Redox Reactions

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Chapter 8 : Redox Reactions

8.1 Classical idea of redox reactions arrow_upward


  • The term “oxidation” is defined as the addition of oxygen/electronegative element to a substance or removal of hydrogen/ electropositive element from a substance.
  • Examples:
  • Addition of oxygen:
  • Removal of Hydrogen:
  • Electro-negative elements:

  • Reduction arrow_upward

  • Reduction is defined as the removal of oxygen/electronegative element from a substance or addition of hydrogen/ electropositive element to a substance.

  • Redox Reactions

  • Redox reactions form an important class of reactions in which oxidation and reduction occur simultaneously.
  • Example:
  • The reaction
  • is a redox reaction because sodium is oxidised due to the addition of oxygen and simultaneously oxygen is reduced.

  • 8.2 Redox Reactions in terms of Electron Transfer Reactions arrow_upward

  • Consider the formation of sodium chloride:
  • Each of the above steps is called a half reaction, which explicitly shows involvement of electrons. Sum of the half reactions gives the overall reaction:
  • Half reactions that involve loss of electrons are called oxidation reactions. Similarly, the half reactions that involve gain of electrons are called reduction reactions.
  • Sodium, which is oxidized, acts as a reducing agent because it donates electron to each of the elements interacting with it and thus helps in reducing them.
  • Oxygen is reduced and act as oxidizing agents because it accepts electrons from sodium.
  • Note:

    Oxidation: Loss of electron(s) by any species.

    Reduction: Gain of electron(s) by any species.

    Oxidising agent: Acceptor of electron(s).

    Reducing agent: Donor of electron(s).

    8.2.1 Competitive Electron Transfer Reactions

  • The reaction between metallic zinc and the aqueous solution of copper nitrate is:
  • Zn(s) + Cu2+ (aq)  Zn2+ (aq) + Cu(s)

  • In reaction, zinc has lost electrons to form Zn2+ and, therefore, zinc is oxidised. Copper ion is reduced by gaining electrons from the zinc.
  • Consider an electron transfer reaction of copper metal and silver nitrate solution in water.
  • The solution develops blue color due to the formation of Cu2+ ions on account of the reaction:
  • Here, Cu(s) is oxidised to Cu2+ (aq) and Ag+ (aq) is reduced to Ag(s).
  • By comparison we have come to know that zinc releases electrons to copper and copper releases electrons to silver and, therefore, the electron releasing tendency of the metals is in the order: Zn > Cu > Ag.

  • 8.3 Oxidation Number arrow_upward

  • Oxidation number denotes the oxidation state of an element in a compound ascertained according to a set of rules formulated on the basis that electron pair in a covalent bond belongs entirely to more electronegative element.

  • Rules to determine oxidation number:

  • In elements, in the free or the uncombined state, each atom bears an oxidation number of zero. Evidently each atom in H2 , O2 , Cl2 , O3 , P4 , S8 , Na, Mg, Al has the oxidation number zero.
  • 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,
  • The oxidation number of oxygen in most compounds is –2. But there are two types of exceptions - peroxides and superoxides.
    • In peroxides (e.g., H2 O2 , Na2 O2 ), each oxygen atom is assigned an oxidation number of –1,
    • In superoxides (e.g., KO2 , RbO2 ) each oxygen atom is assigned an oxidation number of –(½).
  • The oxidation number of hydrogen is +1, except when it is bonded to metals in binary compounds. For example, in LiH, NaH, and CaH2 , its oxidation number is –1.
  • Oxidation State is a term that is often used interchangeably with the oxidation number.
  • In CO2 , the oxidation state of carbon is +4.
  • 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.

  • 8.3.1 Types of Redox Reactions arrow_upward

  • There are four types of reactions:
    • Combination reactions
    • Decomposition reactions
    • Displacement reactions
    • Disproportionation reactions
    1. Combination reactions:
  • A combination reaction may be denoted in the manner:
  • All combustion reactions, which make use of elemental dioxygen, as well as other reactions involving elements other than dioxygen, are redox reactions.
  • Examples:

    2. Decomposition reactions
  • A decomposition reaction leads to the breakdown of a compound into two or more components at least one of which must be in the elemental state.
  • Examples:

  • All decomposition reactions are not redox reactions. For example, decomposition of calcium carbonate is not a redox reaction.
  • 3. Displacement reactions
  • In a displacement reaction, an ion (or an atom) in a compound is replaced by an ion (or an atom) of another element. It may be denoted as:
  • Displacement reactions fit into two categories:
    • Metal displacement and Non-metal displacement.
    Metal displacement:
  • A metal in a compound can be displaced by another metal in the uncombined state.
  • Examples:
  • Non-metal displacement:
  • The non-metal displacement redox reactions include hydrogen displacement and a rarely occurring reaction involving oxygen displacement.
  • 4. Disproportionation reactions
  • Disproportionation reactions are a special type of redox reactions.
  • In a disproportionation reaction an element in one oxidation state is simultaneously oxidised and reduced.
  • Example:
  • The decomposition of hydrogen peroxide:

