Redox Reaction

Front 1

What are redox reactions?

Back 1

Those reactions in which oxidation and reduction occurs simultaneously, are called redox reactions.

Front 2

Define oxidation.

Back 2

“oxidation” is defined as the additionof oxygen/electronegative element to asubstance or removal of hydrogen/electropositive element from a substance.

A substance which supplies oxygen or any other electronegative elements or removes hydrogen or any other electropositive element is called an oxidising agent.

Front 3

Define reduction.

Back 3

removalof oxygen/electronegative element from asubstance or addition of hydrogen/electropositive element to a substance.

A substance which supplies hydrogen or any other electropositive element or removes oxygen or any other electronegative element is called a reducing agent.

Front 4

Redox reaction in terms of electronic concept.

Back 4

Sum of the half reaction gives the overall Redox reaction. Half reaction that involve loss of electron are called Oxidation reaction. Half reaction that involve gain of electron are called reduction reaction.

Oxidising agent : Acceptor of electron(s).

Reducing agent : Donor of electron(s).

Front 5

Electron releasing tendency of Zn, Cu and Ag

Back 5

Zn> Cu> Ag

Front 6

Oxidation number

Back 6

The oxidation number is defined as the charge which appears on an atom of the element when all the other atoms attached to it are removed in the form of their ions. Oxidation number is also called the oxidation state.

Oxidation number of an element in its elementary form is zero, for example:- H2, O2, N2, Cl2, P4, Na, Mg, Fe, etc. have zero oxidation number.

For ions composed of only one atom, the oxidation number is equal to the charge on the ion. Thus Na+ ion has an oxidation number of +1, Mg2+ ion, +2, Fe3+ ion, +3,Cl– ion, –1, O2– ion, –2; and so on. In their compounds all alkali metals have oxidation number of +1, and all alkaline earth metals have an oxidation number of+2. Aluminium is regarded to have an oxidation number of +3 in all its compounds. In compounds with non-metals oxidation number of hydrogen is +1 (eg., HCl, H2S, H2O) but in metal hydrides ( eg., LiH, NaH, CaH2, MgH2) the oxidation number of hydrogen is -1.

The oxidation number of oxygen in most compounds is –2. However, we come across two kinds of exceptions here. One arises in the case of peroxides and superoxides, the compounds of oxygen in which oxygen atoms are directly linked to each other. While in peroxides (e.g., H2O2, Na2O2), each oxygen atom is assigned an oxidation number of –1, in superoxides (e.g., KO2,RbO2) each oxygen atom is assigned anoxidation number of –(½). The second exception appears rarely, i.e. when oxygen is bonded to fluorine. In such compounds e.g., oxygen difluoride (OF2) and dioxygen difluoride (O2F2), the oxygen is assigned an oxidation number of +2 and +1,respectively. This is due to the fact that fluorine is the most electronegative element known, always has an oxidation number of -1.

The algebraic sum of the oxidation number of all the atoms in a compound must be zero. In polyatomic ion, the algebraic sum of all the oxidation numbers of atoms of the ion must equal the charge on the ion.

The metallic elements have positive oxidation number and non metallic elements have positive or negative oxidation number. The atoms of transition elements usually display several positive oxidation states.

Front 7

According to oxidation number, define oxidation, reduction, oxidising agent and reducing agent.

Back 7

Oxidation: An increase in the oxidation number of the element in the given substance.

Reduction: A decrease in the oxidation number of the element in the given substance.

Oxidising agent : A reagent which can increase the oxidation number of an element in a given substance. These reagents are called as oxidants also.

Reducing agent: A reagent which lowers the oxidation number of an element in a given substance. These reagents are also called as reductants.

Redox reactions: Reactions which involve change in oxidation number of the interacting species.

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Combination reactions

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A combination reaction may be denoted in the manner:

A + B --> C

Either A and B or both A and B must be in the elemental form for such a reaction to be a redox reaction. All combustion reactions, which make use of elemental dioxygen, as well as other reactions involving elements other than dioxygen, are redox reactions. Some important examples of this category are in the picture.

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Decomposition reactions

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Decomposition reactions are the opposite of combination reactions. Precisely, a decomposition reaction leads to the breakdown of a compound into two or more components are least one of which must be in the elemental state. All decomposition reactions are not redox reactions. For example, decomposition of calcium carbonate is not a redox reaction.

