Redox Reactions

Redox reactions are chemical reactions where one substance loses electrons (oxidation) and another gains electrons (reduction) at the same time. These reactions are important in many natural and industrial processes, including energy production, corrosion, and metabolism. They can be classified into different types, such as combination, decomposition, displacement, and disproportionation reactions. 

Redox reactions also play a key role in electrochemistry, which studies how chemical reactions can produce electricity (like in batteries) and how electricity can drive chemical changes. Understanding redox reactions helps in balancing chemical equations and is essential in fields like chemistry, biology, and environmental science.

What are Redox Reactions?

• A redox reaction is a chemical reaction where electrons are transferred between two substances.

• The oxidation state of the reacting substances changes during the reaction.

• Oxidation happens when a substance loses electrons and its oxidation state increases.

• Reduction happens when a substance gains electrons and its oxidation state decreases.

• The substance that accepts electrons is called an oxidising agent (it gets reduced).

• The substance that donates electrons is called a reducing agent (it gets oxidised).

Rules for Assigning Oxidation States

Oxidation state (OS) is a number that tells us how many electrons an atom has gained, lost, or appears to use when forming a compound. There are some basic rules to figure out the oxidation state of an atom-

• A free element (not in a compound) always has an oxidation state of 0.

• In a neutral molecule, the total oxidation state of all atoms is 0, while in an ion, it equals the charge of the ion.

• Group 1 metals (like sodium and potassium) always have an oxidation state of +1, and Group 2 metals (like calcium and magnesium) always have +2.

• Fluorine always has an oxidation state of -1 in its compounds.

• Hydrogen usually has an oxidation state of +1 in compounds.

• Oxygen generally has an oxidation state of -2 in compounds.

• Elements in Group 17 (halogens) usually have -1, Group 16 (like oxygen and sulfur) usually have -2, and Group 15 (like nitrogen and phosphorus) usually have -3 in binary metal compounds.

Types of Redox Reactions

There are four different types of Redox Reactions. Namely,

• Decomposition Reactions

• Combination Reactions

• Displacement Reactions

• Disproportionation Reactions

Decomposition Reaction-

• The breakdown of a molecule into other compounds is what this reaction entails. Various are some examples of these types of reactions-

• All of the aforementioned reactions lead to the breakdown of smaller chemical compounds in the form of AB →A + B.

• However, there is one exception that proves that not all decomposition reactions are redox reactions.

Combination Reaction-

• These reactions are the inverse of decomposition processes, in that they combine two chemicals to generate a single compound with the formula A + B → AB.

Displacement Reaction-

• An atom or an ion of a compound is replaced by an atom or an ion of another element in this reaction. 

• It can be represented as X + YZ → XZ + Y.

• Displacement reactions can also be divided into two types.

• The displacement of metal Reaction

• Displacement of non-metals Reaction

Metal Displacement-

• A metal present in the compound is displaced by another metal in this reaction. 

• In metallurgical procedures, when pure metals are extracted from their ores, several types of reactions are used.

Non-Metal Displacement-

• We can detect a hydrogen displacement reaction in this type of reaction, as well as rare oxygen displacement events.

Disproportionation Reaction-

• Disproportionation reactions are those that involve only one reactant being oxidised and reduced.

Oxidation Reaction

• The addition of oxygen or the more electronegative element to a compound or the removal of hydrogen or the more electropositive element from a substance is called an oxidation reaction, according to one definition.

• Some instances of oxidation processes are as follows-

• 2S(s) + O2(g) → SO2(g)

• CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)

Reduction Reaction

• Reduction reactions, like oxidation events, are defined as electron gains. 

• During a chemical process, any material that acquires an electron is reduced.

• The reduction reaction is defined as the addition of hydrogen or a more electropositive element to a substance, or the removal of a more electronegative element or oxygen.

• Now, if we look closely at the above reactions, we can see that they all have both reduction and oxidation reactions.

• As the electronegative element chlorine is removed from FeCl3, it undergoes a reduction process. 

• In the same reaction, hydrogen is oxidised due to the presence of chlorine, an electronegative element.

Oxidising and Reducing Agents

• The oxidising agent is a material (atom, ion, or molecule) that acquires electrons and is therefore reduced to a low valency state.

• A reducing agent is a chemical that loses electrons and hence oxidises to a higher valency state.

Oxidising Agent

• Any element or component that takes up electrons from the other compound and reduces itself is an oxidising agent.

• Electronegative components that make up molecules are strong oxidising agents. For example, O2, O3, and X2 (halogens).

• Compounds that include an element in a more oxidised state. For example, KMnO4, K2Cv2O7, HNO3, and KClO3

• Metal and non-metal oxyhydroxides MgO, CuO, CrO3, P4O10, etc.

• Fluorine is the most powerful oxidiser.

Reducing Agent

• A reducing agent is any atom or component that gives electrons to another element and oxidases itself.

• All metals, such as Na, Zn, Fe, and Al.

• A few non-metals are carbon, hydrogen, and sulphur.

• In the presence of water, lithium is the most powerful reducing agent, but in the absence of water, Cesium is the most powerful reducing agent. 

• H2O2, SO2, H2SO3, HNO2, NaNO2, are oxidising as well as reducing agents.

Solved Examples from the Chapter

Question 1- In an acidic media, H2O2 converts Cr2O72- ion to CrO5, and the oxidation state of Cr in CrO5 is

(a) +6

(b) +5

(c) +3

(d) -10

Solution- (a)

• Calculate the oxidation numbers and compare them.

