxidation capacity shows how many electrons are transferred at an oxidation or reduction reaction.

The chlorine atom in chlorine dioxide has an oxidation number of +4. For this reason chlorine dioxide accepts 5

electrons when it is reduced to chloride. When we look at the molecular weight, chlorine dioxide contains 263

% 'available chlorine'; this is more than 2,5 times the oxidation capacity of chlorine.

Table 2: the oxidation potentials of various oxidants.

oxidant oxidation strength oxidation capacity

ozone (O3) 2,07 2 e-

hydrogen peroxide (H2O2) 1,78 2 e-

hypochlorous acid (HOCl) 1,49 2 e-

hypobromous acid (HOBr) 1,33 2 e-

chlorine dioxide (ClO2) 0,95 5 e-

The following comparisons show what happens when chlorine dioxide reacts. First, chlorine dioxide takes up an

electron and reduces to chlorite:

ClO2 + e- ClO2-

The chlorite ion is oxidized and becomes a chloride ion:

ClO2- + 4H+ + 4e- Cl- + 2H2O

These comparisons suggest that chlorine dioxide is reduced to chloride, and that during this reaction it

accepts 5 electrons. The chlorine atom remains, until stable chloride is formed. This explains why no

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chlorinated substances are formed. When chlorine reacts it does not only accept electrons; it also takes part in

addition and substitution reactions. During these reactions, one or more chlorine atoms are added to the foreign

substance.

Contrary to chlorine, chlorine dioxide does not react with ammonia nitrogen (NH3) and hardly reacts with

elementary amines. It does oxidize nitrite (N02) to nitrate (NO3). It does not react by breaking carbon

connections. No mineralization of organic substances takes place. At neutral pH or at high pH values, sulphuric

acid (H2SO3) reduces chlorine dioxide to chlorite ions (ClO2-). Under alkalic circumstances chlorine dioxide is

broken down to chlorite and chlorate (ClO3-).

2ClO2 + 2OH- H2O + ClO3- + ClO2

-

This reaction is catalyzed by hydrogen (H+) ions. The half life of watery solutions of chlorine dioxide

decreases at increasing pH values. At low pH, chlorine dioxide is reduced to chloride ions (Cl- ).

Pure chlorine dioxide gas that is applied to water produces less disinfection byproducts than oxidators, such as

chlorine. Contrary to ozone (O3), pure chlorine dioxide does not produce bromide (Br-) ions into bromate ions

(BrO3-), unless it undergoes photolysis. Additionally chlorine dioxide does not produce large amounts of

aldehydes, ketons, ketonic acids or other disinfection byproducts that originate from the ozonisation of organic

substances.

Disinfecting property of chlorine dioxide:

Drinking water treatment is the main application of disinfection by chlorine dioxide. Chlorine dioxide

is also used in other branches of industry today. Example are sewage water disinfection, industrial process water

treatment, cooling tower water disinfection, industrial air treatment, mussel control, foodstuffs production and

treatment, industrial waste oxidation and gas sterilization of medical equipment.

Chlorine dioxide disinfects through oxidation. It is the only biocide that is a molecular free radical. It

has 19 electrons and has a preference for substances that give off or take up an electron. Chlorine dioxide only

reacts with substances that give off an electron. Chlorine, oppositely, adds a chlorine atom to or substitutes a

chlorine atom from the substance it reacts with.

Substances of organic nature in bacterial cells react with chlorine dioxide, causing several cellular

processes to be interrupted. Chlorine dioxide reacts directly with amino acids and the RNA in the cell. It is not

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clear whether chlorine dioxide attacks the cell structure or the acids inside the cell. The production of proteins is

prevented. Chlorine dioxide affects the cell membrane by changing membrane proteins and fats and by

prevention of inhalation.

When bacteria are eliminated, the cell wall is penetrated by chlorine dioxide. Viruses are eliminated in a

different way; chlorine dioxide reacts with peptone, a water-soluble substance that originates from hydrolisis of

proteins to amino acids. Chlorine dioxide kills viruses by prevention of protein formation. Chlorine dioxide is

more effective against viruses than chlorine or ozone.

Chlorine dioxide is one of a number of disinfectants that are effective against Giardia Lambia and

Cryptosporidium parasites, which are found in drinking water and induce diseases called 'giardiasis' and

'cryptosporidiosis'. The best protection against protozoan parasites such as these is disinfection by a

combination of ozone and chlorine dioxide.

