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Redox Reactions

How fireworks work? How fireworks work? Fireworks have been a traditional part of festive celebrations since they were invented by Chinese thousands of years ago. The chemical reactions that take place in firecrackers are REDOX reactions. Firecrackers consist of gunpowder, oxidants (potassium nitrate, potassium chlorate), reductants (charcoal, sulfur) and metals such as Cu, Na, Ba, Sr, Ti etc. to get colored sparkles: typically blue, golden yellow, green, red and white. The reaction is accompanied by evolution of heat, sound and light.

Learning Objectives

After completing the topic, the student will be able to:

  • Define reduction and oxidation (redox in short)reactions.
  • Represent a redox reaction in terms of electron transfer.
  • Define oxidizing and reducing agents and give examples.
  • Identify the oxidizing or reducing agent in a given reaction.
  • Define oxidation number and calculate the change in valency or oxidation number of species in a redox reaction.
  • Balance a redox reaction using ion–electron method.
  • Discuss everyday examples involving a redox process.
Simple understanding of oxidation and reduction reactions Simple understanding of oxidation and reduction reactions Oxidation refers to the loss of electrons, while reduction refers to the gain of electrons.
Oxidation and Reduction

In chemical reactions, inter molecular bonds between reactants are broken and new bonds are made in the product molecules. This transformation is the simple rearrangement of electrons. Reactions that involve the transfer of one or more electrons from one reactant to another are called oxidation–reduction reactions.

Oxidation is the process in which electrons are removed from an atom or ion.
Reduction involves the addition of electrons to an atom or ion.

Alternately, oxidation reaction would be an addition of oxygen to a substance.
Reduction reaction would mean removal of oxygen from a substance.

Burning of coal is an oxidation process Burning of coal and magnesium Burning of coal (a) and magnesium (b) are rapid oxidation processes while rusting of iron is a slow oxidation process.
Examples for oxidation process

Since years, we have been burning coal to get heat. We also know that for burning of coal we need supply of fresh air. What happens when coal burns? The oxygen from the air reacts with carbon in coal. We can represent this reaction by the chemical equation as follows:

C + O2 CO2 (g)

Carbon dioxide is the product. If sufficient supply of fresh air is not available, the reaction may be as follows:

2C + O2 2CO (g)

Carbon monoxide is the product. Both the products, CO2 and CO, being gases escape into the atmosphere.

Another common experience is burning of magnesium wire, which is popular in the fireworks.

2Mg + O2 2MgO
Burning match stick glows brighter confirms the release of oxygen Decomposition of red mercuric oxide (right) Mercuric oxide on heating at a temperature of 400 °C turns black and releases oxygen into the balloon.
Examples for reduction process

A common example of reduction is the burning of wood. Wood is mainly cellulose, a substance made up of carbon and hydrogen (C6H10O5)n. As wood burns, its carbon–hydrogen bonds break and the carbon and hydrogen atoms form new bonds with oxygen molecules in the air to create carbon dioxide and water. The formation of these products involves the transfer of electrons from one atom to another and so is by definition an oxidation–reduction reaction. What remains behind is charcoal – which is mostly carbon, C.

C6H12O6 + 6O2 6CO2 + 6H2O

Mercuric oxide on heating reduces to mercury i.e, oxygen is removed.

2HgO 2Hg + O2

If you keep a metal oxide in a test tube and heat it, then bring a burning matchstick near the mouth of the test tube, you will find the matchstick glows brighter. This indicates that the gas escaping from the test tube is oxygen (refer figure at left).

Action of hydrogen on copper oxide Action of hydrogen on copper oxide A hot piece of copper in air forms a black oxide coating which is reduced to copper in hydrogen. The reaction is reversible. A distinct line is observed at the interface between air and hydrogen (click figure for animation).
Representation of redox reaction
Representation of redox reaction In this reaction, we see that oxidation and reduction reaction are occurring together.
Oxidation-reductions reactions

Certain reactions where in hydrogen is removed, are also grouped as oxidation reactions. Similarly, if hydrogen is added, the reactions are called reduction reactions.

Br2 + H2S S + 2HBr

In this reaction H2S is oxidized to S, because hydrogen is removed. Simultaneously, bromine is reduced to HBr because hydrogen is added. In this reaction, the oxidation–reduction reactions occur simultaneously on different reactants. These are called redox reactions.

