Redox Reactions and Electrochemistry: Comprehensive NEET Chemistry Notes

1. Redox Reactions

1.1 Classical Idea of Redox Reactions – Oxidation and Reduction Reactions

Redox reactions involve the simultaneous occurrence of oxidation and reduction. Traditionally, oxidation was defined as the addition of oxygen to a substance or the removal of hydrogen. Reduction was the removal of oxygen or the addition of hydrogen.

Key Reactions:

• $2\text{Mg (s)}+{\text{O}}_2\text{(g)}\to 2\text{MgO (s)}$
• ${\text{CH}}_4\text{(g)}+2{\text{O}}_2\text{(g)}\to {\text{CO}}_2\text{(g)}+2{\text{H}}_2\text{O (l)}$

Oxidation Definition Expansion:

• Removal of hydrogen or electropositive elements (e.g., potassium from potassium ferrocyanide) can also be considered oxidation.

Reduction Definition Expansion:

• The process involves the removal of oxygen/electronegative elements or the addition of hydrogen/electropositive elements.

1.2 Redox Reactions in Terms of Electron Transfer

Electron transfer is fundamental in redox reactions. When sodium reacts with chlorine, sodium loses an electron (oxidized), and chlorine gains an electron (reduced).

Example:

• $2\text{Na (s)}\to 2{\text{Na}}^{+}\text{(g)}+2{\text{e}}^{-}$
• ${\text{Cl}}_2\text{(g)}+2{\text{e}}^{-}\to 2{\text{Cl}}^{-}\text{(g)}$

Mnemonic:

LEO (Loss of Electrons is Oxidation) says GER (Gain of Electrons is Reduction).

1.3 Oxidation Number Concept

The oxidation number helps in tracking the electron shifts in reactions involving covalent compounds.

Rules for Assigning Oxidation Numbers:

1. The oxidation number of an element in its elemental form is zero.
2. For monoatomic ions, the oxidation number equals the charge on the ion.
3. Oxygen usually has an oxidation number of -2, except in peroxides and superoxides.
4. Hydrogen has an oxidation number of +1, except when bonded to metals in binary compounds.

Example:

• ${\text{H}}_2{\text{O}}_2\text{(aq)}\to 2{\text{H}}_2\text{O (l)}+{\text{O}}_2\text{(g)}$ (Oxygen goes from -1 in ${\text{H}}_2{\text{O}}_2$ to 0 in ${\text{O}}_2$ and -2 in ${\text{H}}_2\text{O}$)

1.4 Types of Redox Reactions

1. Combination Reactions:
2. • Involves two or more substances combining to form a single product.
   • Example: $\text{C (s)}+{\text{O}}_2\text{(g)}\to {\text{CO}}_2\text{(g)}$
2. Decomposition Reactions:
3. • A compound breaks down into two or more components.
   • Example: $2{\text{H}}_2\text{O (l)}\to 2{\text{H}}_2\text{(g)}+{\text{O}}_2\text{(g)}$
3. Displacement Reactions:
4. • An ion in a compound is replaced by an ion or atom of another element.
   • Example: $\text{Zn (s)}+{\text{CuSO}}_4\text{(aq)}\to {\text{ZnSO}}_4\text{(aq)}+\text{Cu (s)}$

Did You Know?

The electron transfer principle is used in designing Galvanic cells, where chemical reactions produce electrical energy.

Quick Recap

• Redox reactions involve simultaneous oxidation and reduction.
• Oxidation is the loss of electrons, and reduction is the gain of electrons.
• Oxidation number helps track electron shifts in reactions.

2. Electrochemistry

2.1 Electrochemical Cells

Electrochemical cells convert chemical energy into electrical energy or vice versa.

Key Types of Cells:

• Galvanic Cells (Voltaic Cells): Generate electrical energy from spontaneous redox reactions.
• Electrolytic Cells: Use electrical energy to drive non-spontaneous chemical reactions.

Example of a Galvanic Cell:

• Daniell cell, where zinc is oxidized at the anode, and copper is reduced at the cathode.