Chlorine dioxide as a disinfectant has the advantage that it directly reacts with the cell wall of

microorganisms. This reaction is not dependent on reaction time or concentration. In contrast to non-oxidizing

disinfectants, chlorine dioxide kills microorganisms even when they are inactive. Therefore the chlorine dioxide

concentration needed to effectively kill microorganisms is lower than non-oxidizing disinfectant concentrations.

Microorganisms cannot built up any resistance against chlorine dioxide.

Chlorine dioxide remains gaseous in solution. The chlorine dioxide molecule is powerful and has the

ability to go through the entire system. Chlorine dioxide can penetrate the slime layers of bacteria, because

chlorine dioxide easily dissolves, even in hydrocarbons and emulsions. Chlorine dioxide oxidizes the

polysaccharide matrix that keeps the bio film together. During this reaction chlorine dioxide is reduced to

chlorite ions. These are divided up into pieces of bio film that remain steady. When the bio film starts to grow

again, an acid environment is formed and the chlorite ions are transformed into chlorine dioxide. This chlorine

dioxide removes the remaining bio film.

The reaction process of chlorine dioxide with bacteria and other substances takes place in two steps.

During this process disinfection byproducts are formed that remain in the water. In the first stage the chlorine

dioxide molecule accepts an electron and chlorite is formed (ClO3). In the second stage chlorine dioxide accepts

4 electrons and forms chloride (Cl-). In the water some chlorate (ClO3), which is formed by the production of

chlorine dioxide, can also be found. Both chlorate and chlorite are oxidizing agents. Chlorine dioxide, chlorate

and chlorite dissociate into sodium chloride (NaCl). In the 1950's the biocidal capability of chlorine dioxide,

especially at high pH values, was known. For drinking water treatment it was primary used to remove inorganic

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components, for example manganese and iron, to remove tastes and odors and to reduce chlorine related

disinfection byproducts.

For drinking water treatment chlorine dioxide can be used both as a disinfectant and as an oxidizing

agent. It can be used for both pre-oxidation and post-oxidation steps. By adding chlorine dioxide in the pre-

oxidation stage of surface water treatment, the growth of algae and bacteria can be prevented in the following

stages. Chlorine dioxide oxidizes floating particles and aids the coagulation process and the removal of turbidity

from water.

Chlorine dioxide is a powerful disinfectant for bacteria and viruses. The byproduct, chlorite (ClO 2-), is a

weak bactericidal agent. In water chlorine dioxide is active as a biocide for at least 48 hours, its activity probaly

outranges that of chlorine.Chlorine dioxide prevents the growth of bacteria in the drinking water distribution

network. It is also active against the formation of bio film in the distribution network. Bio film is usually hard to

defeat. It forms a protective layer over pathogenic microorganisms. Most disinfectants cannot reach those

protected pathogens. However, chlorine dioxide removes bio films and kills pathogenic microorganisms.

Chlorine dioxide also prevent bio film formation, because it remains active in the system for a long time.

For the pre- oxidation and reduction of organic substances between 0,5 and 2 mg/L of chlorine dioxide is

required at a contact time between 15 and 30 minutes. Water quality determines the required contact time. For

post- disinfection, concentrations between 0,2 and 0,4 mg/L are applied. The residual byproduct concentration

of chlorite is very low and there are no risks for human health.

Application in disinfecting swimming pools and cooling towers :

For swimming pool disinfection the combination of chlorine (Cl2) and chlorine dioxide (ClO2) can be

applied. Chlorine dioxide is added to the water. Chlorine is already present in the water as hypochlorous acid

(HOCl) and hypochlorite ions (OCl-). Chlorine dioxide breaks down substances, such as phenols. The

advantages of chlorine dioxide are that it can be used at low concentrations to disinfect water, that it hardly

reacts with organic matter, and that little disinfection byproducts are formed.

Chlorine dioxide is used to disinfect the water that flows through cooling towers. It also removes bio

films and prevents bio film formation in cooling towers. The removal of bio film prevents damage to and

corrosion of equipment and piping and causes the pumping efficiency to be improved. Chlorine dioxide is also

effective in removing Legionella bacteria. The circumstances in cooling towers are ideal for the growth of

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Legionella bacteria. Chlorine dioxide has the advantage that it is effective at a pH between 5 and 10 and that no

acids are required to adjust the pH.