Redox reactions

Thus, oxidation may be defined as a chemical reaction in which oxygen is added or hydrogen is removed and reduction is the chemical reaction in which hydrogen is added or oxygen is removed.

Another example:

When copper oxide is heated with hydrogen, copper metal and water are formed.

CuO + H2 Cu + H2O

In this reaction, hydrogen (H2 ) changes to water (H2O) i.e, oxygen is being added to hydrogen. By simple definition of oxidation, we can say that hydrogen is being oxidized to water. Further, CuO is changing to Cu i.e, oxygen is being removed from the copper oxide. By simple definition of reduction, we can say that copper oxide is being reduced to copper.

Oxidation is loss of electrons and Reduction is gain of electrons Oxidation is loss of electrons and Reduction is gain of electrons
Oxidation-Reduction Process

Oxidation and reduction can also be defined in terms of loss or gain of electrons.

Oxidation is the process whereby a reactant loses one or more electrons. Reduction is the opposite process whereby a reactant gains one or more electrons. Oxidation and reduction are complementary and simultaneous processes. They always occur together, you cannot have one without the other. The electrons lost by one chemical in an oxidation reaction don't simply disappear, they are gained by another chemical in a reduction reaction.

Oxidation and reduction of  Chlorine and Sodium respectively Oxidation and reduction of Chlorine and Sodium respectively The Na starts 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.
The Cl2 also starts with an oxidation number of zero (0). But it ends up with an oxidation number of 1−. It has been reduced from chlorine atoms to negative chloride ions.

Oxidation-reduction in NaCl
2Na + Cl2 2NaCl

To see how electrons are transferred in this reaction, we shall look at each reactant individually. Each electrically neutral sodium atom loses an electron and changes to a positively charged ion. It is therefore oxidized.

2Na 2Na+ + 2e    Oxidation

Each electrically neutral chlorine molecule changes to two negatively charged chloride ions by gaining electrons. It is therefore reduced.

Cl2 + 2e 2Cl    Reduction

The net result is that the two electrons lost by the sodium atoms are transferred to the chlorine atoms. Each equation shown above is called a half-reaction and together represent the whole redox (oxidation–reduction) process.

Attaining stable inert gas configuration Attaining stable inert gas configuration Sodium losing its 2s electron to chlorine. Chlorine, short of electron completes its 3p shell by accepting this electron .

Half-reactions are useful for showing which reactant loses electrons and which reactant gains them, which is why half-reactions are used throughout this chapter. Because the sodium causes reduction of chlorine, the sodium is acting as a reducing agent. Note that in behaving as a reducing agent, the sodium is oxidized – it loses electrons.

Conversely, chlorine causes oxidation of sodium and so it is acting as an oxidizing agent. Again, note that an oxidizing agent itself reduces.

Sodium gets inert gas configuration of Neon (1s22s22p6) by losing one electron.
Chlorine gest inert gas configuration of Argon (1s22s22p63s23p6) by gaining one electron.

Mnemonics: Just remember that loss of electrons is oxidation, and gain of electrons is reduction. Here is a mnemonic adopted from a popular children's story: Leo the lion went “ger”. Another mnemonic is OIL RIG: Oxidation Is Loss-of-electrons and Reduction Is Gain-of-electrons.

Oxidation and reduction tendencies Oxidation and reduction tendencies The ability of an atom to gain or lose electrons is a function of its position in the periodic table. Those at the upper right tend to gain electrons, and those at the lower left tend to lose them.
Oxidation and reduction tendencies

Different elements have different oxidation and reduction tendencies – some lose electrons more readily, while others gain electrons more readily, as the figure illustrates.

The tendency to do one or the other is a function of how strongly the atom's nucleus holds electrons. Greater the effective nuclear charge, greater the tendency of the atom to gain electrons. Because the atoms of elements at the upper right of the periodic table have the strongest effective nuclear charges (with the noble gases excluded), these atoms have the greatest tendency to gain electrons and hence behave as oxidizing agents.

The atoms of elements at the lower left of the periodic table have the weakest effective nuclear charges and therefore the greatest tendency to lose electrons and behave as reducing agents.

Ammonia molecule Ammonia molecule Nitrogen combines with 3 hydrogen atoms. So valency of nitrogen is three and valency of each hydrogen atom is one. The bar on nitrogen represents unpaired free electrons (lone pair).
Change of valency

Valency is the combining capacity of an element and depends on the availability of free electrons, especially in outer orbital of the atom. Certain elements like Fe, Sn with incomplete sub–orbitals, change their valency during the reaction. These can be understood in terms of electron transfer and hence oxidation-reduction reactions.

