2. Electrochemistry

Introduction to Electrochemistry

Electrochemistry studies the relationship between chemical energy and electrical energy and how one converts into the other. It focuses on reactions where electrons are transferred, frequently involving redox reactions.

Types of Conductors

• Metallic Conductors: Conduct electricity via electron flow; do not undergo chemical change (e.g., metals).

• Electrolytic Conductors (Electrolytes): Conduct through ion movement in molten or aqueous states and undergo chemical decomposition; can be strong (fully dissociate) or weak (partially dissociate) electrolytes

Electrochemical Cells

1. Galvanic (Voltaic) Cells

• Convert chemical energy to electrical.

• Consist of two half-cells with different metal electrodes in their own ionic solution, connected by a salt bridge

  • Anode: Oxidation occurs; electrons are released.

  • Cathode: Reduction occurs; electrons are gained.

2. Electrolytic Cells

• Convert electrical energy to chemical.

• Electrodes immersed in electrolyte; electricity forces a non-spontaneous reaction (e.g., electrolysis of NaCl)

Salt Bridge Role

The salt bridge connects the two half-cells, maintaining electrical neutrality by allowing ion flow while preventing the mixing of solutions.

In a Daniell cell, a salt bridge is a critical component that connects the two half-cells (anode and cathode) containing the electrolytic solutions. It typically consists of a U-shaped tube filled with a concentrated solution of an inert electrolyte like $\text{KCl}$ or $\text{KNO}_3$ in agar-agar gel.

The salt bridge serves two primary purposes to keep the cell functioning:

1. Maintaining Electrical Neutrality

As the redox reaction progresses, ions are produced or consumed in the half-cells:

• At the Anode: Zinc atoms oxidize into $\text{Zn}^{2+}$ ions, causing a buildup of positive charge in the solution.

• At the Cathode: $\text{Cu}^{2+}$ ions are reduced to solid copper, leaving behind an excess of negative sulfate ions ($\text{SO}_4^{2-}$). The salt bridge provides a path for ions to migrate into these half-cells to neutralize the excess charges, preventing the reaction from stopping prematurely.

2. Completing the Electrical Circuit

The salt bridge allows for the flow of current by completing the internal circuit of the cell. While electrons flow through the external wire from the zinc anode to the copper cathode, the migration of ions through the salt bridge ensures a continuous loop of moving charge.

Summary of Functions

• Ion Migration: It allows ions to move from one half-cell to another without the solutions actually mixing.

• Prevents Polarization: By keeping the solutions neutral, it ensures that the electrical potential remains steady so the cell can continue providing power.

Gibbs Energy

Gibbs Free Energy (denoted by the symbol $G$) is a thermodynamic quantity used to predict the spontaneity of a chemical reaction and determine the maximum useful work that can be obtained from an electrochemical cell.

Gibbs Energy and Spontaneity

In electrochemistry, the change in Gibbs energy ($\Delta G$) indicates whether a reaction will occur naturally or require external energy:

• Spontaneous Reaction (\Delta G < 0): The Gibbs energy decreases, meaning the reaction occurs naturally and releases energy that can be converted into electrical work, as seen in Galvanic cells.

• Non-Spontaneous Reaction ($\Delta G > 0$): The Gibbs energy increases, meaning the reaction cannot occur on its own and requires an external input of electrical energy to proceed, as seen in Electrolytic cells.

Relation to Cell Potential

There is a direct mathematical relationship between the Gibbs energy change and the electrical potential produced by a cell:

$\Delta G = -n F E_{cell}$

• $n$: The number of moles of electrons transferred in the redox reaction.

• $F$: Faraday's constant (approximately $96485\text{ C/mol}$).

• $E_{cell}$: The electrical potential or electromotive force (EMF) of the cell.

Why it Matters

• Work Capacity: The decrease in Gibbs energy ($-\Delta G$) is equal to the maximum electrical work that can be done by a galvanic cell.