The amount of disinfectant required needs to be determined first. This amount can be determined by

adding disinfectant to the water and measuring the amount that remains after a defined contact time. The

amount of chlorine dioxide that is dosed depends upon the contact time, the pH, the temperature and the amount

of pollution that is present in the water.

Advantages of the use of chlorine dioxide :

The interest in the use of chlorine dioxide as an alternative for or addition to chlorine for the disinfection

of water has increased in the last few years. Chlorine dioxide is a very effective bacterial disinfectant and it is

even more effective than chlorine for the disinfection of water that contains viruses. Chlorine dioxide has

regained attention because it is effectively deactivates the chlorine-resistant pathogens Giardia and

Cryptosporidium. Chlorine dioxide removes and prevents bio film. Disinfection with chlorine dioxide does not

cause odor nuisance. It destroys phenols, which can cause odor and taste problems. Chlorine dioxide is more

effective for the removal of iron and manganese than chlorine, especially when these are found in complex

substances.

The use of chlorine dioxide instead of chlorine prevents the formation of harmful halogenated disinfection

byproducts, for example trihalomethanes and halogenated acidic acids. Chlorine dioxide does not react with

ammonia nitrogen, amines or other oxidizable organic matter. Chlorine dioxide removes substances that can

form trihalomethanes and improves coagulation. It does not oxidize bromide into bromine. When bromide

containing water is treated with chlorine or ozone, bromide is oxidized into bromine and hypobromous acid.

After that these react with organic material to form brominated disinfection byproducts, for example

bromoform.

The use of chlorine dioxide reduces the health risk of microbial pollutions in water and at the same time

decreases the risk of chemical pollutions and byproducts. Chlorine dioxide is a more effective disinfectant than

chlorine, causing the required concentration to kill microorganisms to be much lower. The required contact time

is also very low.

Influence of the pH value on chlorine dioxide efficiency :

Contrary to chlorine, chlorine dioxide is effective at a pH of between 5 and 10. The efficiency increases at

high pH values, while the active forms of chlorine are greatly influenced by pH. Under normal circumstances

49

chlorine dioxide does not hydrolyze. This is why the oxidation potential is high and the disinfection capacity is

not influenced by pH. Both temperature and alkalinity of the water do not influence the efficiency. At the

concentrations required for disinfection, chlorine dioxide is not corrosive. Chlorine dioxide is more water-

soluble than chlorine.

2.9 Aging

Oxidation reactions occur when life essential oxygen combusts within the human body and produces by

products referred to as oxygen free radicals. When an oxidation reaction occurs in metals such as iron, we are

aware that "rusting" occurs. When this same process occurs in living system, it is called aging.

Free radicals produced by oxidation reactions are incomplete molecules that have lost an electron. When

an oxygen molecule loses an electron, it is called singlet oxygen because only one of its electrons remains.

Oxygen in this state is not stable. In an attempt restore balance, the free radical tries to steal an electron away

from a nearby molecule, or donate its remaining electron to a nearby molecule. In doing so, the radical creates

molecular instability that damages, disrupts, and even destroys nearby cells. If DNA is involved, the problem

intensifies and genetic cell mutations may occur (a theory for the common cause of cancer). Uninhibited over

time, free radical damage builds in the body, thus causing aging.

Free radicals are not only produced inside our bodies, but free radicals are also ingested through

smoking, eating certain foods, water and air pollution, x-rays, extended exposure to the sun and a variety of

other poisons we are exposed to in our every day environment.

Oxidation may also be linked with the effects of aging in humans, as well as with other conditions such

as cancer, hardening of the arteries, and rheumatoid arthritis. It appears that oxygen molecules and other

oxidizing agents, always hungry for electrons, extract these from the membranes in human cells. Over time, this

can cause a gradual breakdown in the body's immune system.

To overcome the effects of oxidation, some doctors and scientists recommend antioxidants-natural

reducing agents such as vitamin C and vitamin E. The vitamin C in lemon juice can be used to prevent

oxidizing on the cut surface of an apple, to keep it from turning brown. Perhaps, some experts maintain, natural

reducing agents can also slow the pace of oxidation in the human body.

In some of the normal chemical reactions that take place in the human body, strong oxidizing agents, such

as hydrogen peroxide, are formed. These highly reactive substances cause chemical changes in cell DNA that 50

can be damaging unless the changes are reversed. Fortunately, in healthy cells, normal repair reactions occur

that convert the altered DNA back to its normal form.

