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We can observe many chemical reactions in our surroundings such as burning, dissolution, respiration, photosynthesis or corrosion etc. A chemical reaction can be recognized with its components; reactants, reagent and product.

For example in a chemical reaction two reactant molecules; A & B are reacted together to form AB. Therefore both A and B are reactants that are converted in to product AB. The product AB can act as reactant for other chemical reaction such as can decompose to A and B or can react with other reactant molecule ‘C’ to form a new product.
For example, the reaction of magnesium metal with oxygen forms magnesium oxide. On the contrary, the decomposition reaction is just opposite to that and involves the cleavage of reactant molecules to small fragment such as decomposition of calcium carbonate results the formation of calcium oxide and release carbon dioxide gas. The displacement reactions involve the replacement of one of the component of reactant molecule by other. If it happens once, it will known as single displacement reaction while double displacement reaction involves the replacement of both components of recant molecules.

For example the reaction of zinc metal with hydrochloric acid forms zinc chloride and releases the hydrogen gas or reaction solid aluminum metal with aqueous solution of silver nitrate forms aluminum nitrate and solid silver metal. The reaction equation can be written as
Zn + 2 HCl $\rightarrow$ ZnCl2 + H2
3AgNO3 + Al $\rightarrow$ Al(NO3)3 + 3Ag

Another type of chemical reaction is Oxidation-reduction reaction or also known as Redox reaction that is very important, occurs in our surroundings. The redox reaction involves several biological processes such as metabolic pathways in living systems, photosynthesis and respiration etc.

In photosynthesis process, carbon dioxide is reduced to organic matter in the presence of solar energy. On the contrary, the decomposition of organic matters in the presence of oxygen releases a good amount of energy through an oxidation reaction in non-photosynthetic organisms. Overall a chemical reaction involves the rearrangement of atoms of molecules to form new molecules. The total mass on the both side of chemical equation remains same before and after the chemical reaction.

The balance of a chemical reaction with the help of stoichiometric coefficients represents the proportions of whole molecules involve during the chemical reaction. For example, in the given reaction, one molecule of methane reacts with two molecules of oxygen to form one molecule of carbon dioxide with two molecules of water. We can write the balanced reaction as given below.

Balanced Chemical Reaction

Definition

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As we know that a redox reaction is a combination of two reactions; oxidation and reduction, each of these two reactions occur separately but related to each other. Overall if one species oxidized in a chemical reaction, another gets reduced.

Each half reaction represents either oxidation or reduction process. There is a driving force for both reactions to occur that is called as reaction potential. For oxidation reactions, the driving force is known as oxidation potential while for reduction reactions, it is called as reduction potential.

Both oxidation and reduction potentials are related to each other with opposite signs. Therefore the total redox potential is the sum of oxidation and reduction potential that can measure under standard conditions such as one bar pressure and at pH = 7.

The redox potential can define as a measurement of the affinity of a substance towards electrons or the electron negativity of the substance compare to hydrogen atom whose electron negativity is considered as zero.

Standard Redox Potential

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We know that the redox potential is sum of oxidation and reduction potential, but in general we are using the standard redox potential only. The standard redox potential or standard cell potential indicates the potential difference between the cathode and anode under standard conditions such as a 298 K temperature, 1 atmospheric pressure with one molar solution and denoted as E0.

The standard reduction potentials are written for reduction reactions only, for example the standard reduction potential of an element ‘A’ with charge x can be written as given below.
Ax+ + x * e- $\rightarrow$ A (s)

Like for copper (II) ion;

Cu2+ (1M ) + 2* e- $\rightarrow$ Cu(s); E0º (reduction) = +0.340 V

Similarly the standard oxidation potential indicates the tendency of a substance to oxidize under standard conditions such as the standard oxidation potential for copper can be written as given below.

Cu (s) $\rightarrow$ Cu2+ (1M) + 2* e-; E0º (oxidation) = -0.340 V

The magnitude of standard oxidation and reduction potential remains same only the sign gets opposite for the same chemical substance. Therefore we can write.

