CHSE Odisha Class 12 Chemistry Notes Chapter 3 Electrochemistry

Chapter 3: Electrochemistry

Introduction

Electrochemistry is the branch of chemistry that deals with the relationship between electrical energy and chemical changes. It involves the study of redox reactions where oxidation and reduction occur simultaneously, and the conversion of chemical energy into electrical energy or vice versa.

1. Electrochemical Cells

An electrochemical cell is a device that converts chemical energy into electrical energy or electrical energy into chemical energy. There are two main types of electrochemical cells:

• Galvanic (Voltaic) Cells: These generate electrical energy from spontaneous redox reactions.

• Electrolytic Cells: These use electrical energy to drive non-spontaneous chemical reactions.

1.1 Galvanic Cell (Voltaic Cell)

• Definition: A galvanic cell converts chemical energy into electrical energy using a spontaneous redox reaction.

• Example: Daniel cell.

• Components:

  • Anode (negative electrode): Oxidation occurs here.

  • Cathode (positive electrode): Reduction occurs here.

  • Salt Bridge: Maintains electrical neutrality by allowing the exchange of ions.

  Cell Representation (Daniel Cell):

  $$\text{Zn} (s) | \text{Zn}^{2+} (aq) || \text{Cu}^{2+} (aq) | \text{Cu} (s)$$

• Cell Potential (EMF of the Cell): It is the potential difference between the two electrodes and is given by:

  $$E^\circ_\text{cell} = E^\circ_\text{cathode} – E^\circ_\text{anode}$$

  ​

1.2 Electrolytic Cell

• Definition: Electrolytic cells use external electrical energy to drive a non-spontaneous reaction.

• Example: Electrolysis of molten NaCl.

• Key Points:

  • Anode is positive and Cathode is negative (opposite to galvanic cells).

  • Oxidation occurs at the anode, and reduction occurs at the cathode.

2. Electrochemical Series

The electrochemical series is a list of standard electrode potentials arranged in order of increasing reduction potential. The key points include:

• Elements at the top (like Li) are strong reducing agents (they have negative standard reduction potentials).

• Elements at the bottom (like F₂) are strong oxidizing agents (they have positive standard reduction potentials).

Applications:

• To predict the feasibility of redox reactions.

• To determine the standard EMF of a cell.

3. Nernst Equation

The Nernst equation relates the cell potential to the concentration of the reactants and products.

Nernst Equation for a general cell reaction:

$$E_\text{cell} = E^\circ_\text{cell} – \frac{0.0591}{n} \log \frac{[ \text{Products} ]}{[ \text{Reactants} ]}$$

​

Where:

• $E_\text{cell}$

  = Cell potential at non-standard conditions.

• $E^\circ_\text{cell}$

  = Standard cell potential.

• $n$

  = Number of moles of electrons transferred.

• $[ \text{Products} ] / [ \text{Reactants} ]$

  $[\text{Products}]/[\text{Reactants}]$ = Reaction quotient.

Applications of Nernst Equation:

• Calculation of EMF under non-standard conditions.

• Determination of equilibrium constants for redox reactions.

4. Conductance of Electrolytic Solutions

4.1 Electrolytic Conductance

The ability of an electrolyte solution to conduct electricity is termed as electrolytic conductance. It is influenced by factors such as concentration, temperature, and the nature of the electrolyte.

4.2 Specific Conductance (κ)

It is the conductance of a solution of unit length and unit cross-sectional area.

$$κ = \frac{1}{ρ}$$

​

Where

$κ$

is specific conductance and

$ρ$

ρ is resistivity.

4.3 Molar Conductance (Λ_m)

Molar conductance is defined as the conductance of all the ions produced by one mole of an electrolyte in a solution.

$$Λ_m = \frac{κ \times 1000}{C}$$

​

Where:

• $Λ_m$

  ​ = Molar conductance.

• $κ$

  = Specific conductance.

• $C$

  = Concentration of the solution in mol/L.

4.4 Variation of Conductance with Concentration

• Strong Electrolytes: Conductance decreases slightly with dilution.

• Weak Electrolytes: Conductance increases significantly with dilution due to increased ionization.

4.5 Kohlrausch’s Law of Independent Migration of Ions

It states that at infinite dilution, each ion contributes independently to the total molar conductance of an electrolyte.

$$Λ_m^\circ = λ^\circ_+ + λ^\circ_-$$

Where: $Λ_m^\circ$

${\mathrm{\Lambda }}_m^{\circ }$

= Limiting molar conductance.