The repair mechanisms are thought to slow down with age. Some medical researchers believe that this

slowing down of DNA repair is connected to certain diseases associated with aging, such as cancer, heart

disease, cataracts, and brain dysfunction.

Substances called antioxidants that are found in food react with oxidizing agents (such as hydrogen

peroxide) and thus remove them from our system. This is believed to slow the alteration of DNA, so the slower

rate of normal repair can balance it.

Vitamins C and E are antioxidants, and foods that contain relatively high amounts of them are considered

important in slowing some of the medical problems that come from aging. Five servings of fruits and vegetables

per day are thought to supply enough antioxidants to provide reasonable protection from the damage done by

oxidizing agents.

2.10 Biogas production (anaerobic digestion)

Biogas typically refers to a gas produced by the biological breakdown of organic matter in the absence

of oxygen. Organic waste such as dead plant and animal material, animal dung, and kitchen waste can be

converted into a gaseous fuel called biogas. Biogas originates from biogenic material and is a type of biofuel.

51

Biogas is produced by the anaerobic digestion or fermentation of biodegradable materials such as

biomass, manure, sewage, municipal waste, green waste, plant material, and crops. Biogas comprises primarily

methane (CH4) and carbon dioxide (CO2) and may have small amounts of hydrogen sulphide (H2S), moisture

and siloxanes.

The gases methane, hydrogen, and carbon monoxide (CO) can be combusted or oxidized with oxygen.

This energy release allows biogas to be used as a fuel. Biogas can be used as a fuel in any country for any

heating purpose, such as cooking. It can also be used in anaerobic digesters where it is typically used in a gas

engine to convert the energy in the gas into electricity and heat. Biogas can be compressed, much like natural

gas, and used to power motor vehicles. In the UK, for example, biogas is estimated to have the potential to

replace around 17% of vehicle fuel.[] Biogas is a renewable fuel, so it qualifies for renewable energy subsidies

in some parts of the world. Biogas can also be cleaned and upgraded to natural gas standards when it becomes

biomethane.

Composition:

Typical composition of biogas

Compound Chem  %

Methane CH4 50–75

Carbon dioxide CO2 25–50

Nitrogen N2 0–10

Hydrogen H2 0–1

Hydrogen sulfide H2S 0–3

Oxygen O2 0–0

The composition of biogas varies depending upon the origin of the anaerobic digestion process. Landfill

gas typically has methane concentrations around 50%. Advanced waste treatment technologies can produce

biogas with 55–75% CH4, which for reactors with free liquids can be increased to 80-90% methane using in-situ

gas purification techniques As-produced, biogas also contains water vapor. The fractional volume of water

vapor is a function of biogas temperature; correction of measured gas volume for both water vapor content and

thermal expansion is easily done via a simple mathematic algorithm which yields the standardized volume of

dry biogas.

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In some cases, biogas contains siloxanes. These siloxanes are formed from the anaerobic decomposition

of materials commonly found in soaps and detergents. During combustion of biogas containing siloxanes,

silicon is released and can combine with free oxygen or various other elements in the combustion gas. Deposits

are formed containing mostly silica (SiO2) or silicates (SixOy) and can also contain calcium, sulfur, zinc,

phosphorus. Such white mineral deposits accumulate to a surface thickness of several millimeters and must be

removed by chemical or mechanical means.

Anaerobic digestion is the breakdown of organic material by micro-organisms in the absence of oxygen.

Although this takes place naturally within a landfill, the term normally describes an artificially accelerated

operation in closed vessels, resulting in a relatively stable solid residue. Biogas is generated during anaerobic

digestion (AD) - mostly methane and carbon dioxide - this gas can be used as a chemical feedstock or as a fuel.

Anaerobic digestion can treat many biodegradable wastes, including wastes that are unsuitable for composting,

such as meat and cooked food.

The process begins with separation of household waste into biodegradable and non-biodegradable waste.