• An increase in the oxidation number represents oxidation.

• A drop in the oxidation number represents reduction.

• In this case, Cr is undergoing oxidation or reduction.
  Cr2O72- → CrO5

• Let the oxidation/reduction number of Cr be x
  ∴ In Cr2O72-, 2x + 7(-2) = -2
  Hence, x = -2 + 14 = +12/2 = +6.

• Similarly, in CrO5, x + 5(-2) = -4.
  Hence, x = +6.

• Therefore, the oxidation number of Cr in CrO5 is +6.

Key points to remember- Accurate calculation of the changes in the oxidation and reduction number after having identified the oxidising agent and the reducing agent respectively, is the most important concept while solving these types of problems.

Question 2- The oxidation number of Fe is +1 in which of the following complexes? 

(a) [FeBr4]–

(b) [Fe(H2O)5NO]SO4

(c) Fe4[Fe(CN)6]3

(d) [Fe(H2O)6]2-

Solution-

• Step 1- Let the oxidation number of Fe in each compound be x. 

• Step 2- Find out the value of x in each compound

• Option A- [FeBr4]–
  x + 4(Br) = -1;
  Now, Br stands for Bromine. And bromine can gain one electron and have a negative charge of -1, because it is a halogen.
  ∴ x + 4(-1) = -1,
  ∴ x = -1 + 4 = 3.
  Thus, the oxidation number of Fe in  [FeBr4]–  is +3.

• Option B- [Fe(H2O)5NO]SO4
  In this compound there are 5 water molecules. Since water in itself does not contain any charge, the oxidation number becomes 0.
  On the other hand the NO compound can have an oxidation state of +1.
  Similarly, the sulphate ion on its own carries a charge of -2, hence its oxidation number is -2. Hence, the charge on the Fe containing complex is +2.
  ∴ x + 5(0) + (+1) = +2,
  ∴ x =  = +1.
  Thus, the oxidation number of Fe in  [FeBr4]–  is +1.

• Similarly, the oxidation number of Fe in Fe4[Fe(CN)6]3 is +3 and +2 for the ion outside the coordination complex and inside the coordination complex respectively. And the oxidation number of Fe in [Fe(H2O)6]2- is -2.

• Hence, the final answer is (b) [Fe(H2O)5NO]SO4.

Key points to remember- Accurate calculation of the changes in the oxidation and reduction number after having identified the oxidising agent and the reducing agent respectively, is the most important concept while solving these types of problems. It might so happen that for certain chemical species the oxidation state will remain unchanged. Care should be taken while handling such chemical species.

Solved Examples of Previous Year Question Papers

Question 1- Given - XNa2HAsO3 +YNaBrO3+ZHCl → NaBr + H3AsO4 + NaCl

The values of X, Y and Z in the above redox reaction are respectively -

(1) 2, 1, 3

(2) 3, 1, 6

(3) 2, 1, 2

(4) 3, 1, 4

Solution- 

• The equation for a balanced equation is shown below-

3Na2HAsO3 + NaBrO3 + 6HCl → NaBr + 3H3AsO4 + 6NaCl

• X, Y, and Z have the values 3, 1, and 6 correspondingly.

• As a result, option (2) is the correct answer.

Question 2- Consider the reaction

H2SO3(aq) + Sn4+(aq) + H2O(l) → Sn2+(aq) + HSO4–(aq) + 3H+(aq)

Which of the following statements is correct?

(1) H2SO3 is the reducing agent because it undergoes oxidation

(2) H2SO3 is the reducing agent because it undergoes reduction

(3) Sn4+ is the reducing agent because it undergoes oxidation

(4) Sn4+ is the oxidising agent because it undergoes oxidation

Solution- 

• The loss of electrons by a molecule during a reaction is referred to as oxidation. 

• Because it undergoes oxidation, H2SO3 is the reducing agent in the above equation.

• As a result, option 1 is the correct answer.

Question 3- In which of the following reactions H2O2 acts as a reducing agent ?

(1) H2O2 + 2H+ + 2e– → 2H2O

(2) H2O2 - 2e– → O2 + 2H+

(3) H2O2 + 2e– → 2OH–

(4) H2O2 + 2OH– - 2e– → O2 + 2H2O

(1) (1), (3)

(2) (2), (4)

(3) (1), (2)

(4) (3), (4)

Solution- 

• In a redox chemical process, a reducing agent is an element or molecule that loses an electron to an electron recipient. 

• H2O2 functions as a reducing agent in (2) and (4).

• As a result, option (2) is the correct answer.

Practice Questions

Question 1- Which of the following chemicals has the highest reduction potential?

(a) H2S

(b) HNO2

(c) SnCl2

(d) H2SO3

Answer- (a) H2S

Question 2- Which element's oxidation number changes the most during the reaction of oxalic acid, potassium chlorate, and sulphuric acid?

(a) H

(b) S

(c) C

(d) Cl

Answer- (d) Cl

Conclusion

The oxidation states of the reactants change in redox reactions, which are oxidation-reduction chemical reactions. The reduction-oxidation process is referred to as redox. Redox reactions can be broken down into two categories- reduction and oxidation.

The oxidation and reduction reactions always happen at the same time in a redox reaction, also known as an Oxidation-Reduction process. In a chemical reaction, the oxidising agent is the substance that is being reduced, while the reducing agent is the substance that is being oxidised.

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