$λ^\circ_+$

and

$λ^\circ_-$

are the limiting conductances of cations and anions, respectively.

5. Electrolysis and Faraday’s Laws

Electrolysis is a process where electrical energy is used to drive a non-spontaneous chemical reaction. This is carried out in an electrolytic cell, which consists of electrodes (anode and cathode) immersed in an electrolyte solution.

1. Electrolysis

1.1. Components of Electrolytic Cell

• Electrolyte: A substance that undergoes decomposition during electrolysis. It can be in molten form or dissolved in water (aqueous solution).

• Electrodes:

  • Anode (Positive electrode): Oxidation occurs here.

  • Cathode (Negative electrode): Reduction occurs here.

• External Power Source: Provides electrical energy to drive the non-spontaneous reaction.

1.2. Electrolytic Process

When electricity is passed through the electrolyte:

• At the Anode: Oxidation takes place, and electrons are released. $$\text{A}^{-} \rightarrow \text{A} + e^-$$

   

• At the Cathode: Reduction occurs, and electrons are accepted. $$\text{B}^{+} + e^- \rightarrow \text{B}$$

   

Example: Electrolysis of molten sodium chloride (NaCl):

• At Cathode: Na⁺ ions are reduced to sodium metal. $$Na^+ + e^- \rightarrow Na$$

   

• At Anode: Cl⁻ ions are oxidized to chlorine gas. $$2Cl^- \rightarrow Cl_2 + 2e^-$$

   

2. Faraday’s Laws of Electrolysis

Michael Faraday formulated two important laws that govern the process of electrolysis.

2.1. Faraday’s First Law of Electrolysis

This law states that:

• The amount of substance deposited or liberated at an electrode is directly proportional to the quantity of electricity (charge) passed through the electrolyte.

Mathematically:

$$m \propto Q$$

$$m = Z \cdot Q$$

 

Where:

• $m$

  = Mass of the substance deposited or liberated.

• $Q$

  = Quantity of charge passed (in coulombs).

• $Z$

  = Electrochemical equivalent (a constant for each substance).

Since $Q = I \cdot t$

, where $I$

$I$ is the current (in amperes) and $t$

is the time (in seconds):

$$m = Z \cdot I \cdot t$$

 

• Electrochemical equivalent $Z$

  is the mass of a substance deposited when one coulomb of charge is passed.

2.2. Faraday’s Second Law of Electrolysis

This law states that:

• When the same quantity of electricity is passed through different electrolytes, the masses of substances deposited or liberated are directly proportional to their equivalent weights.

Mathematically:

$$\frac{m_1}{m_2}=\frac{E_1}{E_{\mathrm{2<span\; style="font-family:\; -apple-system,\; system-ui,\; BlinkMacSystemFont,\; \text{'}Segoe\; UI\text{'},\; Helvetica,\; Arial,\; sans-serif,\; \text{'}Apple\; Color\; Emoji\text{'},\; \text{'}Segoe\; UI\; Emoji\text{'},\; \text{'}Segoe\; UI\; Symbol\text{'};\; font-size:\; inherit;"></span><span\; class="vlist-s"\; style="font-size:\; inherit;\; font-family:\; -apple-system,\; system-ui,\; BlinkMacSystemFont,\; \text{'}Segoe\; UI\text{'},\; Helvetica,\; Arial,\; sans-serif,\; \text{'}Apple\; Color\; Emoji\text{'},\; \text{'}Segoe\; UI\; Emoji\text{'},\; \text{'}Segoe\; UI\; Symbol\text{'};"></span>}}}$$

Where:

• $m_1$

  and $m_2$

  are the masses of substances deposited at two different electrodes.

• $E_1$

  ​ and $E_2$

  are the equivalent weights of the respective substances.

Equivalent weight ( $E$

$E$) of a substance is given by:

$$E=\frac{\text{Molar mass}}{\mathrm{n<span\; style="font-size:\; inherit;\; font-family:\; -apple-system,\; system-ui,\; BlinkMacSystemFont,\; \text{'}Segoe\; UI\text{'},\; Helvetica,\; Arial,\; sans-serif,\; \text{'}Apple\; Color\; Emoji\text{'},\; \text{'}Segoe\; UI\; Emoji\text{'},\; \text{'}Segoe\; UI\; Symbol\text{'};\; text-transform:\; none;"></span>}}$$

Where $n$

is the number of electrons involved in the redox reaction.