The biodegradable material is shredded, slurried and then screened and pasteurised to start the process of

killing harmful pathogens. It is then pumped into the digester where bacteria break down the material and form

biogas, leaving a digestate. There are four main process stages in anaerobic digestion, as follows

Hydrolysis:

Hydrolysis is a chemical reaction in which the breakdown of water occurs to form

H+ cations and OH- anions. Hydrolysis is often used to break down larger polymers,

often in the presence of an acidic catalyst. In anaerobic digestion, hydrolysis is the

essential first step, as Biomass is normally comprised of very large organic polymers,

which are otherwise unusable. Through hydrolysis, these large polymers, namely

proteins, fats and carbohydrates, are broken down into smaller molecules such as amino

acids, fatty acids, and simple sugars. While some of the products of hydrolysis, including

hydrogen and acetate, may be used by methanogens later in the anaerobic digestion

process, the majority of the molecules, which are still relatively large, must be further

broken down in the process of acidogenesis so that they may be used to create

methane.

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Insoluble organic polymers such as carbohydrates, cellulose, proteins and fats are broken down and

liquefied by enzymes produced by hydrolytic bacteria. Carbohydrates, proteins and lipids are hydrolysed to

sugars which then decompose further to form carbon dioxide, hydrogen, ammonia and organic acids. Proteins

decompose to form ammonia, carboxylic acids and carbon dioxide. During this phase gas concentrations may

rise to levels of 80 per cent carbon dioxide and 20 per cent hydrogen.

Acidogenesis:

Acidogenesis is the next step of anaerobic digestion in which acidogenic microorganisms further break

down the Biomass products after hydrolysis. These fermentative bacteria produce an acidic environment in the

digestive tank while creating ammonia, H2, CO2, H2S, shorter volatile fatty acids, carbonic acids, alcohols, as

well as trace amounts of other byproducts Organic acids formed in the hydrolysis and fermentation stage are

converted by acetogenic micro-organisms to acetic acid. At the end of this stage carbon dioxide and hydrogen

concentrations begin to decrease. While acidogenic bacteria further breaks down the organic matter, it is still

too large and unusable for the ultimate goal of methane production, so the biomass must next undergo the

process of acetogenesis.

Acetogenesis:

In general, acetogenesis is the creation of acetate, a derivative of acetic acid, from carbon and energy

sources by acetogens. These microorganisms catabolize many of the products created in acidogenesis into acetic

acid, CO2 and H2. Acetogens break down the Biomass to a point to which Methanogens can utilize much of the

remaining material to create Methane as a Biofuel.

Methanogenesis:

Methanogenesis constitutes the final stage of anaerobic digestion in which, Methane (60%) and

carbon dioxide (40%) are produced from the organic acids and their derivatives produced in the acidogenic

phase. This step involves the reduction of organic acids to methane.The methane is a useful fuel source and

methanogenic bacteria play a further role in maintaining wider breakdown processes.

Efficient mixing of the contents of the digester improves the contact between the material and the

resident bacteria. Mixing of the waste slurry in the digester is important in maintaining a high rate of anaerobic 54

biodegradation and a high production level of gas. The mixing process disperses the incoming waste within the

digesting sludge, improving contact with the micro-organisms. Monitoring the acidity within the digester is

necessary to provide optimum conditions for the balanced growth of bacteria. Monitoring takes place in the

reactor using probes. The concentration of volatile fatty acids is an important parameter for monitoring as this

can be the first indicator that digestion is not progressing normally. Methanogens create methane from the final

products of acetogenesis as well as from some of the intermediate products from hydrolysis and acidogenesis.

There are two general pathways involving the use of acetic acid and carbon dioxide, the two main products of

the first three steps of anaerobic digestion, to create methane in methanogenesis

CO2 + 4 H2 → CH4 + 2H2O

CH3COOH → CH4 + CO2

While CO2 can be converted into methane and water through the reaction, the main mechanism to create

methane in methanogenesis is the path involving acetic acid. This path creates methane and CO 2, the two main

products of anaerobic digestion.

2.11 Weathering of rock

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Weathering is a process of the disintegration and degeneration of rocks minerals or soils as a result of

direct contact with the atmosphere of the Earth. When weathering takes place as a result of chemical reactions,

it is known as chemical weathering. In this process the rock disintegrates chemically as the chemicals in the

atmospheric agents react with the chemicals of the rock and the resultant reaction brings about the weathering

of the rock. The rate of weathering differs with variation in the chemical composition and structure.

Chemical reactions break down the bonds holding the rocks together, causing them to fall apart,

forming smaller and smaller pieces. Chemical weathering is much more common in locations where there is a

lot of water. This is because water is important to many of the chemical reactions that can take place. Warmer

temperatures are also more friendly to chemical weathering. The most common types of chemical weathering

are oxidation, hydrolysis and carbonation.Chemical weathering takes place in almost all types of rocks. Smaller

rocks are more susceptible, because they have a greater amount of surface area. Since the chemical reactions

occur largely on the surface of the rocks, therefore the smaller the fragments, the greater the surface area per

unit volume available for reaction.

