Balancing of Redox Reactions

Redox reactions, as they are known, are energy-producing processes that occur in both industry and the human body. The transfer of one or more electrons from one species to another is at the heart of a redox reaction. The species that loses electrons is called oxidized, whereas the species that gains electrons is called reduced. The importance of oxidation-reduction (redox) processes is that they are the primary sources of energy on our planet, both natural and manufactured. Large amounts of energy are generally released when molecules are oxidized by removing hydrogen or combining with oxygen.

Details on Examples, Types, Applications, Balancing of Redox Reactions-

Now we will take a look at what do we mean by redox reactions, balancing redox reactions, types of redox reactions and lastly at applications of redox reactions. 

Redox reactions-

An oxidation-reduction reaction is any chemical process in which the oxidation number of a participating chemical species changes. The term encompasses a wide variety of activities. Fire, metal corrosion and dissolving, fruit browning, and basic life functions like breathing and photosynthesis are all examples of oxidation-reduction reactions.

The transfer of oxygen, hydrogen, or electrons is involved in the majority of oxidation-reduction (redox) reactions, and all three processes have two fundamental characteristics:

(1) Every oxidation process involves a reciprocal reduction.

(2) They include a typical net chemical change, in which an atom or electron is moved from one unit of matter to another.

Balancing redox reaction types of redox reactions-

When a chemical reaction happens, the mass of the products should match the mass of the reactants, according to the rule of conservation of mass. As a result, the number of atoms in each element does not change as a result of the chemical reaction. As a consequence, the chemical equation describing the reaction must be balanced. When the number of atoms in the reactants side equals the number of atoms in the products side, the chemical equation is balanced.

There are 2 methods-

1. Oxidation number method 

2. Half reaction method 

We have to determine the oxidation numbers of all atoms using the oxidation number approach. The atoms that have altered are then multiplied by tiny whole numbers. Making the entire electron loss equal to the total electron gain. The remaining atoms are then balanced. Determine the oxidation numbers and write two half-reactions using the half-reaction approach. Then you multiply them by tiny whole numbers to equalize the electron loss and gain. Then you combine the two partial reactions and balance the remaining atoms.

We must know what the oxidation numbers are and what they become throughout the reaction in both approaches. Occasionally, one way is more convenient than the other.

1.Oxidation number method-

We will see how to balance Fe2O3(s)+CO(g)→Fe(s)+CO2(g)  

Step 1- Give each of the atoms in the equation an oxidation number and write it above the atom.

Fe+32O-23(s)+C+2O−2(g)→Fe(s)+C+4O2−2(g)

Step 2-Determine which atoms are oxidized and which are reduced. The carbon atom is oxidized in the previous equation because its oxidation rises from +2 to +4. Since its oxidation number drops from +3 to 0, the iron atom is being reduced.

Step 3- Draw a line connecting the atoms that are changing oxidation numbers. Write the oxidation-number change on the line.The oxidation number of the carbon atom grows by 2, while the oxidation number of the iron atom lowers by 3. In a redox process, the oxidation number of the carbon atom increases by 2, while the oxidation number of the iron atom decreases by 3.

The number of electrons lost does not match the number of electrons gained in the equation as presented. These must be equivalent in a balanced redox equation. As a result, the rise in one atom’s oxidation number must be balanced by the reduction in the oxidation number of the other.

Step 4- Using coefficients, make the total oxidation number rise equal to the total oxidation number decrease. The least common multiple of 2 and 3 in this example is 6. As a result, the rise in oxidation number should be multiplied by 3, while the reduction in oxidation number should be multiplied by 2. The coefficient is also applied to the equation’s formulations. 

As a result, a 3 is inserted in front of the CO and the CO2. On the right side of the equation, a 2 is inserted in front of the Fe. Because the subscript of 2 following the Fe indicates that there are already two iron atoms, the Fe2O3 does not need a coefficient.

Step 5- Verify that both the atoms and the charge are balanced. A coefficient may be needed in front of a molecular formula that was not engaged in the redox process on occasion. The equation is now balanced in the present case.

Fe2O3(s)+3CO(g)→2Fe(s)+3CO2(g)

2.Half reactions method-

Step 1- Make a list of the imbalanced equations.

Step 2- Split the redox process into two halves.

a) Assign oxidation numbers to each atom

b) Identify and write down all redox couples in the reaction.

c) To make two half-reactions, combine these redox couples.

Step 3- In each half reaction, balance the atoms.

a) Balance all atoms except H and O

b) Balance oxygen atoms with H2O 

c) Balance hydrogen atoms with H+

d) For every H+, add one OH- to either side of a basic medium.

Step 4: Use e-juice to balance the charge.

Step 5: In the half-reactions, make electron gain equal to electron loss.

Step 6: Combine the half-reactions.

Step 7: Make the equation easier to understand.

Co3+ (aq) + Ni(s) → Co2+(aq)+Ni2+(aq)

Will be balanced as Ni(s)+2Co3+(aq)→Ni2+(aq)+2Co2+(aq) 

Types of redox reactions-

1.Combination reaction- A combination reaction is a chemical reaction in which two or more chemicals combine to form a single new product. The reaction is called a redox reaction if one (or both) of the reactants are in the elemental state.

This kind of reaction is known as a redox reaction because both species (A & B) or either of the species (A & B) are present in their elemental form. This kind of redox reaction is the polar opposite of the breakdown process (oxidation and reduction).

2Na(s)+Cl2(g)→2NaCl(s)

2.Decompostion-When a complex is broken down into two or more simple components, the decomposition reaction occurs. The decomposition reaction is the polar opposite of the combination process. Redox reactions aren’t always present in decomposition processes.

A redox reaction occurs when at least one of the breakdown process’s products is in the elemental state.

2 H2O2→ 2 H2O + O

3.Displacement reaction- An atom (or ion) from one element or compound is replaced by an atom (or ion) from another element or compound.

Mg(s)+CuSO4(aq)+Mg(s)+CuSO4(aq)+Mg(s)+Cu

Cu+MgSO4 (aq) (s) 

4.Disproportionation reaction- This is a reaction in which the same element is oxidized and reduced. Such reactions occur when one of the reactive elements has at least three oxidation states.

In addition, the reactive molecule’s element must be in the intermediate oxidation state, while oxidation and reduction need the lower and higher oxidation states, respectively. A common example of such a process is the breakdown of hydrogen peroxide, in which oxygen species are disproportionated.

Mn2O3 degrades to Mn2+ and MnO2.

Applications of redox reactions-

1.Metal extraction- Using an appropriate reducing agent, metal oxides may be converted to metals. In the blast furnace, for example, ferrous oxide is reduced to iron using coke as a reducing agent.

2.Chemical manufacturing- Electrolysis, which is based on redox reactions, produces several chemicals that we utilize in our everyday lives, such as caustic soda, chlorine, and fluorine.

3.Analysis quantitative- Redox reactions are highly helpful in redox titrations for quantitative analysis. The reactions between oxidizing and reducing agents are used in these titrations to estimate the quantity of unknown compounds in solutions. In the pharmaceutical industry, these responses are helpful.

Conclusion-

In this article we read about redox reactions, different methods to balance such equations, types of redox reactions and lastly applications of these reactions. The importance of oxidation-reduction (redox) processes is that they are the primary sources of energy on our planet, both natural and manufactured. Large amounts of energy are generally released when molecules are oxidized by removing hydrogen or combining with oxygen. The species that gets electrons is reduced and is generally the oxidizing agent, whereas the species that loses electrons is oxidized and is usually the reducing agent. Photosynthesis, respiration, combustion, and corrosion are examples of common redox processes.