Redox Reactions Class 11 Notes Chemistry Chapter 8

Redox Reactions Class 11 Notes

This article provides you with Redox Reactions Class 11 Notes of Chemistry. The notes on redox reactions of class 11 chemistry have been prepared with great care keeping in mind the effectiveness of it for the students. This article provides the revision notes of the Redox Reactions chapter of Class 11 Chemistry for the students so that they can give a quick glance of the chapter.

Redox Reactions

The transfer of electrons from one chemical substance to another takes place through a number of different types of reactions. This electron – transfer reactions are termed as oxidation-reduction or redox-reactions.

Oxidation and Reduction Reaction

Oxidation may be defined as a process which involves the addition of oxygen or addition of non-metal or increase in +ve valency, or removal of hydrogen, removal of metal, loss of electrons. It may also be defined as the increase in oxidation number. In this process, electrons are liberated and are known as de-electronation.

Reduction is just reverse of oxidation. Reduction is a process which involves removal of oxygen or removal of non-metal or decrease in +ve valency, or addition of hydrogen, an addition of metal, gain of electrons. It may also be defined as the decrease in oxidation number. In this process, electrons are gained and are known as electronation.

Oxidizing and Reducing Agent:

An oxidizing agent or oxidant is a substance which supplies oxygen or any other electronegative element or removes hydrogen or any other electropositive element. An oxidizing agent after carrying out oxidation is itself reduced in a chemical reaction.

A reducing agent or reductant is a substance which supplies hydrogen or any other electropositive element or removes oxygen or any other electronegative element. A reducing agent after carrying out reduction is itself oxidised in a chemical reaction.

An element in a compound can function as an oxidizing agent if it is in its highest possible oxidation state.

(a) An element in a compound can function as an oxidizing agent if it is in its highest possible oxidation state. e.g. KMnO4, K2Cr2O7, HNO3, H2SO4, HClO4 etc.

(b) An element in a compound can function as a reducing agent if it is in its lowest possible oxidation state. e.g. H2S, FeSO4, Na2S2O3, SnCl2 etc.

(c) An element in a compound can function as an oxidizing agent as well as reducing agent if it is in its intermediate oxidation state. e.g. H2O2, H2SO3, HNO3, SO2 etc.

(d) A highly electronegative element in a compound can function as a powerful oxidizing agent if it is in its higher oxidation state. e.g. KClO4, KClO3, KIO3 etc.

(e) A highly electronegative element in a compound can function as a powerful reducing agent if it is in its lower oxidation state. e.g. I , Br , N3 etc.

MODERN CONCEPT OF OXIDATION AND REDUCTION

According to the modern concept, loss of electrons is oxidation whereas gain of electrons is reduction. Oxidation and Reduction can be represented in a general way as follows:

Redox Reactions Class 11 Notes

Redox reaction involves two half-reactions, one involving loss of electron or electrons (oxidation) and the other involving gain of electron or electrons (reduction).

ION ELECTRON METHOD FOR BALANCING REDOX REACTIONS

This method involves the following steps:

(a) Divide the complete equations into two half reactions-

(i) One representing oxidation
(ii) The other representing reduction

(b) Balance the atoms in each half-reaction separately according to the following steps-

(i) Balance all atoms other than oxygen and hydrogen
(ii) To balance oxygen and hydrogen

(c) Acidic Medium

(i) Add H2O to the side which is oxygen deficient to balance oxygen atoms
(ii) Add H+ to the side which is hydrogen deficient to balance H atoms

(d) Basic Medium

(i) Add OH¯ to the side which has less negative charge
(ii) Add H2O to the side which is oxygen deficient to balance oxygen atoms
(iii) Add H+ to the side which is hydrogen deficient

ION ELECTRON METHOD

Oxidation State

It is defined as the charge (real or imaginary) which an atom appears to have when it is in combination. In the case of electrovalent compounds, the oxidation number of an element or radical is the same as the charge on the ion.