The effectiveness of chemical weathering is closely related to the mineral composition of rocks. E.g.

quartz responds far slowly to the chemical attack than olivine or pyroxene.

Chemical Processes of weathering involving following reactions:

1.Hydration: Chemical combination of water molecules with a particular substance or mineral leading to a

change in structure.Soil forming minerals in rocks do not contain any water and they undergo hydration when

exposed to humid conditions. Up on hydration there is swelling and increase in volume of minerals. The

minerals lose their luster and become soft.

It is one of the most common processes in nature and works with secondary minerals, such as aluminium oxide

and iron oxide minerals and gypsum. For example,

2Fe2O3 + 3HOH        2Fe2O3 .3H2O

   (Hematite) (Red)            (Limonite) (Yellow)

Al2O3 + 3HOH           Al2O3 .3H2O

   (Bauxite)                      (Hyd. aluminium Oxide)

CaSO4 + 2H2O           CaSO4 .2H2O

  (Anhydrite)                    (Gypsum)

 

3(MgO.FeO.SiO2) + 2H2O  3MgO.2SiO2.2H2O + SiO2 + 3H2O

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(Olivine) (Serpentine)

 

2. Hydrolysis: Most important process in chemical weathering. It is due to the dissociation of H2O into H+ and

OH- ions which chemically combine with minerals and bring about changes, such as exchange, decomposition

of crystalline structure and formation of new compounds. Water acts as a weak acid on silicate minerals.

 

KAlSi3O8 + H2O              HAlSi3O8 + KOH

(Orthoclase)                       (Acid silt clay)

 

HAlSi3O8 + 8 HOH    Al2O3 .3H2O        +      6 H2SiO3

(Recombination)         (Hyd. Alum. oxide)        (Silicic acid)

This reaction is important because of two reasons.

a).clay, bases and Silicic acid - the substances formed in these reactions - are available to plants

b).water often containing CO2 (absorbed from atmosphere), reacts with the minerals directly to produce

insoluble clay minerals, positively charged metal ions (Ca++, Mg++, Na+, K+ ) and negatively charged ions

(OH-, HCO3-) and some soluble silica – all these ions are made available for plant growth.

3. Solution: Some substances present in the rocks are directly soluble in water. The soluble substances are

removed by the continuous action of water and the rock no longer remains solid and form holes, rills or rough

surface and ultimately falls into pieces or decomposes. The action is considerably increased when the water is

acidified by the dissolution of organic and inorganic acids. (e.g) halites, NaCl

NaCl + H2O -> Na+, Cl- , H2O (dissolved ions with water)

4. Carbonation: Carbon dioxide when dissolved in water it forms carbonic acid.

 

2H2O + CO2       H2CO3

This carbonic acid attacks many rocks and minerals and brings them into solution. The carbonated water has an

etching effect up on some rocks, especially lime stone. The removal of cement that holds sand particles together

leads to their disintegration.

CaCO3     +   H2CO3        Ca(HCO3)2

(Calcite) slightly soluble          (Ca-bicarbonate) readily soluble

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 5.Oxidation : The process where the metals present in the rock combine with oxygen and water to form oxides

like goethite, hematite and limonite is called oxidation of rock. These oxides make the rock weak and it

consequently crumbles to form smaller rock particles. This process can also be termed as rusting.

The absorption of oxygen is usually from O2 dissolved in soil water and that present in atmosphere.

The oxidation is more active in the presence of moisture and results in hydrated oxides. For example, minerals

containing Fe and Mg.

 

4FeO (Ferrous oxide) + O2       2Fe2O3 (Ferric oxide)

4Fe3O4 (Magnetite) + O2             6Fe2O3 (Hematite)

2Fe2O3 (Hematite) + 3H2O      2Fe2O3 .3H2O(Limonite)

6. Reduction: The process of removal of oxygen and is the reverse of oxidation and is equally important in

changing soil colour to grey, blue or green as ferric iron is converted to ferrous iron compounds which is called

as reduction. Under the conditions of excess water or water logged condition (less or no oxygen), reduction

takes place.