3. Faraday’s Constant

Faraday’s constant ( $F$

$F$) represents the total charge carried by one mole of electrons.

$$F=N_A\cdot e$$

Where:

1. $N_A = 6.022 \times 10^{23} \text{ mol}^{-1}$

   (Avogadro’s number).

2. $e = 1.602 \times 10^{-19} \text{ C}$

   (charge of one electron).

Hence,

$$F = 96500 \text{ C mol}^{-1}$$

 

This means that one mole of electrons carries a charge of 96500 C.

4. Applications of Electrolysis

• Electroplating: Electrolysis is used to deposit a thin layer of metal (e.g., gold, silver, chromium) onto another metal surface.

• Electrorefining: Used to purify metals (e.g., copper).

• Production of Chemicals: Electrolysis is used to produce substances like chlorine, sodium hydroxide (NaOH), and hydrogen gas.

• Metal Extraction: Electrolysis is used to extract metals from their ores (e.g., extraction of aluminum from alumina in the Hall-Héroult process).

Key Equations

1. Faraday’s First Law: $$m = Z \cdot I \cdot t$$

    

2. Quantity of charge: $$Q = I \cdot t$$

    

3. Faraday’s Second Law: $$\frac{m_1}{m_2} = \frac{E_1}{E_2}$$

   ​​

Summary

• Electrolysis is the process of using electrical energy to bring about a chemical reaction.

• Faraday’s First Law relates the amount of substance deposited to the total charge passed through the electrolyte.

• Faraday’s Second Law compares the masses of different substances deposited when the same amount of charge is passed.

• These laws form the foundation for various industrial applications like electroplating, refining, and chemical production.

5.3 Electrolysis of Aqueous Solutions

• Involves competing reactions at electrodes (e.g., electrolysis of water).

1. Conductance (G)

Conductance is the measure of how easily electricity flows through a material. The higher the conductance, the easier it is for current to pass through. It is the reciprocal (inverse) of resistance.

Formula for Conductance:

$$G = \frac{1}{R}$$

​

Where:

$G$

is the conductance (measured in siemens, S).

$R$

is the resistance (measured in ohms,

$Ω$

Ω).

Unit of Conductance:

• The SI unit of conductance is the siemens (S), formerly known as mho (ohm spelled backward).

Factors Affecting Conductance:

1. Nature of the material: Metals typically have high conductance, while insulators have low conductance.

2. Length of the conductor: Longer conductors have lower conductance.

3. Cross-sectional area: A larger cross-sectional area increases conductance.

4. Temperature: For metals, conductance decreases with an increase in temperature, while for electrolytes, conductance increases with temperature.

5. Concentration of electrolyte: In electrolytic solutions, the conductance depends on the concentration of ions.

2. Specific Conductance (κ)

Specific conductance (κ), also known as conductivity, is the conductance of a solution between two electrodes placed 1 cm apart with a unit cross-sectional area. It measures how well a solution can conduct electricity.

Formula for Specific Conductance:

$$κ = \frac{G \cdot l}{A}$$

​

Where:

$κ$

is the specific conductance or conductivity (measured in S/m or S/cm).

$G$

is the conductance of the solution.

$l$

is the length between the electrodes (in cm).

$A$

is the cross-sectional area of the electrodes (in cm²).

Unit of Specific Conductance:

• The SI unit is siemens per meter (S/m) or siemens per centimeter (S/cm).

Factors Affecting Specific Conductance:

1. Concentration of ions: Higher concentration of ions in the solution increases the conductivity.

2. Temperature: As temperature increases, conductivity typically increases because the mobility of ions increases.

3. Molar Conductance (Λ $_m$

m​)

Molar conductance is the conductance of all the ions produced by one mole of an electrolyte in a solution. It is calculated by dividing the specific conductance by the molar concentration of the electrolyte.

Formula for Molar Conductance:

$$Λ_m = \frac{κ \times 1000}{C}$$

​

Where:

$Λ_m$

​ is the molar conductance (measured in S cm²/mol).

$κ$

is the specific conductance (S/cm).

$C$

is the concentration of the solution in mol/L.

Unit of Molar Conductance:

• Siemens square centimeters per mole (S cm²/mol).