Decrease in valency is termed as reduction. In the below reaction, Fe valency changes from + 3 to + 2.

2FeCl3 + H2S 2FeCl2 + 2HCl + S

Increase in valency is described as oxidation. In the below reaction, Sn valency changes from + 2 to + 4.

SnCl2 + Cl2    SnCl4

Valency of Magnesium Valency of magnesium and oxygen atoms The magnesium atom attains +2 valency and oxygen atom attains −2 valency while participating in a chemical reaction.

Mg atom (Z = 12) has 2 electrons in the K–shell, 8 electrons in the L–shell and M–shell has 2 electrons. By loss of two electrons Mg is converted to (Mg++) i.e. magnesium ion. It acquires two positive charges by this reaction.

2Mg 2Mg++ +  4e

An oxygen atom, close by, first gets converted into an oxide ion (O−−), by borrowing two negative charges from the magnesium ion as

O2 + 4e 2O2−

So the net reaction may be represented as

2Mg + O2 2(Mg++O−−) 2MgO

Thus, in the formation of magnesium oxide, magnesium atom is oxidized by the loss of two electrons.

Valence electrons in redox reactions

In the formation of KCl, potassium losses it's 2s valence electron to chlorine. Thus the chemical reaction between the potassium and chlorine is a result of transfer of valence electrons.

2K 2K+ + 2e

Cl2 + 2e 2Cl

Net reaction is:  2K + Cl2 2(K+Cl)

The electrons lost by potassium are gained by chlorine. Hence it is a redox reaction.

When nascent hydrogen (i.e. Hydrogen in atomic form H, not as molecule H2), is passed through ferric chloride (FeCl3) solution, it is reduced to ferrous chloride.

H (atom) H+ + e

Fe3+ + 3Cl + e Fe2+ + 3Cl

(Ions, which do not take part in a reaction, are known as spectator ions and disregarded in the net reaction. Here Cl ions are disregarded)

Net reaction is:  Fe+++ + H      Fe++ + H+

In the above reaction, nascent Hydrogen is said to be oxidized (H becomes H+ ion by loss of one electron). Fe3+ is reduced to Fe2+ by gain of one electron.

When one metal displaces another one from its salt, the metal is said to be oxidized and the one, which gets displaced is said to be reduced.

CuSO4 + Fe FeSO4 + Cu

Here Fe is oxidized, Cu is reduced and SO42− ion is a spectator ion.

Silver is reduced and Copper is oxidized Silver is reduced and copper is oxidized A typical reaction with silver nitrate solution is to suspend a rod of copper in it and leave for a few hours. The silver nitrate reacts with copper to form hairlike crystals of silver metal and a blue solution of copper nitrate.
Redox reaction
Redox reaction in storage cell battery
Oxidizing and reducing agents

A substance which brings out oxidation of another substance, is called as an oxidizing agent or an oxidant. Oxidizing agent is also defined as a substance, which forces another substance to lose electrons and, in turn, it gains the electrons. For example: O2, Cl2. Non–metals are borrowers of electrons and can therefore act as oxidants.

Substances, which force another substance to gain electrons and itself lose them, are called reducing agents or reductants. All metals and hydrogen give up electrons and form +ve ions.

Hydrogen Loses

There also compounds which acts either as an oxidizing or a reducing agent. Compounds like hydrogen, hydrogen peroxide are examples of this kind. Hydrogen acts as oxidizing agent when it reacts with metals (as in eqn (i) ) and as reducing agent when it reacts with non-metals (as in eqn (ii)).

2 Na (S) + H2 (g) 2 NaH (g)    -----(i)

H2 (g) + Cl2 (g) 2 HCl (g)    -----(ii)

CuO + H2 Cu + H2O    -----(iii)

In eqn (iii), CuO is an oxidizing agent and H2 is a reducing agent.

H2S + Cl2 S + 2HCl    -----(iv)

In eqn (iv), H2S is a reducing agent and Cl2 is an oxidizing agent.

Oxidizing and reducing agents are routinely used for extraction of metals and non–metals from ores. Also these reactions are seen in all storage cells or batteries.