 

2Fe2O3(Hematite) - O2 4FeO(Ferrous oxide) - reduced form

 

During chemical weathering igneous and metamorphic rocks can be regarded as involving destruction of

primary minerals and the production of secondary minerals.

 

In sedimentary rocks, which is made up of primary and secondary minerals, weathering acts initially to destroy

any relatively weak bonding agents (FeO) and the particles are freed and can be individually subjected to

weathering.

2.12 Photo-oxidation

Many people who wear eye glasses prefer those made with photochromic lenses or glass lenses which

darken when exposed to bright light. These eyeglasses eliminate the need for sunglasses as they can reduce up

to 80% of transmitted light. The basis of this change in color in response to light can be explained in terms of

oxidation-reduction reactions. Glass consists of a complex matrix of silicates which is ordinarily transparent to

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visible light. In photochromic lenses, silver chloride (AgCl) and copper (I) chloride (CuCl) crystals are added

during the manufacturing of the glass while it is in the molten state and these crystals become uniformly

embedded in the glass as it solidifies. One characteristic of silver chloride is its suscepitibility to oxidation and

reduction by light as described below.

Cl- Cl + e-

oxidation

 

Ag+ + e- Ag

reduction

The chloride ions are oxidized to produce chlorine atoms and an electron. The electron is then transferred

to silver ions to produce silver atoms. These atoms cluster together and block the transmittance of light, causing

the lenses to darken. This process occurs almost instantaneously. As the degree of "darkening" is dependent on

the intensity of the light, these photochromic lenses are quite convenient and all but eliminate the need for an

extra pair of sunglasses.

The photochromic process would not be useful unless it were reversible. The presence of copper (I)

chloride reverses the darkening process in the following way. When the lenses are removed from light, the

following reactions occur:

Cl + Cu+ Cu+ + Cl-

oxidizing agent reducing agent oxidized species reduced species

The chlorine atoms formed by the exposure to light are reduced by the copper ions, preventing their

escape as gaseous atoms from the matrix. The copper (+1) ion is oxidized to produce copper (+2) ions, which

then reacts with the silver atoms as shown.

Cu+2 + Ag Cu+1 + Ag+

oxidizing agent reducing agent reduced species oxidized species

The net effect of these reactions is that the lenses become transparent again as the silver and chloride atoms are

converted to their original oxidized and reduced states.

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2.13 Photography

Early nineteenth century photographers produced crude images using papers impregnated with silver

nitrate or silver chloride. Their "photographs" darkened with time; a method to prevent the continued reaction of

light with the Ag-treated photographic papers had yet to be discovered. In 1839, however, Louis J. M Daguerre

invented light-fast images. His procedure relied on silver halide photochemistry, but included a process for

making the image permanent. Treatment of the exposed photographic plate (copper covered with a surface layer

of AgI) with mercury vapors, followed by washing with sodium hyposulfite (Na2S2O3), dissolved the silver

iodide from the unexposed portion of the plate. William Henry Fox Talbot improved process for coating silver

halides directly on paper in combination with a hyposulfite fixative replaced the daguerreotype by the end of the

nineteenth century. Although technologically more advanced, the basic procedures developed by Fox Talbot,

the "Inventor of Modern Photography," are used in all silver-based photography today. Modern silver-based

photography relies on oxidation-reduction chemistry to capture the image.

Chemical Reactions Involved in Photographic Processes

A. Silver-based photographic processes. Capturing light to produce an image utilizes two properties of the

silver cation: (1) Ag+ is reduced to silver metal in the presence of a halide which can be oxidized

photochemically (i.e., a photon ejects an electron from the halide). (2) Although the halide salts of silver, AgX,

have very low aqueous solubility, many complex ions of Ag+ (such as that formed with hyposulfite) do dissolve

in water. The media-specific solubility of silver halide salts make the initial image permanent. The key reactions

are outlined below:

1. Forming the image by exposure to light (hυ ) : A very small number of X - ions in the AgX crystals in the film

are oxidized to X. The electrons released from this oxidation reduce the Ag+ to silver metal in the surrounding

AgX crystal.

X - + hυ X + e-

Ag+ + e- Ag

B. Development:

The small number of Ag metal atoms formed (the latent image) act as a catalyst and sensitizes the

surrounding halide salt so that, in the presence of a developer i.e.a reducing agent, the sensitized AgX is

reduced, to produce black silver metal in the area exposed to light. Modern developers contain one of many

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reducing agents for this process. The most common is hydroquinone, which reacts with Ag+(in AgX), as shown

in equation below:

Note that the above reaction occurs in basic medium (OH-). The development can be stopped, therefore, by

dipping the photographic film in acid. The most common "stopper" contains acetic acid. In this experiment, we

will remove excess reagents by washing the exposed "film" in water before fixing the image.