Oxidation Number

Oxidation number of an element in a particular compound represents the number of electrons lost or gained by an element during its change from the free state into that compound or Oxidation number of an element in a particular compound represent the extent of oxidation or reduction of an element during its change from the free state into that compound. It is given a positive sign if electrons are lost and a negative sign if electrons are gained.

Oxidation number represents real charge in case of ionic compounds. However, in covalent compounds, it represents an imaginary charge.

Rules for Calculation of Oxidation Number

Following rules have been arbitrarily adopted to decide oxidation number of elements on the basis of their periodic properties.

(a) In uncombined state or free state, the oxidation number of an element is zero.

(b) In combined state oxidation number of-

(i) F is always –1.

(ii) O is –2. In peroxide, it is –1, in superoxides, it is –1/2. However, in F2O it is +2.

(iii) H is +1. In ionic hydrides, it is –1. (i.e., IA, IIA and IIIA metals).

(iv) Halogens as halide are always –1.

(v) Sulphur as sulphide is always –2.

(vi) Metal is always +ve.

(vii) Alkali metals (Li, Na, K, Rb, Cs, Fr) is always +1.

(viii) Alkaline earth metals (Be, Mg, Ca, Sr, Ba, Ra) is always +2.

(c) The algebraic sum of the oxidation number of all the atoms in a compound is equal to zero. e.g. KMnO4.

(Ox. no. of K) + (Ox. no. of Mn) + (Ox. no. of O) × 4 = 0
(+1) + (+7) + 4x (–2) = 0

(d) The algebraic sum of all the oxidation no. of elements in a radical is equal to the net charge on the radical. e.g. CO32.

(Oxidation no. of C) + 3 × (Oxidation no. of O) = –2(4) + 3x (–2) = –2

(e) Oxidation number can be zero, +ve, –ve (integer or fraction)

(f) Maximum oxidation no. of an element is = Group no. (Except O and F). Minimum oxidation no. of an element is = Group no. –8 (Except metals).

Redox reactions involve oxidation and reduction both. Oxidation means loss of electrons and reduction means a gain of electrons. Thus redox reactions involve electron transfer and the number of electrons lost are the same as the number of electrons gained during the reaction. This aspect of redox reaction can serve as the basis of a pattern for balancing redox reactions.

Oxidation number of Mn in KMnO4: Let the oxidation number of Mn be x. Now we know that the oxidation numbers of K is +1 and that of O is –2.

K          Mn              O4            or             K               Mn           O4

+ 1        + x             4 x (-2)                       +1              +x             -8

Now to the sum of oxidation numbers of all atoms in the formula of the compound must be zero, i.e. +1 +x – 8 = 0. Hence, the oxidation number of Mn in KMnO4 is +7.

Balancing of Redox Reactions by Oxidation State Method 

This method is based on the fact that the number of electrons gained during reduction must be equal to the number of electrons lost during oxidation. Following steps must be followed while balancing redox equations by this method.

(a) Write the skeleton equation of all the reactants and products of the reaction.

(b) With the help of the oxidation number of elements, find out which atom is undergoing oxidation/reduction, and write separate equations for the atom undergoing oxidation/reduction.

(c) Add the respective electrons on the right for oxidation and on the left for reduction equation. Note that the net charge on the left and right side should be equal.

(d) Multiply the oxidation and reduction reactions by suitable integers so as to equalize the total electrons lost in one reaction to that of the total electrons gained by other reaction.

(e) Transfer the coefficients of the oxidizing and reducing agents and their products as determined in the above step to the concerned molecule or ion.

(f) By inspection, put the proper coefficient for the other formulae of species not undergoing oxidation and reduction to balance the equation.

Balancing of Redox Reactions by Oxidation State Method 

TYPES OF REDOX REACTIONS

The redox reactions are of the following types:

(a) Combination reactions: A compound is formed by chemical combination of two or more elements. The combination of an element or compound with oxygen is called combustion. The combustion and several other combinations which involve the change in oxidation state are called redox reactions.

combination Reaction

(b) Decomposition reactions: Decomposition is the reverse process of combination, it involves the breakdown of the compound into two or more components. The product of decomposition must contain at least one component in the elemental state.

decomposition Reaction

In the above example, there is no change in the oxidation state of potassium. Thus, it should be noted that the decomposition does not result in a change in the oxidation number of each element.