Permanganate ion- A versatile oxidizing agent Permanganate ion – A versatile oxidizing agent Permanganate solutions are purple in color and are stable in neutral or slightly alkaline media. The exact chemical reaction is dependent upon the organic contaminants present and the oxidant utilized.
Permanganate ion

Permanganate ion (MnO4) in permanganate solution reacts differently in acidic, basic and neutral solutions. This is because of the electrons which get transferred during the reaction. The oxidation reactions are written as a half–reaction. The permanganate ion is a very versatile oxidizing agent.

The color of the permanganate solutions indicates whether the reaction involves one electron transfer or more than one–electron transfer. If the color of the solution is deep violet, then the permanganate ion is said to be in its ionic state (MnO4). If the color turns green, then the permanganate ion changes to another ionic state (MnO2), mostly observed in basic solutions. If the color changes to colorless, then the permanganate ion is no more in its ionic form. The Mn2+ ions catalyze the reaction, mostly observed in acidic solutions.

In neutral or slightly acidic solutions, the ionic form of permanganate changes to non–ionic form of permanganate (MnO2) product, which is in the form of a black precipitate.

The reactions of permanganate ion, in acidic, basic and neutral solutions The reactions of permanganate ion, (MnO4) in acidic, basic and neutral solutions
Reactions of permanganate ion

The following are the reactions of permanganate ion, (MnO4) in acidic, basic and neutral solutions. In acid solutions, the half–reaction of permanganate ion is:

MnO4 + 8H+ + 5e Mn2+ + 4H2O

Five electrons are used in this reaction, and the soluble Mn2+ ion is colorless. This reaction is slow and the Mn2+ ion catalyzes the reaction.

In basic solutions, the half‐reaction is:

MnO4 + 4H+ + 3e MnO2 + 2H2O

The reaction mixture changes color from the deep violet of permanganate ion, (MnO4) to green color for the MnO2 ion.

In neutral solutions, the half–reaction is:

MnO4 + e MnO42−

Neutral solutions mainly involve water, which dissociate into H+ ions and react with three electrons of permanganate ion to give an insoluble black precipitate of MnO2.

Oxidation number

In a complex reaction, a chemical charge alone is not enough to tell whether a substance is undergoing oxidation or reduction. For example:

Oxidation Number

In such situations, oxidation number is considered important to understand if oxidation or reduction has taken place.

Definition: The oxidation number of an element represents the numerical value and sign of charge on the atom of an element, when it is in the form of an ion. Thus, the oxidation state or number means the same thing as the charge on the ion. The oxidation state of a free (i.e. un‐combined) element is zero, as it has no charge on it. Oxidation numbers of elements in some ionic compounds are listed below.

Oxidation Numbers of Elements

In a covalent compound, there is no actual loss or gain of electrons between atoms. So, strictly speaking the concept of oxidation number cannot be applied to covalent compounds. But as a matter of convenience, the concept is applied to covalent compounds with certain restrictions. Since the shared electron pair is attracted towards one or the other atom involved, this is justified.

Determination of oxidation number Sulfur's oxidation number = +6 Common oxidation states of sulfur range from −2 to +6 . Oxidation state shows the total number of electrons which have been removed from an element (a positive oxidation state) or added to an element (a negative oxidation state) to get to its present state.
Electronegativity and oxidation number

Over the years, the definition of oxidation–reduction has been broadened to include processes, which involve combinations of atoms in which there is no clear–cut transfer of electrons between them. An understanding of this behavior is provided by the concept of electronegativity. According to this concept, each kind of atom has a certain attraction for the electrons involved in a chemical bond. This “electron–attracting” power of each atom can be listed numerically on an electronegativity scale. Fluorine, which has the greatest attraction for electrons in bond–forming situations, is assigned the highest value on this scale. More electronegative element will attract the shared electrons more, hence it has −ve oxidation number where as the other atom has + ve oxidation number.

The electronegative elements (non–metals) may be arranged in decreasing order of their electronegativity as follows:

F > O > Cl > N > Br > S > C> I > P > H

Thus the Oxidation number of different elements in a covalent compound can be shown as:

Oxidation Numbers of Covalent Compounds
Covalent Compounds
Sulfur's oxidation number = +4 Sulfur's oxidation number = +4

Oxidation number of sulfur depends on the compound it is in. For example: In H2SO4 oxidation number is +6, in H2SO3 oxidation is +4 and in H2S2O3 oxidation is +2.