• Equilibrium: If $\Delta G = 0$, the system is at equilibrium, and no net work can be performed by the cell.

Electrode Potential

• Measured as the tendency of a half-cell to gain or lose electrons.

• Standard Hydrogen Electrode (SHE) is the reference electrode, with potential defined as 0 V.

• Emf (cell potential) is the potential difference between electrodes and determines the direction and magnitude of electron flow.

Electrical potential

Based on the provided materials, electrical potential in a cell (often referred to as cell potential) is the force that drives the movement of electrons through an external circuit.

In an electrochemical cell, this potential arises from the difference in the ability of the two electrodes to lose or gain electrons.

Key Concepts of Cell Potential

• Driving Force: It represents the "push" given to electrons by the chemical reactions occurring at the electrodes.

• Origin: It is created by the simultaneous processes of oxidation (loss of electrons) at the anode and reduction (gain of electrons) at the cathode.

• Energy Conversion: In a galvanic cell, this potential allows for the conversion of chemical energy into electrical energy.

• Spontaneity: For a galvanic cell like the Daniell cell, the electrical potential is generated by a spontaneous redox reaction.

Practical Example: The Daniell Cell

In a Daniell cell, the electrical potential is specifically generated by the reaction between zinc and copper:

• Anode (Oxidation): Zinc atoms lose electrons to become $Zn^{2+}$ ions.

• Cathode (Reduction): Copper ions gain those electrons to become solid copper.

• Result: The difference in electrical potential between these two electrodes allows electricity to flow through an external wire.

Spontaneous and Non Spontaneous Chemical Reaction

In the context of electrochemistry, a non-spontaneous chemical reaction is a process that does not occur on its own under standard conditions. Unlike spontaneous reactions, which release energy to perform work, these reactions require an external input of energy (usually electrical energy) to proceed.

Key Characteristics

• Energy Requirement: These reactions cannot happen without a continuous supply of energy from an external source.

• Electrolytic Cells: While galvanic cells use spontaneous reactions to produce electricity, electrolytic cells use electricity to force non-spontaneous chemical changes to occur.

• Direction: A non-spontaneous reaction is often the reverse of a spontaneous one. For example, while the discharge of a battery is spontaneous, recharging it involves a non-spontaneous reaction driven by an external charger.

Comparison with Spontaneous Reactions

Metal Atom vs Metal Ion

Nernst Equation

Used to calculate electrode/cell potential under non-standard conditions:

E=E${^\circ}$− 0.0591 n $log$ ⁡QE = E${^\circ}$ - ${\frac{0.0591}{n}}$ $\log$ Q

Where E${^\circ}$ is the standard electrode potential, $n$ is the number of electrons transferred, and $Q$ is the reaction quotient.

Conductance in Electrolytic Solutions

• Specific Conductance (KKK): Conductivity of a solution with unit volume.

• Molar Conductivity (Λm\Lambda_mΛm): Conductivity per mole of electrolyte.

• Molar conductivity increases with dilution; for strong electrolytes, it reaches a limiting value (Kohlrausch's Law).vedantu

Batteries and Cells

• Primary Cells: Non-rechargeable (e.g., Leclanche cell).vedantu

• Secondary Cells: Rechargeable (e.g., Lead-acid battery).vedantu

Corrosion

Electrochemical rusting of metals, particularly iron, is discussed as an application of spontaneous electrode reactions.vedantu

Important Formulas

• Cell reaction: Oxidation at anode + reduction at cathode.

• Cell EMF: Ecell=Ecathode−EanodeE_{cell} = E_{cathode} - E_{anode}Ecell=Ecathode−Eanode

Key Definitions and Concepts

• Anode: Where oxidation takes place.learncbse

• Cathode: Where reduction takes place.learncbse

• Electrolysis: Decomposition by electrical energy.byjus

• Electrode potential: Voltage at a particular electrode compared to SHE.slideshare+1

Applications

• Metal extraction, electroplating, batteries, corrosion prevention.ncerthelp

Difference between cells