Important Notes on Redox Reaction

Redox Reactions

Redox reaction is a reaction mechanism wherein the oxidation state of the metal changes after the reaction. The oxidation state may increase or decrease depending on the type of the Redox Reaction. Formation of the rust on the metal surface is one such example wherein iron gets oxidised into iron oxide. Similarly, another example is ghee preparation, wherein the reduction reaction takes place. These reactions - oxidation and reduction reactions together are known as redox reactions. 

1. The concept of Oxidation and Reduction

The concept of Oxidation and Reduction can be understood simply as addition or removal of oxygen from a compound. Further, we can understand this concept in various terms. 

• According to oxygen transfer 
  • Oxidation is the gain of oxygen 
  • Reduction is the loss of oxygen  
  • 
  • Fe (III) is converting into Fe (O) - decrease. Similarly, the oxidation state of carbon is increased. 
  • This is known as the redox reaction as oxidation and reduction is happening simultaneously.
• According to Hydrogen Transfer 
  • The loss of Hydrogen is the Oxidation 
  • The gain of Hydrogen is the reduction 
  • 
• According to electron transfer - The loss of electrons is Oxidation and the gain of electrons is the reduction.   
  • 
  • The above redox reactions can be considered as two half reactions. 
    • 2Na -> 2 Na⁺ (loss of electron)
    • Cl₂ -> 2Cl^- (gain of electron) 
  • Summation of two half-reactions will give redox reaction.  

2. The concept of Oxidising and Reducing Agent 

• Oxidizing agents gets self-reduced by accepting electrons and oxidises another compound. Oxidising agents supplies oxygen or any other electronegative elements to another substance. Thus, it is a substance that can accept one or more electrons. 
• Reducing agents gets self-oxidised by donating electrons and reduces another compound. Reducing agents supplies hydrogen or any other electropositive elements to another substance. Thus, it is a substance that can give one or more electrons. 
• When oxidation happens, at the same time reduction also happens. Thus, oxidation-reduction reactions happen simultaneously always. 
• Example: 

Summary of the concept - Remember it!

Reducing agent / Oxidation - Loss of the electron 

Oxidising agent / Reduction - Gain of the electron

 3. Oxidation Number

The oxidation number is the charge on the atom of the element when the attached atoms are removed in form of their ions. 

Steps to calculate the Oxidation of an atom in a molecule or ion

• An element in its elementary form has zero oxidation number. For example, H₂, N₂, O₂, P₄, Na, Fe etc. 
• For a single monoatomic ion, the oxidation state(number) is same as the charge on the ion. For example, Mg⁺² has +2 charge, thus the oxidation state is +2. 
• The oxidation number of hydrogen: It varies with the compound depending on whether the compounds contains metal/ non - metal. When the compound comprises of non - metals, the oxidation state of hydrogen is +1. For example, HCl, H₂O etc. When the compound comprises of metal, then the oxidation number of hydrogen is -1. For example, CaH₂, MgH₂ etc. 
• The oxidation number of oxygen: Usually the oxidation state of oxygen is -2. However, there are exceptions where the oxidation state varies depending on the atom present in the compound. Exceptions: In case of peroxides (H₂O₂, BaO₂ etc.). In the case of OF₂, O₂F₂, the oxidation number of oxygen is +2 and +1 respectively, because Fluorine is the most electronegative element and always has an oxidation number of -1. 
• Metal has a positive oxidation number, while non-metal has a negative oxidation number. For example, NaCl, Na has +1 oxidation number and Cl has -1 oxidation number. 
  • In a case when there are two non- metallic atoms, the negative oxidation number is assigned to the atom with higher electronegativity whereas the positive oxidation number is assigned to the atom with comparatively lower electronegativity. 
• In the case of neutral compounds, the sum of the oxidation number of all the atoms involved vanishes (equals to zero). 
• In the case of ions, the sum of the oxidation number of all the atoms involved equals to the charge of the ion (positive or negative) 

Nomenclature of compounds where oxidation number is shown

Many metal compounds show more than one oxidation states. These compounds are named with certain rules to distinguish from one another. The rule is: Oxidation state of the metal is shown in the parenthesis after the name of the metal. 
For example FeCl₂ - Iron (II) Chloride, FeCl₃ - Iron (III) Chloride etc. 