Oxidation number of covalent compounds

In covalent compounds, elements having higher electronegativity have a −ve oxidation number while elements having a lower electronegativity have a +ve oxidation number.

Electro Negativity

Thus, to include the cases of covalent compounds too, the definition of the oxidation number may be modified. Oxidation number of an element is defined as the numerical value and sign of the charge, which the atom of the element appears to have in a molecule of the compound.

How to calculate oxidation number of nitrogen in NO3ion:

Oxidation Number of Nitrogen

So, X = (6) − 1 = +5

Determining the oxidation number of carbon Determining the oxidation number of carbon in C2H5OH

C2H5OH = 2(x) + 5(1) + (−2) + (1)
     0 =2x + 5 − 2 + 1
     x = −2
∴ The oxidation state of carbon in ethane is: − 2.

Variable oxidation numbers

Some elements are capable of having several oxidation numbers i.e., in different compounds the oxidation number is different. Examples:

1. Sulfur:   Sulphur

2. Nitrogen:   Nitrogen

The increase in oxidation number of an atom or ion of an element is oxidation and the decrease in oxidation number is reduction. Examples:

Oxidation Number of Mg is Increased

In above reaction, the oxidation number of Mg is increased from zero to two. So Mg is oxidized.
Oxygen is reduced because the oxidation number of oxygen has been lowered from zero to minus two. Thus Mg is a reducing agent and oxygen is an oxidizing agent.

Oxidation Number

In above reaction, CO2 is reduced and hydrogen is oxidized. Thus CO2 is an oxidizing agent and H2 is reducing agent.

Structural elucidation of chemical bonds Structural elucidation of chemical bonds Fractional oxidation value does not mean that the number of electrons shared are in fraction. It is considered only to satisfy the possible structure of compound with existing chemical bonds
Fractional oxidation value

Certain compounds like Fe3O4, C3O2, N3H, etc., exhibits a fractional value of oxidation state. This is due to a state exhibited by the element to satisfy the structural elucidation of chemical bonding The stability of such compounds also need to be examined.

Some more species with fractional oxidation value are: S4O6-2, Mn3O4, Pb3O4, Br3O8 etc.

Calculating the oxidation state

Oxidation state of iron in its oxide: Fe3O4:
Consider the oxidation state of Fe as 'x' and -2 state for oxygen.
Then, Fe3O4 = 3x + 4(-2)
             0 = 3x - 8
             x = 8/3

Methods of balancing chemical equations Methods of balancing chemical equations A chemical equation is a theoretical or written representation of what happens during a chemical reaction. In a chemical reaction, the number of atoms that are present in the reactants has to balance the number of atoms that are present in the products.
Balancing of reactions

The following rules provide another method of balancing a chemical reaction:

  • Write the skeletal equation representing the actual chemical change.
  • Write the oxidation number of each element in the reactants and the products below the corresponding symbol.
  • Pick out the elements undergoing change in oxidation number.
  • Write the skeletal equation once again and indicate the oxidation numbers of only those elements, which undergo change in oxidation numbers.
  • Find the total oxidation number changes for each element and write these on lines connecting the species.
  • Invert these numbers for the species involved so that the increase in oxidation number becomes equal to the total decrease in oxidation number.
  • Finally balance the other species by inspection. For reaction in acidic solution, add H+ or H2O or both to the equation as needed. For reactions in basic solution, add OH or H2O or both to the equation as needed. For ionic equations, charges must be balanced on both sides of the equation.
Ion - Electron Method

In order to balance the complex oxidation–reduction reactions within a length of time, chemists have developed a simple and most versatile method known as Ion–electron method. The reactions which occur in acid solutions (containing H+ ions), can be balanced by using Ion–electron method, which involves the following steps explained for an acidic solution.

Standard solutions of iodine are prepared by reacting iodide ions with the iodate ion in acid solution. The skeleton reaction is

I + IO3 I2

Step 1:The 2 reactants giving the same product are separated bytheir half-reactions.

I I2
IO3 I2

Step 2:The atoms on the reactant side are balanced by using the coefficient x, which is equal to the number of atoms in the product side (right side) in both half–reactions.

2I I2
2IO3 I2

Step 3:Water molecules (H2O) are added to any one or both half–reactions on the reactant side of the reaction, which contain oxygen, in order to balance the oxygen atoms on the product side of the reaction. This step applies only to the second half–reaction since the first one contains no oxygen. Six water molecules are needed to balance the six oxygens.