Equation:

• Anode: $\text{Zn (s)}\to {\text{Zn}}^{2+}\text{(aq)}+2{\text{e}}^{-}$
• Cathode: ${\text{Cu}}^{2+}\text{(aq)}+2{\text{e}}^{-}\to \text{Cu (s)}$

Real-life Application:

Galvanic cells are the basis of batteries used in everyday devices like remote controls and flashlights.

2.2 Standard Electrode Potential (E°)

The standard electrode potential is a measure of the tendency of a redox couple to gain or lose electrons.

Standard Hydrogen Electrode (SHE):

• Used as a reference with a potential of 0.00 V.

Important Electrode Potentials:

• ${\text{F}}_2/{\text{F}}^{-}:+2.87\text{V}$ (strong oxidizing agent)
• ${\text{Li}}^{+}/\text{Li}:-3.04\text{V}$ (strong reducing agent)

2.3 Nernst Equation

The Nernst equation relates the cell potential to the standard electrode potential, temperature, and reaction quotient.

Formula: $E_{\text{cell}}=E_{\text{cell}}^{\circ }-\frac{RT}{nF}\ln Q$

Example Application:

The Nernst equation is used to calculate the cell potential under non-standard conditions.

Common Misconception:

Many students confuse the anode and cathode in electrochemical cells. Remember, oxidation always occurs at the anode, and reduction occurs at the cathode.

Quick Recap

• Electrochemical cells convert chemical energy to electrical energy.
• Galvanic cells are spontaneous, while electrolytic cells require energy input.
• Standard electrode potential indicates the tendency to gain or lose electrons.

Practice Questions

1. Balance the following redox reaction in acidic medium: ${\text{MnO}}_4^{-}\text{(aq)}+{\text{Fe}}^{2+}\text{(aq)}\to {\text{Mn}}^{2+}\text{(aq)}+{\text{Fe}}^{3+}\text{(aq)}$
2. Calculate the cell potential for a Galvanic cell using the following half-reactions:
3. • ${\text{Zn}}^{2+}\text{(aq)}+2{\text{e}}^{-}\to \text{Zn (s)}{E}^{\circ }=-0.76\text{V}$
   • ${\text{Cu}}^{2+}\text{(aq)}+2{\text{e}}^{-}\to \text{Cu (s)}{E}^{\circ }=+0.34\text{V}$
3. Explain the role of the salt bridge in a Galvanic cell.
4. Determine the oxidation number of chromium in ${\text{K}}_2{\text{Cr}}_2{\text{O}}_7$.
5. What will be the effect on the cell potential if the concentration of ${\text{Zn}}^{2+}$ is increased in a Daniell cell?

Solutions:

1. Balanced equation: ${\text{MnO}}_4^{-}\text{(aq)}+5{\text{Fe}}^{2+}\text{(aq)}+8{\text{H}}^{+}\text{(aq)}\to {\text{Mn}}^{2+}\text{(aq)}+5{\text{Fe}}^{3+}\text{(aq)}+4{\text{H}}_2\text{O (l)}$
2. $E_{\text{cell}}=E_{\text{cathode}}^{\circ }-E_{\text{anode}}^{\circ }=0.34\text{V}-(-0.76\text{V})=1.10\text{V}$
3. The salt bridge maintains electrical neutrality by allowing ions to flow between the two solutions, preventing charge buildup.
4. Chromium in ${\text{K}}_2{\text{Cr}}_2{\text{O}}_7$ has an oxidation number of +6.
5. Increasing the concentration of ${\text{Zn}}^{2+}$ will decrease the cell potential according to the Nernst equation.

Glossary

• Redox Reaction: A chemical reaction involving the transfer of electrons between two species.
• Electrochemical Cell: A device that converts chemical energy into electrical energy or vice versa.
• Standard Electrode Potential (E°): The potential of a half-cell under standard conditions (298 K, 1 atm, 1 M solutions).
• Nernst Equation: An equation used to calculate the potential of an electrochemical cell under non-standard conditions.

This guide provides a structured approach to mastering the key concepts of Redox Reactions and Electrochemistry, which are critical for NEET Chemistry preparation.