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Displacement reactions

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In a displacement reaction, an ion in a compound is replaced by an ion of another element. Displacement reactions fit into two categories: metal displacement and non-metal displacement.

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Metal displacement

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A metal in a compound can be displaced by another metal in the uncombined state. Metal displacement reactions find many applications in metallurgical processes in which pure metals are obtained from their compounds in ores. In each reaction, the reducing metal is a better reducing agent than the one that is being reduced which evidently shows more capability to lose electrons as compared to the one that is reduced.

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Non metal displacement

Back 12

The non-metal displacement redox reactions include hydrogen displacement and a rarely occurring reaction involving oxygen displacement.

i) All alkali metals and some alkaline earth metals (Ca, Sr, and Ba) which are very good reductants, will displace hydrogen from cold water.

Hint 12

ii) Less active metals such as magnesium and iron react with steam to produce dihydrogen gas.

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Non metal displacement (continued)

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iii) Many metals, including those which do not react with cold water, are capable of displacing hydrogen from acids. Dihydrogen from acids may even be produced by such metals which do not react with steam. Cadmium and tin are the examples of such metals.

A few examplesfor the displacement of hydrogen from acids are:

Zn(s) + 2HCl(aq)------> ZnCl2 (aq) + H2 (g)

Mg (s) + 2HCl (aq)-----> MgCl2 (aq) + H2 (g)

Fe(s) + 2HCl(aq)-----> FeCl2(aq) + H2(g)

The above reactions are used to prepare dihydrogen gas in the laboratory. Here, the reactivity of metals is reflected in the rate of hydrogen gas evolution, which is the slowest for the least active metal Fe, and the fastest for the most reactive metal, Mg.

Very less active metals, which may occur in the native state such as silver (Ag), and gold (Au)do not react even with hydrochloric acid.

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Reactivity of fluorine

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Like metals, activity series also exists for the halogens. The power of these elements as oxidising agents decreases as we move down from fluorine to iodine in group17 of the periodic table. This implies that fluorine is so reactive that it can replace chloride, bromide and iodide ions in solution. In fact, fluorine is so reactive that it attacks water and displaces the oxygen of water :

2H2O (l) + 2F2 (g)-----> 4HF(aq) + O2(g)

It is for this reason that the displacement reactions of chlorine, bromine and iodine using fluorine are not generally carried out in aqueous solution.

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Layer test

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As Br2 and I2 are coloured and dissolve in CCl4, can easily be identified from the colour of the solution. This forms the basis of identifying Br– and I – in the laboratory through the test popularly known as ‘Layer Test’.

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Disproportionation reactions

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In a disproportionation reaction an element in one oxidation state is simultaneously oxidised and reduced. One of the reacting substances in a disproportionation reaction always contains an element that can exist in at least three oxidation states. The element in the form of reacting substance is in the intermediate oxidation state; and both higher and lower oxidation states of that element are formed in the reaction. The decomposition of hydrogen peroxide is a familiar example of the reaction, where oxygen experiences disproportionation.

-1 -2 -1

2H2O2 (aq)-------> 2H2O(l) + O2(g)

Here the oxygen of peroxide, which is present in –1 state, is converted to zero oxidation state in O2 and decreases to –2 oxidation state in H2O.

Phosphorous, sulphur and chlorine undergo disproportionation in the alkaline medium as shown below :

P4(s) + 3OH–(aq)+ 3H2O(l)------> PH3(g) + 3H2PO2–(aq)

S8(s) + 12 OH– (aq)-------> 4S2– (aq) + 2S2O32–(aq)+ 6H2O(l)

Cl2 (g) + 2 OH– (aq)--------> ClO– (aq) + Cl– (aq) +H2O (l)

The last reaction describes the formation of household bleaching agents. The hypochlorite ion (ClO–) formed in the reaction oxidises the colour-bearing stains of the substances to colourless compounds.

Among halogens, fluorine does not show a disproportionation tendency.

The reaction thattakes place in the case of fluorine is as follows:

2 F2(g) + 2OH–(aq)------> 2 F–(aq) + OF2(g) + H2O(l)

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Fractional oxidation number

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Hint 17