E
0º (reduction) = -E0º (oxidation)

  • The standard redox potential can determine with the help of standard hydrogen electrode whose electrode potential is considered as zero. Therefore the total cell potential will be either oxidation or reduction potential under standard conditions.
  • Standard hydrogen is composed of a glass tube filled with hydrogen gas and contains a platinum wire with platinum electrode, dipped in one molar solution of hydrogen ions at 298 K temperature and one atmosphere pressure of gas.
  • The oxidation and reduction potential of hydrogen is considered as zero. Hence it can use for the determination of standard potential of certain substances with the help of a galvanic cell.
  • The difference in the cell potential is actually the oxidation or reduction potential of another substance that is used with the standard hydrogen electrode. If the unknown substance is reduced and hydrogen is oxidized, the standard electrode potential represents the standard reduction potential.
  • On the contrary, if the unknown substance is oxidized and hydrogen is reduced, the potential represents the oxidation potential of that chemical substance.
For example; if we take a solution of copper ions with standard hydrogen electrode, the cell potential gives the value of the standard reduction potential of copper. The electrons flow from hydrogen electrode to copper solution and reduced it to solid copper that deposited on the electrode. The reaction occurs in the reaction system is as follows

Cu2+(aq) + 2e-
$\rightarrow$ Cu(s) E = +0.34 V

The standard reduction potential is generally used for the determination of standard cell potential or standard redox potential that we can check in standard redox potential table. The relation between standard reduction potential and standard oxidation potential with the standard cell potential can be written as below.

Ecellº = E0º (reduction) + E0º (oxidation)
Ecellº = E0º(cathode) + E0º (anode)

The activity series of elements is also based on the standard reduction potential of them from low value to high value (negative to positive). The most negative value of the standard reduction potential of elements makes it a good reducing agent while the element with positive value of standard potential will be good oxidizing agent.

Discontinuity

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Like Gibbs free energy, the standard redox potential is also not an absolute value. It can measure with the help of standard hydrogen electrode taken as standard with zero potential value at standard temperature and pressure values with one mole of the solution. The overall cell potential is the sum of oxidation and reduction potential of the substances involve in the chemical processes.

E0º ox represents the oxidation potential of half oxidation reaction while E0º red is potential of reduction half reaction. The relation between the redox potential under standard and non-standard condition can be shown with the help of Nerns't equation as given below.

E = E0º – RT/nF ln Q

Here R is a gas constant, T represents the absolute temperature in Kelvin, n shows the number of electrons transferred in the reaction; F is the Faraday number and Q represents the reaction quotient.

Table

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The redox potential chart represents the redox potential values for various redox reactions. As discussed, usually we use the standard reduction potential of the elements. Some of the chemical substances with their oxidized and reduced forms and standard reduction potential values are given below.


Sl. No. Oxidized species Number of electrons involve in the reaction Reduced species Standard Reduction Potential (Volts)
1 Ag+
1 Ag 0.7996
2 AgBr 1 Ag + Br- 0.0713
3 AgNO2 1 Ag2+ NO2- 0.564
4 Ag2S 2 2 Ag + S2- -0.691
5 Ag2S + 2 H+ 2 2 Ag + H2S -0.0366
6 Al3+ 3 Al -1.662
7 Cl2 2 2 Cl-(g) 1.3583
8 Cr2O72- 6 2 Cr3+ 1.36
9 F2 + 2 H+ 2 2HF 3.053
10 Fe2+ 2 Fe -0.447
11 H2O2 + 2 H+ 2 2 H2O 1.776
12 Li+ 1 Li -3.0401
13 Na+ 1 Na -2.71
14 O2 + 2 H+ 2 2 H2O2 0.695
15 PbSO4 2 Pb + SO42 -0.3588
16 SO42- + H2 2 SO32- + 2 OH- -0.93
17 Zn2+ 2 Zn -0.7618
18 Br2() 2 2 Br- 1.066
19 AgI 1 Ag + I- -0.1522
20 Ag3+ 3 Ag2+ 1.8