The case of Fractional Oxidation States

There are a few cases where the oxidation state is not the whole number, instead, it is a whole number. However, this is not possible as per the classical definition, because electrons cannot be shared or transferred in a fraction.  This case is understood after understanding the structure of such species. It has been seen that, one element is involved in different bonding and this makes the overall oxidation states to be fractional. In actual, the oxidation number is the whole number only, but on an average, the oxidation state is in the fractional state. 

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4. Types of Redox Reactions

a. Combination Reactions

When two or more substances combine to form a single substance are called as combination reactions. If combination reactions have to be redox, then one or more substance must be in the elemental form. 

b. Decomposition Reactions

This is opposite to that of Combination Reaction. In decompositions reactions, a single substance breaks into two or more simpler substances. At least one out of the formed substances has to be in the elemental form. 

Note: All Decomposition reactions are not redox reactions because in these reactions there is no change in the oxidation state of elements. 

c. Displacement Reactions 

In displacement reactions, one ion of a compound is replaced by an ion of the other element. Displacement reactions are of two types: 

a) Metal Displacement Reaction 

In the case of the metal displacement reaction, the metal of the compound is replaced by some other metal (in the form of elemental state). This reaction is more useful in metallurgy where certain metals are isolated from their ores using reduction mechanism 

b) Non - metal displacement reaction 

Non - metal displacement reaction mainly comprises of hydrogen displacement or oxygen replacement reactions.

i) Good reducing agents such as alkali metals and alkaline earth metals will displace hydrogen from the cold water. 

ii) Dihydrogen gas will be produced when less active metals such as Mg and Fe react with steam 

iii) Those metals that do not react with cold water displaces H₂ from acids. 

iv) Reactive metals that occur in their native state do not react even with dil. HCl such as Silver (Ag) and Gold(Au) etc. 

c) Reactivity of non - metals 

Reactivity of non - metals depends on their oxidising power. For halogens, oxidising powers decreases as we move down the group 17 from Fluorine on the top to I at the bottom. Since oxidising powers decreases, thus it can be said that Fluorine is the strongest oxidising agent.
Thus, fluorine can displace in their aqueous solutions, Cl₂ , Br₂, I₂ from Cl^- , Br^- and I^- ions.
Similarly, chlorine can displace bromide and iodide ions in aqueous solutions. Chlorine cannot displace fluoride ions.
Bromine can displace iodine from iodide ion in the aqueous solution.
These concepts are widely used in laboratory and industry. These reactions are used in getting halogens from their corresponding halides with suitable oxidising agents. There are no oxidising agents available to oxidise F^- ions to F₂ as fluorine itself is the strongest oxidising agent.

d. Disproportionation Reaction

Disproportionation reaction is a reaction in which the same element is oxidised as well as reduced simultaneously. For element to be involved in disproportionation reaction, the reacting species must have at least three oxidation states. This reaction happens only in the basic medium. 

In the above reaction, the oxidation number of oxygen in H₂O₂ which is -1 changes to -2 in H₂0 and changes to 0 in O₂

Halogen (except Fluorine) goes under disproportionation reaction. Fluorine does not go under disproportionation reaction as fluorine is the strongest oxidising agent and thus does not show positive oxidation states. Similarly, phosphorous, sulphur also goes under a disproportionation reaction.

5. Balancing of Redox Reactions

Two important methods through which a Redox Reaction can be balanced are: 

i) Oxidation number method 

ii) Ion-electron method 

i) Oxidation Number Method 

Following are the steps that should be used to balance the redox reactions using the oxidation number method 

• Write down the reactants and products of the reaction in its skeleton form. 
• Identify the elements whose oxidation number are changing and mention its oxidation states above the symbol 
• Calculate the increase or decrease int eh oxidation number per atom. Identify oxidising and reducing agents. 
  • In a case where more than one atom of the same element is involved, count the total decrease or increase in the oxidation number by multiplying this decrease or increase in oxidation number per atom by the number of atoms undergoing that change. 
• To equalise the total decrease or increase in the oxidation number, multiply the oxidising and reducing agents formulae by suitable integers. 
• The next step is to balance all atoms other than 'H' and 'O'
• Finally, to balance 'O', add H₂O molecules using hit and trial method. 
• Ionic Reactions case: 
  • Acidic Medium 
    