1. "Fixing" the image (making it permanent): Unexposed AgX on the photography film (plate or paper) is

removed by complex formation with thiosulfate.

2. The soluble complex, [Ag(S2O3)2-]23-, can be readily washed away to leave only the dark

silver metal image.

AgX + 2 S2O32- [Ag(S2O3

2- ]23-

The process described above forms the negative in conventional black-and-white photography; light shining

through the negative produces the final photograph (the positive) using this same chemistry.

2. Toning (coloring) the image. The silver-based black and white photographs may be altered by toning, using

chemistry to produce different colored images. For example,reactions of Ag with thiosulfate in acid solution

produces sulfur that then reacts with the Ag-image to yield the brown Ag2S of sepia photographs.

B. Cyanotypes. The blue photographs on formation of insoluble Prussian Blue through photoreduction of

Fe(III) to Fe(II) are called cyanotypes. Toning reactions may alter the color of the initial image. Photoreaction

and formation of insoluble Prussian Blue. The two photoactive iron compounds most commonly used for

photography are ferric ammonium citrate, Fe(III)NH4(C6H6O72-)2, and ferric ammonium oxalate, Fe(III)

(NH4(C2O42- )2.

Citrate and oxalate are the anions generated by loss of two acidic protons from citric and oxalic acids,

respectively. Interaction of light with these anions leads to their oxidation and releases carbon dioxide and an

electron (equation 4a) which then reduces Fe(III) to Fe(II) (4b). Although we will be using ferric ammonium

citrate in this lab, the redox chemistry is more readily apparent in the reaction of oxalate as shown in equation

4a:

61

C2O42- + hυ 2 CO2 + 2 e- (4a)

Fe3+ + 1 e- Fe2+ (4b)

Overall reaction

2 Fe3+ + C2O42- + hυ 2 Fe2+ + 2 CO2 (4c)

The Fe(II) formed in (4c) combines with CN- present in the solution to form the complex [Fe(CN)6 ]4- which, in

turn, gives the insoluble blue Prussian blue, Fe(III)4[Fe(CN)6]3, adhering to the fibers of the cloth or paper on

which the reagents had been coated:

Fe2+ + 6 CN- [Fe(CN)6]4-

3 [Fe(CN)6]4- + 4 Fe3+ Fe(III)4[Fe(CN)6]3

The blue color results from interaction between the iron in two different oxidation states. (Similar

compounds, all containing both Fe(III) and [Fe(CN)6]4-, but with either K+ or NH4+ ions also are blue or

greenish-blue in color and are also called Prussian Blue as well. Only compounds containing iron in these two

oxidation states of iron are blue).

Conclusion:

Oxidation and reduction are the unique type of reactions which involves the transfer of electron or the

transfer of atom i.e. hydrogen or oxygen, from one species to another. Both oxidation and reduction reactions

are dependent on one another, which means that we cannot expect oxidation reaction without the reduction 62

process or vice versa.So that there must be presence of both oxidizing and reducing species for the redox

reaction to take place.

Oxidation and reduction reaction are having wide range of applications due to the fact that most of them

are energy yielding reactions. For example, combustion ,redox reactions in electrochemical cell and metabolic

processes etc.Redox reactions are very essential to sustain life on earth. As they play a major in respiration,

photosynthesis and metabolic processes one cannot expect life without the redox reactions.

One can observe that, most of the redox reactions carried out by the nature with ease and less expense of

energy, which intern leads to generation of large amount of energy. On contrary to this it is found to be very

difficult to carry out those redox reactions in laboratories under normal conditions and require extreme

conditions.For instance, the nitrogen fixation occur in nature, under normal temperature and pressure. But in

order to carry out the same process in laboratory, it require to maintain 500-550 ◦C and above 300 atmospheric

pressure.

All the redox reaction may not be useful,i.e.some of them may leads to undesired products as in case of

corrosion, in which the metals get destructed due to the oxidation of the same.

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5. Atkins, P. W., and J. A. Beran. General Chemistry. 2nd ed. New York: Scientific American Books,

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