Redox Reactions

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The study of oxidation and reduction reactions is very important to understand several natural concepts such as photosynthesis, corrosion, rancidity, burning etc. Both oxidation and reduction reactions are complementary to each other and combine known as Redox reaction. We can explain the oxidation reaction as the reaction of addition of oxygen or increment in the oxidation number or removal of electrons. The reverse reaction is known as reduction i.e. removal of oxygen, addition of electron or reduction in oxidation number. We can use any one of these concept to explain the oxidation and reduction reactions.

Both reactions are always coupled with each other as one of the molecule loose electrons and oxidized, another accept the same electrons to reduce. The substance that gets reduced in the process is called as oxidizing agent while the substance that gets reduced by gaining electrons is called as reducing agent as it gets oxidized during process.

Let’s take some common examples of oxidation and reduction reaction. The reaction of solid sodium metal with chlorine gas forms solid sodium chloride. Here sodium is a pure metal and chlorine gas is also in elemental state with oxidation state of zero. The reaction between both reactants forms sodium chloride in which sodium exits in +1 oxidation state and chlorine is in -1 oxidation state.

Therefore sodium is oxidized by chlorine which acts as oxidizing agent and chlorine is reduced to chloride ion, hence sodium acts as reducing agent. The reaction of metals and non-metals with oxygen is also a good example of redox reactions such as the reaction of magnesium with oxygen s results the formation of magnesium oxide, reaction of iron with oxygen forms ferrous oxide and carbon with oxygen forms carbon dioxide.

In all the given reactions, metals are exiting in their elemental stage with zero oxidation number and same with oxygen molecule. But in the product side, oxygen exits in oxide form, hence the oxidation number decreases from 0 to 2- which show the reduction of oxygen. The metal atoms oxidize from 0 to +1 and act as reducing agent. The last reaction of formation of carbon dioxide form carbon and oxygen is an example of combustion reaction that is also a redox reaction. Here the oxidation number of carbon changes from zero to 4+ that indicates the oxidation and the oxidation number of oxygen changes from zero to 2- that is reduction. Therefore carbon acts as reducing agent and oxygen acts as oxidizing agent.

Redox Reactions Examples

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Let’s discuss some common examples of redox reactions. We can observe many examples of redox reactions during the study of heat and work.
  • The burning of natural gas is an oxidation-reduction reaction that releases around 800 kj/mol of energy with carbon dioxide and water.
CH4(g) + 2 O2(g) $\rightarrow$ CO2(g) + 2 H2O(g)
  • Similarly the oxidation of glucose molecules to carbon dioxide and water during respiration reaction or decomposition of fatty acids to CO2 and water is also oxidation-reduction reaction.
Both of these reactions are exothermic in nature and release a good amount of energy in living bodies. Another example of oxidation-reduction reaction is the tarnishing of silver in the presence of atmospheric oxygen gas. Similar example is corrosion of iron in the presence of oxygen to form ferrous oxide and ferric oxide which are deposited on the metal surface and called as rust. The reaction of molten iron with water also forms ferrous ions with hydrogen gas.

All these redox reactions occur at slow speed at room temperature while can speed up under specific reaction conditions such as temperature, pressure or even catalyst. The corrosion of iron surface is a slow process that involves the reaction of solid metal with moisture and atmospheric oxygen to form oxides of iron in hydrated form. This hydrated oxide further oxidized to form hydrate form of ferric oxide. This oxidation-reduction process can prevent to apply a coat of paint on the metallic surface which prevent the attack of moisture and atmospheric oxygen on the metal surface.

Let’s discuss the change in the oxidation number of different components of reactions in a redox reaction of methane with oxygen to form water and carbon dioxide.