2I I2
2IO3 I2 + 6H2O
(Contd ...)

Step 4:Hydrogen atoms (H+ ions) on the product side, in one or both half-reactions are balanced by adding the number of H+ ions on the reactant side. This step also applies only to the second half-reaction, where 12 H+ ions are required.

2I I2
12H+ + 2IO3 I2 + 6H2O

Step 5:All the ions on both half–reactions are balanced, by adding the number of electrons (e) either on the product side or reactant side, such that the total charge is zero on both sides of two half–reactions.

2I I2 + 2e
10e + 12H+ + 2IO3 I2 + 6H2O

Step 6:The electrons are equalized in the two half–reactions, by multiplying with a suitable number (5 in this case) such that both half–reactions have same number of electrons.

10I 5I2 + 10e
10e + 12H+ + 2IO3 I2 + 6H2O

Now, the reactants and products of the two half–reactions are added resulting in an equation.

10I + 10e + 12H+ + 2IO3 5I2 + 10e + I2 + 6H2O

Cancel the number of electrons from each side. Divide by a common factor (2 in this case) to obtain the final balanced equation.

5I + 6H+ + IO3 3I2 + 3H2O

This completes balancing by Ion-electron method.

The reactions which occur in basic or neutral solutions (containing OH ions), can be balanced by the same Ion–electron method, which involves an additional step in order to balance the OH ions.

Step 7: Add one OH ion for each H+ ion to both sides of the equation in step 6. Combine the H+ and OH ions on one side of the reaction into water molecules. Cancel H2O molecules that appear on both sides of equation.

Check: All the atoms must be balanced in the reaction. In addition, in ionic reactions the charges must also balance.

Steel hulls in ships have zinc blocks Steel hulls in ship - Oxidation reaction Steel hulls in ships have zinc blocks attached because they oxidize and release electrons. There electrons get consumed by the steel and prevent corrosion
Everyday examples

Oxidation–reduction reactions have many applications in our lives. Some of these applications are so common, that we take them for granted. Few examples:

  • Photography
  • Photosynthesis
  • Metabolism
  • Combustion
  • Bleaching agents
  • Nitrogen fixation
  • Dry Cell
  • Photosensitive goggles

The above applications are explained in some detail in the next sections.

Camera Redox reactions The light energy causes electrons to be ejected from a few of the bromide ions, oxidizing them to elemental bromine.
Photo Negative
Developed film is a positive film The negative is dark where Ag+ions have been reduced to metallic silver. Light projected through the negative is captured on photographic paper as a positive image.

Lay some wax paper on the back of an open unloaded camera. Hold the shutter open and then focus. You have an image. Let the shutter close, however, and the image is gone. This is the same image that forms on the photographic film inside a loaded camera. The difference between the film and the wax paper is that the film is able to retain the image after the shutter has closed. How does it do that? The answer has to do with oxidation–reduction chemistry.

A simplified explanation of how a black and white photograph is produced is given below:

1. Un‐exposed black and white photographic film is a transparent strip of plastic, coated with a gel containing micro crystals of silver bromide, AgBr. Light reflected from the subject being photographed passes through the camera lens and is focused on these micro crystals. The light causes many of the bromide ions in the micro crystals to oxidize. The electrons set loose by this oxidation are transferred to the silver ions, which are thereby reduced to opaque silver atoms. The more light received by a microcrystal, the greater the number of opaque silver atoms formed. In this way, the photographic image is encoded, and the film is said to be exposed.

2. The light reflected from the subject does not typically result in the formation of enough silver atoms to make a visible image. The more silver atoms a microcrystal contains, however, the more susceptible it is to further oxidation–reduction reactions. To make a visible image, the photographer puts the film in a light–tight container to prevent further exposure. Then the film is treated with a reducing agent, such as hydroquinone, C6H6O2, which reveals the encoded image by causing the formation of many more opaque silver atoms. Through this step the image develops.

3. The reduction of silver ions by the hydroquinone developing solution is stopped by treating the film with a solution of sodium thiosulfate, Na2S2O3 (also called either hypo or fixing solution). The thiosulfate ion, S2O32–, binds with any un‐reduced silver ions to form a water–soluble salt. Subsequent washing with water removes everything except the silver atoms adhering to the film, which are most abundant where the greatest amount of light hit the film when the photograph was taken. The film is now fixed.

