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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
1
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
2
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)
4
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
8
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
9
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
10
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
11
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
12
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.
13
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.
14
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.
15
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 .
16
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
17
(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
18
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
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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.
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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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