CH4 + O2 $\rightarrow$ H2O + CO2

The oxidation number of carbon in methane is 4- that turn to 4+ in carbon dioxide molecule. Therefore the increment in the oxidation number shows the oxidation. The oxidation number of oxygen changes from zero to 2- that is reduced; hence it acts as an oxidizing agent.

The change in the oxidation number of atoms is associated with release in energy in the form of heat. Overall the increment in the number of X-O bonds or reduction in the X-H bond represents the oxidation reaction, here X represents an element. The oxidation-reduction reaction also occurs in batteries, or during plating or deposition of a pure metal from a solution of that metal such as silver coating on copper surface can be done by dipping a copper rod in silver nitrate solution that form solid silver deposited on copper rod. Here copper is oxidized and acts as reducing agent while silver ions are reduced and act as oxidizing agent.

In the reverse reaction; solid silver with copper ion’s solution forms solid copper metal and silver ions in the solution. For this reverse reaction, solid silver oxidizes and acts as reducing agent while copper ions are reduced and act as an oxidizing agent. Therefore, the copper metal should deposit on the surface of solid silver. But this reverse reaction cannot occur that can be explained with the activity series of metals.

Similarly the reaction of solid zinc with a copper ion solution forms solid copper and zinc ions in the solution. Here zinc oxidizes to zinc (II) ions and copper ions reduce to solid copper metal. This redox reaction is generally used in the batteries.

Redox Equations

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A complete redox equation involves the oxidizing agent and reducing agent in the reaction with appropriate stoichiometric coefficient. Compare to other chemical reactions, it is a bit difficult to balance the redox equations. There are mainly two methods to balance the redox reaction, one is oxidation number method and another is half reaction or ion-electron method.

In the balancing of redox equation, we mainly need the oxidation number of reactants and products involve in the reaction. As we know that any redox reaction is a combination of oxidation and reduction reactions, therefore the complete redox reaction can divide into two half reactions. In these half reactions, each reaction represents either oxidation or reduction process.

Let’s learn how to balance a redox reaction with the help of these half reactions. The method of balancing a redox reaction with the help of half reactions is known as an ion-electron method. We have to balance each half reaction separately and then combine them to make a complete balance oxidation-reduction reaction.

During this method, we have to balance the charges and the number of ions on both the sides of the reaction. Each oxidation or reduction reaction is an Ionic equation that has to balance and combine to get molecular equation.

For example; the reaction of an aqueous solution of nitric acid with solid copper forms cupric nitrate, nitric oxide and water as given below.

Cu + HNO3 $\rightarrow$ Cu (NO3)2 + NO + H2O

Now first we have to write the Ionic equation of the given molecular equation in which nitric acid and copper nitrite can be written in Ionic form, therefore the Ionic equation would be

Cu + H+ + NO3- $\rightarrow$ Cu2+ + 2NO3- + NO + H2O

Now check the oxidation numbers of all the species involve in the reaction. We know that the oxidation and reduction process. The oxidation number of copper gets changed to zero to 2+ while for hydrogen, it remains same. But the oxidation number of nitrogen changes from -2 to +2 in the half reaction.

Cu + H++ NO3- $\rightarrow$ Cu2+ + 2NO3- + NO + H2O
Hence the overall change in the oxidation state will be;

Cu $\rightarrow$ Cu2+ ; NO3- $\rightarrow$ NO

Now balance all the atoms excluding oxygen and hydrogen as it is much easier to balance the oxygen and hydrogen at the end of the method. In acidic medium, oxygen can balance with the addition of water molecules while in alkaline medium; hydroxyl ions are used to balance the oxygen atom.