4. Because the silver atoms are opaque, the film appears as a negative, which is dark where the subject was light and light where the subject was dark.

5. Light is projected through the negative onto photographic paper, which is developed using the same reactions that produced the negative. The resulting developed image is a negative of the negative – in other words, a positive print.

Color photographic film is coated with a variety of chemicals that respond to light of different frequencies (colors). There are more oxidation–reduction reactions involved in the developing a color photograph, but the basic principle is the same, the selective reduction of only those chemicals exposed to light. Digital photography, by contrast, is an outgrowth of photo–voltaic cells, which are made of metalloids, such as silicon, that lose electrons upon exposure to light.

Photosynthesis Photosynthesis 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.

An example of naturally occurring biological oxidation–reduction reaction is the process of photosynthesis. It is a very complex process carried out by green plants and blue–green algae. These organisms are able to harness the energy contained in sunlight, and via a series of oxidation–reduction reactions, produce oxygen and sugar. This may be utilized for energy as well as the synthesis of other compounds.

The overall equation for the photosynthetic process may be expressed as:

In Presence Of Glucose

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.

Electron transport in photosynthesis Electron transport in photosynthesis is a sequence of redox reactions Electron transport chains are the cellular mechanisms used for extracting energy from sunlight in photosynthesis and also from redox reactions, such as the oxidation of sugars (respiration). The electron transport chain in the mitochondrion is the site of oxidative phosphorylation in eukaryotes. The NADH and succinate generated in the citric acid cycle are oxidized, providing energy to power ATP synthase.
Photosynthesis mechanism

Photosynthesis is the net result of two processes: Light reaction and dark reaction.

The first process involving the splitting of water is really an oxidative process that requires light, and is often referred to as the “light reaction”. This reaction may be written as:

Light Reaction

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:

Nicotinamide Adenine Dinucleotide Phosphate

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.


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 + 24H+ + 24e C6H12O6 + 6H2O

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”.

Glucose metabolism Glucose metabolism involves many redox reactions Glucose is a much stronger reducing agent than water Therefore, the transfer of electrons from glucose to O2 is a thermodynamically downhill, energy–releasing process. Cells use the energy released by the glucose/oxygen redox process to carry out a wide variety of energy–requiring activities.

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: anabolism (reactions involving the synthesis of compounds) and catabolism (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 a process can be described by the oxidation of glucose as shown below. Note that this equation is the reverse of the photosynthetic equation.

C6H12O6 + 6O2 6CO2 + 6H2O + Energy

Note also that in this reaction, the carbon atoms in glucose are oxidized, undergoing an increase in oxidation state (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 Triphosfate (ATP) and Nicotinamide Adenine Dinucleotide (NADH).

Redox reactions in combustion of methane The combustion of methane to carbon dioxide is an oxidation of carbon because the oxidation number of carbon increases from −4 to +4.

The combustion or the burning of fuels, is perhaps the most common and obvious example of oxidation and reduction. Combustion is also that process which converts the potential energy of fuels into kinetic energy (heat and light). Most fuels (gasoline, diesel oil, propane, etc.) are compounds comprised primarily of carbon and hydrogen. These hydrocarbons represent an excellent source of potential energy, which is released as heat during the combustion process.

A common example is the oxidation of propane, the fuel used for gas ranges:

C3H8 + 5O2 4H2O + 3CO2 + 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 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.

Combustion \96 Example of redox reaction Combustion – Example of redox reaction Combustion reactions are good examples of redox reactions where one molecule gains oxygen (is oxidized) and one molecule gains hydrogen (is reduced).
Alcohols and Hydrogen as fuels

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 (H2), which readily reacts with oxygen forming water as shown:

2H2 + O2 2H2O + Energy

The simplicity and “non polluting” 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.

Bleaching Bleaching action A common household bleach is a solution of sodium hypochlorite which is an oxidizing agent. The compound being bleached (e.g. stains on clothing etc) are the reducing agents. Other bleaches such as hydrogen peroxide (use for hair) and sodium percarbonate (Napisan) are also both oxidizing agents. Sodium thionite used for fabric dye bleaching is a reducing agent.
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 of the commercial bleaches are oxidizing agents, such as sodium hypochlorite (NaOCl) or hydrogen peroxide (H2O2), which are quite effective in “ decolorizing ” substances via oxidation. The action of these bleaches can be illustrated in the following simplified way:

The decolorizing action of bleaches is due to their ability to remove the 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.