(c) Displacement reactions: The reactions in which an atom or ion in a compound is displaced by another atom or ion are called displacement reactions. The displacement reactions are of 2 types:

(i) Metal displacement: In these reactions, a metal in a compound is replaced by another metal in an uncombined state. It is found that a metal with a stronger reducing character can displace the other metal having a weaker reducing character.

metal displacement Reaction

(ii) Non-metal displacement: These displacement reactions generally involve redox reactions, where the hydrogen is displaced. Alkali and alkaline earth metals are highly electropositive, they displace hydrogen from cold water.

non metal displacement Reaction

(d) Disproportionation and Oxidation–Reduction:  A reaction in which the same species may act simultaneously as an oxidizing agent with the result that a part of it gets oxidized to a higher state and rest of it is reduced to lower state of oxidation, i.e., a reaction in which a substance undergoes simultaneous oxidation and reduction is called disproportionation reaction and the substance is said to disproportionate.

The following are some of the examples of disproportionation:

disproportionation reaction

(e) The oxidation state of chlorine lies between –1 to +7; thus out of ClO, ClO2, ClO3, ClO4; ClO4 does not undergo disproportionation because in this oxidation state of chlorine is highest, i.e., +7. Disproportionation of the other oxoanions are:

Redox Reactions Class 11 Notes

Redox Reaction as the Basis for Titrations

Titration: It is an operation forming the basis of volumetric analysis. The addition of a measured amount of a solution of one reagent (called the titrant) from a burette to a definite amount of another reagent (called analyte) until the reaction between them is complete, i.e., till the second reagent (analyte) is completely used up, i.e., up to end point. The end point of the titration is normally detected by a sudden change in colour of the solution.

Indicator: These compound mixed in the solution in very small amount, which responses the sudden change in colour of the solution and show the end point of the titration. In an acid-base titration, the indicators used are either weak organic acid or weak organic bases. Some examples are

Acidic Indicator: Phenolphthalein, litmus paper etc.
Basic Indicator: Methyl orange, methyl red, etc.

Redox Titrations involves the titration of an oxidizing agent against a reducing agent or vice versa.

If the reagent itself is intensely coloured eg, MnO4; it can act as a self-indicator. If there is no auto-colour change, there are indicators which are oxidised immediately after the last drop of the reactant has reacted, producing a dramatic colour change eg. Cr2O72- oxidises the indicator substance diphenylamine just after the equivalence point to produce an intense blue colour, thus signalling the end point.  Starch is used as an indicator in case of those reagents which can either oxidise I ions such as Cu2+ or reduce I2 such as S2O32- ions as it gives intense blue colour with molecular iodine.

Redox Reactions and Electrode Processes

Redox couple: It is defined as having together the oxidised and reduced forms of a substance taking pat is an oxidation or reduction half-reaction i.e., a metal dipped in the solution of its own ions.

Electrode potential: The potential difference set up between the metal and its own ions in the solution is called the electrode potential. In general, it is the tendency of an electrode to gain or lose an electron. Standard electrode potential (E0): If the concentration of each species taking part in the electrode reaction is unity and further the reaction is carried out at 298 K, then the potential of each electrode is called standard electrode potential. Standard electrode potential of hydrogen is taken as 0.00 volts by convention.

Electrochemical series is a series in which a list of oxidising agents is arranged in decreasing order of their strength. It is also called activity or electromotive series.

A negative E0 means that the redox couple is a stronger reducing agent than the H+/H2 couple. A positive E0 means that the redox couple is a weaker reducing agent than the H+/H2 couple.

 

This article has tried to highlight the key points of redox reactions in the form of revision notes for class 11 students in order to understand the basic concepts of the chapter. The notes on redox reactions have not only been prepared for class 11 but also for the different competitive exams such as iit jee, neet, etc.

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