Variation with Concentration:

• For strong electrolytes (like NaCl), molar conductance decreases slightly with an increase in concentration due to interionic interactions.

• For weak electrolytes (like acetic acid), molar conductance increases significantly upon dilution because ionization increases.

4. Resistance (R)

Resistance is the measure of how much a material opposes the flow of electric current. Higher resistance means that the material does not allow electricity to pass easily.

Formula for Resistance:

$$R = \frac{V}{I}$$

​

Where:

$R$

is the resistance (measured in ohms,

$Ω$

Ω).

$V$

is the voltage across the conductor (in volts).

$I$

is the current through the conductor (in amperes).

Unit of Resistance:

• The SI unit of resistance is the ohm (Ω).

Factors Affecting Resistance:

1. Length of the conductor: Longer conductors have higher resistance.

2. Cross-sectional area: A larger cross-sectional area reduces resistance.

3. Material: Different materials have different resistivities. Metals like copper have low resistance, while insulators like rubber have high resistance.

4. Temperature: For most materials, resistance increases with an increase in temperature.

5. Relation Between Resistance and Conductance

Resistance and conductance are inversely related:

$$G = \frac{1}{R} \quad \text{or} \quad R = \frac{1}{G}$$

​

This means that a material with high resistance has low conductance and vice versa.

6. Resistivity (ρ) and Conductivity (κ)

Resistivity (ρ) is the resistance of a material of unit length and unit cross-sectional area. It is a material property that quantifies how strongly a material opposes the flow of electric current.

Formula for Resistivity:

$$R = ρ \cdot \frac{l}{A}$$
​

Where:

$R$

is the resistance.

$ρ$

is the resistivity (measured in ohm-meter,

$Ω \cdot m$

Ω⋅m).

$l$

is the length of the conductor (in meters).

$A$

is the cross-sectional area (in square meters).

Formula for Conductivity:

Conductivity (

$κ$

κ) is the inverse of resistivity:

$$\kappa =\frac{1}{\mathrm{\rho <span\; style="font-size:\; inherit;\; font-family:\; -apple-system,\; system-ui,\; BlinkMacSystemFont,\; \text{'}Segoe\; UI\text{'},\; Helvetica,\; Arial,\; sans-serif,\; \text{'}Apple\; Color\; Emoji\text{'},\; \text{'}Segoe\; UI\; Emoji\text{'},\; \text{'}Segoe\; UI\; Symbol\text{'};\; text-transform:\; none;"></span>}}$$

Where

$κ$

is the specific conductance (conductivity) and

$ρ$

is the resistivity.

---

Summary of Key Concepts

Concept	Definition	Unit	Formula
Conductance (G)	Measure of how easily current flows through a material	Siemens (S)	$G = \frac{1}{R}$ ​
Resistance (R)	Measure of how much a material opposes current flow	Ohms (Ω)	$R = \frac{V}{I}$ ​
Specific Conductance	Conductance of a solution between electrodes 1 cm apart with unit area	Siemens per meter (S/m)	$κ = \frac{G \cdot l}{A}$ ​
Molar Conductance	Conductance of all ions produced by one mole of an electrolyte	S cm²/mol	$Λ_m = \frac{κ \times 1000}{C}$ ​​
Resistivity (ρ)	Resistance of a material of unit length and cross-sectional area	Ohm-meter (Ω⋅m)	$R = ρ \cdot \frac{l}{A}$ ​
Conductivity (κ)	Inverse of resistivity, indicating how well a material conducts electricity	Siemens per meter (S/m)	$κ = \frac{1}{ρ}$ ​

6. Batteries

6.1 Primary Cells

• Non-rechargeable cells.

• Example: Dry cell (Leclanché cell).

6.2 Secondary Cells

• Rechargeable cells.

• Example: Lead-acid battery, nickel-cadmium battery.

7. Fuel Cells

Fuel cells convert the chemical energy of a fuel directly into electrical energy. A common example is the hydrogen-oxygen fuel cell.

Reaction:

$$\text{Anode: } 2H_2 (g) \rightarrow 4H^+ (aq) + 4e^-$$

$$\text{Cathode: } O_2 (g) + 4H^+ (aq) + 4e^- \rightarrow 2H_2O (l)$$

Total reaction:

$$2H_2 (g) + O_2 (g) \rightarrow 2H_2O (l)$$

8. Corrosion

Corrosion is the degradation of metals as a result of electrochemical reactions with their environment. The most common example is rusting of iron.