OCl + 2e + HOH Cl + 2OH

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.

Pea plant Green peas are excellent soil builders - plants that have the ability to absorb nitrogen from the surrounding atmosphere and attach it in the form of nodules to its roots; also known as "nitrogen fixing". As soon as the green peas are harvested and pea plants are tilled under, other nitrogen loving companion plants such as beans will benefit from all that nitrogen that has been gathered up in the soil.
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 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. Nitrogen–fixing bacteria present in the soil carry out this complex process. Although nitrogen–fixation involves a number of oxidation–reduction reactions that occur sequentially, the reaction which describes its reduction can be written in a simplified way as:

N2 + 6e + 8H+ 2NH4+ (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.

Root nodules Nitrogen fixation process Many Fabaceae host bacteria in their roots within structures called root nodules. These bacteria, known as rhizobia, have the ability to take nitrogen gas (N2) out of the air and convert it to a form of nitrogen that is usable to the host plant (NO3 or NH3). The legume, acting as a host, and rhizobia, acting as a provider of usable nitrate, form a symbiotic relationship.

Ammonia is formed 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:

Reduction Reactions Involving Nitrate and Nitrite ions
Dry Cell Principle of redox reactions in dry cell batteries 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).
Dry cell

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 wristwatches 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.

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

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

Zn(s) Zn+2 + 2e

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 + 2OH ZnO + H2O + 2e
2MnO2 + 2e + H2O Mn2O3 + 2OH

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

Photosensitive Goggles Photosensitive Goggles When this glass is exposed to UV light in the wavelength range 280–320nm, a latent image is formed. The glass remains transparent at this stage, but its absorption in the UV range of the spectrum increases. This increased absorption is only detectable using UV transmission spectroscopy. The reason behind this is suggested to be an oxidation-reduction reaction that occurs inside the glass during exposure in which cerium ions are oxidized to a more stable state and silver ions are reduced to silver.
Know why?
Why silver salts are stored in amber coloured bottles?

Silver salts are photo sensitive. When they are exposed to light, silver salt is reduced to give metallic silver. So they are stored in amber coloured bottles.

Photosensitive goggles

Many people wear photosensitive goggles. These are made up of lenses, which darken when exposed to bright light.

The basis of this change in color in response to light can be explained in terms of oxidation–reduction reactions. Lens material (glass or plastics) consists of a complex matrix of silicates, which is ordinarily transparent to visible light. In photosensitive 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 susceptibility to oxidation and reduction by light as described below.

Silver Chloride is Susceptibility

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 transmission of light, causing the lenses to darken. This process occurs almost instantaneously. The degree of “darkening” is dependent on the intensity of the light.

This 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 reaction occurs:

Darkening Process

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.

Copper Reacts with Silver

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.

 Flow Chart of Oxidation Reduction reactions An overview of redox reactions The formal name for a redox reaction is “oxidation reduction reaction,” and we can see that “redox” is just shorthand for the words reduction and oxidation. Thus, in a redox reaction, two things happen. These two have to happen together. We cannot have an oxidation reaction without a corresponding reduction reaction.
Redox reactions

Why do copper vessels and articles get a greenish tint when exposed to air?
Copper reacts with carbon dioxide in air to form copper carbonate. This is greenish in color and is known as patina. To polish the copper vessel and remove the patina, generally tamarind is used. Tamarind has tartaric acid (C4H6O6). The acid reacts with the carbonate to release carbon dioxide and water. The copper surface is restored to a bright and shiny surface.

Why do silver-wares turn black on exposure to air?
Silver reacts with molecules in air. With oxygen it becomes silver oxide. A layer of silver oxide, which is whitish in color, will always be present on any metal surface, including on silver-wares. Air contains small amounts of hydrogen sulphide; silver will react with it to become silver sulphide, which is black in color. Hence silver-wares turn black on exposure to air. This is generally removed with mild acids which can react with sulphur and remove the sulphide layer.

Summarizing oxidation—reduction reactions have many meanings: gain/loss of oxygen, loss/gain of hydrogen, loss/gain of electrons etc.

Then there is an added confusion about what are oxidizing agents and reducing agents. Oxidizing agents attract electrons and bring about reduction reactions in other substances and reducing agents lose electrons and bring about oxidation reactions in other substances! The flow chart may help in visualizing what is happening when oxidation–reduction reactions occur.

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