  • 8.3.2 Balancing of Redox Reactions arrow_upward

  • There are two methods which are used to balance chemical equations for redox processes:
    • Oxidation Number Method
    • Half Reaction Method

    Oxidation Number Method:

  • Write the correct formula for each reactant and product.
  • Identify atoms which undergo change in oxidation number in the reaction by assigning the oxidation number to all elements in the reaction.
  • Calculate the increase or decrease in the oxidation number per atom and for the entire molecule/ion in which it occurs. If these are not equal then multiply by suitable number so that these become equal.
  • Ascertain the involvement of ions if the reaction is taking place in water, add H+ or OH ions to the expression on the appropriate side so that the total ionic charges of reactants and products are equal. If the reaction is carried out in acidic solution, use H+ ions in the equation; if in basic solution, and use OH ions.
  • Make the numbers of hydrogen atoms in the expression on the two sides equal by adding water (H2 O) molecules to the reactants or products. Now, also check the number of oxygen atoms.

  • Half Reaction Method:

  • In this method, the two half equations are balanced separately and then added together to give balanced equation.
  • Produce unbalanced equation for the reaction in ionic form:
  • Separate the equation into half reactions:
  • Balance the atoms other than O and H in each half reaction individually. Here the oxidation half reaction is already balanced with respect to Fe atoms. For the reduction half reaction, we multiply the Cr3+ by 2 to balance Cr atoms.
  • For reactions occurring in acidic medium, add H2 O to balance O atoms and H+ to balance H atoms.
  • Add electrons to one side of the half reaction to balance the charges.
  • The oxidation half reaction is thus rewritten to balance the charge:
  • Now in the reduction half reaction there are net twelve positive charges on the left hand side and only six positive charges on the right hand side. Therefore, we add six electrons on the left side.
  • We add the two half reactions to achieve the overall reaction and cancel the electrons on each side.
  • Verify that the equation contains the same type and number of atoms and the same charges on both sides of the equation.
  • For the reaction in a basic medium, first balance the atoms as is done in acidic medium. Then for each H+ ion, add an equal number of OH ions to both sides of the equation.

  • 8.3.3 Redox Reactions as the Basis for Titrations arrow_upward

  • In acid-base system titration is a method use to find the strength of one solution against the other using a pH sensitive indicator.
  • Similarly, in redox systems, the titration method can be adopted to determine the strength of a reductant/oxidant using a redox sensitive indicator.
  • In one situation, the reagent itself is intensely coloured, e.g., permanganate ion, MnO4 . Here MnO4   acts as the self-indicator. The visible end point in this case is achieved after the last of the reductant (Fe2+ or C2 O4 2– ) is oxidised and the first lasting tinge of pink colour appears at MnO4 concentration as low as 10–6 mol dm–3 .
  • If there is no dramatic auto-colour change (as with MnO4 titration), there are indicators which are oxidised immediately after the last bit of the reactant is consumed, producing a dramatic colour change.
  • There is yet another method which is interesting and quite common. Its use is restricted to those reagents which are able to oxidise I– ions, say, for example, Cu(II):

  • 8.3.4 Limitations of Concept of Oxidation Number

  • Oxidation process is visualised as a decrease in electron density and reduction process as an increase in electron density around the atom(s) involved in the reaction.

  • 8.4 Redox Reactions and Electrode Processes arrow_upward

  • A redox reaction in which oxidation and reduction takes place in the vessel is called direct redox reaction.
  • The transference of electrons from reducing agent to oxidising agent occurs over a very short distance.
  • A redox couple is defined as having together the oxidised and reduced forms of a substance taking part in an oxidation or reduction half reaction. In this experiment the two redox couples are represented as Zn2+ /Zn and Cu2+ /Cu.
  • Put the beaker containing copper sulphate solution and the beaker containing zinc sulphate solution side by side.
  • The zinc and copper rods are connected by a metallic wire with a provision for an ammeter and a switch. The set-up as shown in the figure given below is known as Daniell cell.
  • Electrode Processes
  • Electrochemical conversions that occurat an electrode electrolyte interface.
  • Where a charge is transferred through theinterface and an electriccurrent flows.
  • Depending on the direction of the electron flow (from the electrode to the material or vice versa) a distinction is made betweencathodic and anodic processes, which result in the reduction and oxidation, respectively, of the material.
  • The spatial separation of the oxidationand reduction processes is used in chemical sources of electric current and for electrolysis.

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