Prevention of Corrosion:

• Coating the metal surface (painting, galvanization).

• Using sacrificial anodes.

• Cathodic protection.

Corrosion

Corrosion is the gradual destruction of metals due to chemical or electrochemical reactions with the environment. The most common type of corrosion is the rusting of iron, where iron reacts with oxygen and moisture to form iron oxides.

1. Rusting of Iron

• Rust is a red-brown flaky substance that forms on the surface of iron or steel when it is exposed to air and moisture for a prolonged period.

• The overall reaction involved in rusting: $$4Fe(s) + 3O_2(g) + 6H_2O(l) \rightarrow 4Fe(OH)_3(s)$$
  4Fe(s)+3O2​(g)+6H2​O(l)→4Fe(OH)3​(s)

  • $Fe(OH)_3$

    ​ (iron hydroxide) further dehydrates to form rust ( $Fe_2O_3 \cdot xH_2O$

    ).

Electrochemical Mechanism of Corrosion: Corrosion of metals like iron occurs via electrochemical reactions, where iron acts as an anode and undergoes oxidation:

• Anodic Reaction (Oxidation): $$Fe \rightarrow Fe^{2+} + 2e^-$$

   

• Cathodic Reaction (Reduction): $$O_2 + 4H^+ + 4e^- \rightarrow 2H_2O$$

   

In the presence of moisture, the

$Fe^{2+}$

Fe2+ ions further react with oxygen to form rust.

Prevention of Corrosion

There are several methods to prevent or slow down the corrosion of metals:

1. Barrier Protection

The metal surface is coated with a barrier that prevents air, moisture, and electrolytes from reaching the metal.

• Painting: Coating the metal surface with paint creates a physical barrier.

• Coating with Oil/Grease: Used in machinery to avoid exposure to moisture.

• Electroplating: The metal surface is coated with another metal (e.g., chrome or nickel) that resists corrosion.

2. Galvanization

Galvanization is the process of coating iron or steel with a thin layer of zinc. Zinc is more reactive than iron and corrodes preferentially, protecting the iron underneath. This protection works in two ways:

• Barrier protection: The zinc layer prevents air and moisture from reaching the iron surface.

• Sacrificial protection: Zinc acts as a sacrificial anode, corroding instead of iron, even if the zinc coating is scratched.

3. Cathodic Protection

In this method, the metal to be protected is made the cathode of an electrochemical cell. It is achieved by connecting the metal to a more easily corroded metal, called a sacrificial anode.

• Sacrificial Anodes: Metals like zinc or magnesium are used as sacrificial anodes. These metals corrode instead of the protected metal (iron or steel).

• Impressed Current: An external current is applied to the metal to negate the corrosion potential.

4. Alloying

Corrosion resistance can be increased by mixing metals to form alloys. Some alloys are more resistant to corrosion than pure metals.

• Example: Stainless Steel (iron alloyed with chromium and nickel) is highly resistant to corrosion because the chromium forms a protective oxide layer on the surface.

5. Using Inhibitors

Corrosion inhibitors are chemicals that slow down or prevent corrosion. They are added to the environment (e.g., water, oils, or acids) to reduce the rate of corrosion.

• Example: Chromates and phosphates are used as inhibitors in cooling systems.

Summary of Corrosion Prevention Methods

Method	How It Works	Example
Barrier Protection	Blocks air and moisture	Painting, oil/grease coating
Galvanization	Zinc layer prevents corrosion by acting as a barrier and sacrificial anode	Coating steel with zinc
Cathodic Protection	Metal is made the cathode, using sacrificial anodes or impressed current	Zinc or magnesium anodes for ships
Alloying	Adds elements to improve corrosion resistance	Stainless steel
Using Inhibitors	Chemicals reduce or slow down the corrosion process	Adding inhibitors to cooling systems

Important Formulas

1. Cell potential: $$E^\circ_\text{cell} = E^\circ_\text{cathode} – E^\circ_\text{anode}$$

   ​

2. Nernst equation: $$E_\text{cell} = E^\circ_\text{cell} – \frac{0.0591}{n} \log \frac{[ \text{Products} ]}{[ \text{Reactants} ]}$$

   ​

3. Molar conductance: $$Λ_m = \frac{κ \times 1000}{C}$$

   ​