    • Balance oxygen atoms by adding water (H₂O) molecules to the side where the oxygen atom is in deficiency. 
    • Balance hydrogen atoms by adding H⁺ ions to the side where the hydrogen atom is in deficiency. 
  • Basic Medium
    • Balance oxygen atoms by adding water (H₂O) molecules to the side where the oxygen atom is in deficiency.
    • Balance hydrogen atoms by adding water (H₂O) molecules. The number of such water molecules shall be equal to the number of hydrogen atoms in deficiency. 
    • Equal number of hydroxide ions (OH^-) ions are added to the other side of the equations. 

    ii) Ion-electron method

Ion-electron method is based on the fact that the electrons lost during oxidation half-reaction are equal to the electrons gained during reduction half-reaction. Following are the steps that should be used to balance the redox reactions using the ion-electron method. 

• Write down the reactants and products of the reaction in its skeleton form.
•       Identify the elements whose oxidation number are changing and mention its oxidation states above the symbol 
• Write down two half-reactions simultaneously: Oxidation half-reaction and reduction half reaction
• Balance the two half-reactions separately following the below-mentioned steps: 
  • Balance the atoms of the elements whose oxidation number has changed 
  • To fill the difference in the oxidation number in each half reaction, add electrons to the required side. 
  • To fulfil the oxygen deficiency, add the required number of water molecules to the required side of the reaction. 
  • In the case of an acidic medium, balance H atoms by adding H⁺ ions to the side where the Hydrogen atom is in deficiency. 
  •       Balance oxygen atoms by adding water (H₂O) molecules to the side where the oxygen atom is in deficiency.
  •       Balance hydrogen atoms by adding water (H₂O) molecules. The number of such water molecules shall be equal to the number of                    hydrogen atoms in deficiency. 
  •       Equal number of hydroxide ions (OH^-) ions are added to the other side of the equations. 
  • Multiple the two half-reactions by a suitable integer to ensure that the total number of electrons gained in one-half of the reactions equals the number of electrons lost in the other half of the reaction. 
  • Add these two half-reactions to get the final balanced reaction. 

6. Titrations

Titration method is used to determine the strength of a reductant/oxidant using an indicator (redox-sensitive). Following are the indicators that are used in titrations: 

i) Potassium Permanganate (KMnO₄)

Potassium Permanganate (KMnO₄) is a strong oxidising agent which is used in the estimation of reducing agents. Potassium Permanganate (KMnO₄) is deep in colour, thus no indicator is used. It acts as a self-indicator. When Potassium Permanganate (KMnO₄) is titrated against a reducing agent, the colour of Potassium Permanganate (KMnO₄) disappears. When the reducing agent gets used up, a small addition of Potassium Permanganate (KMnO₄) will give a tinge of pink colour to the solution. 

ii) Acidified Potassium Dichromate (K₂Cr₂O₇) 

       Acidified Potassium Dichromate (K₂Cr₂O₇) is an oxidising agent used in the titration. It is less strong than that of Potassium Permanganate (KMnO₄).                   Unlike, Potassium Permanganate (KMnO₄), Potassium Dichromate (K₂Cr₂O₇) is not a self-indicator as no colour change takes place on its own. Thus,             indicators such as diphenylamine, potassium ferricyanide are used. 

iii) Iodine (I₂) 

Iodine (I₂) is a mild oxidising agent. It can be used for titrating strong reducing agents. 

7. Electrochemical Series 

• The oxidising agents are listed in the decreasing order of their strength i.e., as we move from top to bottom, the tendency of the oxidation half-reaction goes on decreasing.
  
  Application of Electrochemical Series
• The activity of the metal will depend on the standard reduction potential. Lower the standard reduction potential value, higher will be the activity of the metal because the tendency to lose electrons will be higher. 
• Metals having reduction potential lesser than hydrogen (having negative values), will have more tendency to lose electron than hydrogen. If this happens, then H⁺ ions will gain electrons and thus hydrogen can be liberated. This is the criteria for the liberation of the hydrogen.
• A redox reaction is feasible when the higher reduction potential species gains electrons and lower reduction potential release electrons. 
• EMF of the cell = Standard electrode potential of cathode - standard electrode potential of anode.  

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