Since, the given reaction occurs in acidic medium therefore the oxygen atom can be balanced with the help of water molecules. Therefore the half reaction would be

Cu $\rightarrow$ Cu2+ ; NO3- $\rightarrow$ NO + 2H2O

For balancing the hydrogen atoms; hydrogen ions are used in acidic medium while water molecules are used in alkaline medium. So let’s balance the hydrogen ion in the reaction;
Cu $\rightarrow$ Cu2+ ; 4 H+ + NO3- $\rightarrow$ NO + 2H2O

Now we have to balance the charge on each of the half reaction. Since there is a 2+ charge on the right side the copper reaction, therefore for balancing the charge, we must add two electrons on the right side. Similarly, three electrons require on the left side of the nitrate reaction.

Oxidation: Cu(s) $\rightarrow$ Cu2+ + 2e-
Reduction: 3e- + 4 H+ + NO3- $\rightarrow$ NO + 2H2O

Now balance the charge on both the reactions by multiplying with the common factor.

Oxidation: 3 x [Cu(s) $\rightarrow$ Cu2+ + 2e-]
Reduction: 2 x [3e- + 4 H+ + NO3- $\rightarrow$ NO + 2H2O]

Now addition of both of these reactions will give the complete redox equation that is balance on both sides.

3Cu + 8HNO3 $\rightarrow$ 3Cu (NO3)2 + 2NO + 4H2O

Fuel Cell

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  1. A device that can convert the chemical energy to electrical energy through chemical reactions is called as fuel cell.
  2. It contains two oppositely charged electrodes; anode and cathode, dipped in a certain solution that is known as electrolyte.
  3. The reaction between electrodes and solution generates electricity.
  4. Some of the fuel cells also contain catalyst that can speed up the reactions and produce electricity with fast speed.
  5. There are two basic fuels in the fuel cells; hydrogen and oxygen. Therefore these kinds of cell produce minimum pollution and generate harmless side products such as water.
  6. The combination of many fuel cells in a series produces a large amount of electricity that can use in various applications like illumination of light, electric motor etc.
On the basis of type of fuel or operating technique, fuel cells can be different types. Usually, the hydrogen atoms are used in the fuel cell to ionize at the anode and form positively charged hydrogen ions. The electrons from hydrogen atom flow through wire to generate the direct current.

The oxygen gas that acts as another fuel for fuel cell, present at the cathode and combine with electrons and hydrogen ions to form water molecules.

Fuel Cell

The presence of electrolyte helps in the flow of ions and electrons between both electrodes. The only need of the fuel cells is the regular supply of hydrogen and oxygen gas as they consume and form water molecules as a side product. The fuel cells are more efficient in terms of their energy, less waste material and more supply with no limitation of thermodynamic laws.

Generally fuel cells are designed for use in vehicles to produce less than 1.16 volts of electricity which is enough to power a vehicle. That is the reason; the fuel cells are assembled into a fuel cell stack. The number and size of fuel cells determine the power generation from the assembled cell.

Solid Oxide Fuel Cell

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Different fuel cells contain different electrolytes like alkali (potassium hydroxide) in the alkali fuel cell, carbonate salts of sodium or magnesium in molten carbonate cells, phosphoric acid in a phosphoric acid fuel cell, polymer in proton exchange membrane fuel cells and metal oxides in solid oxide fuel cells.

The solid oxide fuel cells contain a ceramic compound of metal oxides such as calcium or zirconium oxides as an electrolyte for the chemical reaction. The efficiency value for solid oxide fuel cell is 60% and can only 1273 K of temperature. No doubt the output of such fuel cell is up to 100 kW and the recycling process of waste heat provides additional electricity.
Solid Oxide Fuel Cell

Due to high temperature, no reformer is required to generate the hydrogen gas for fueling the cell. The presence of high temperature limits the uses of fuel cells and responsible for the larger size. The warmed air at high temperature enters from the cathode side of the fuel cell and steam enters from another side of fuel cell that is anode side to produce reformed fuel.

The layer of electrolyte is placed in between permeable electrodes and impermeable for both fuels. The oxygen ions react with fuel to form water, carbon dioxide with electricity. The water which is formed as side product can recycle again to reform the fuel.