Electrochemistry

Electrochemistry is a branch of chemistry that deals with the study of the relationship between electricity and chemical reactions. It explores how chemical reactions can either produce or consume electrical energy. The fundamental processes involved in electrochemistry revolve around the transfer of electrons between substances undergoing oxidation and reduction reactions, commonly known as redox reactions.

1.0What is Electrochemistry?

Electrochemistry explores chemical processes where charges separate and transfer. This can happen uniformly in a solution or on the surfaces of electrodes. To maintain balance, two or more charge transfer reactions occur, usually at different electrodes in a cell. These electrodes are connected by ionic transport in the solution and external paths like wires.

In an electrochemical cell, products from the two electrodes can be separated. If the combined energy change at both electrodes is negative, it releases electrical energy, useful in batteries. Conversely, if it's positive, external electrical energy can drive reactions, leading to chemical transformations in processes like electrolysis.

2.0Electrochemical Cells 

These are devices that enable the conversion of chemical energy into electrical energy and vice versa. Galvanic (voltaic) cells spontaneously generate electrical energy through redox reactions, while electrolytic cells use external electrical energy to drive non-spontaneous reactions. Let’s learn more about cells.

Galvanic Cells (voltaic cells)

A galvanic cell, also known as a voltaic cell, operates based on the principle of converting chemical energy into electrical energy through spontaneous redox reactions. It involves two half-cells, each containing an electrode immersed in an electrolyte. The potential difference between these electrodes generates an electric current.

Parts of Galvanic Cell:

• Anode: The electrode where oxidation occurs. Electrons are released from the anode into the external circuit.
• Cathode: The electrode where reduction occurs. Electrons from the external circuit are accepted by the cathode.
• Electrolyte: A solution containing ions that allow the flow of electric current between the anode and cathode.
• Salt Bridge: A bridge that connects the two half-cells, allowing ion flow to maintain charge neutrality. It prevents the mixing of electrolytes.
• External Circuit: The pathway through which electrons flow from the anode to the cathode, generating an electric current.

Working of Galvanic Cell:

• Oxidation at Anode: At the anode, the metal undergoes oxidation, losing electrons and forming ions. For example, in a copper-zinc galvanic cell, zinc (Zn) might oxidize to form Zn²⁺ ions.

Zn(s)   →  Zn²⁺(aq)   +   2e⁻

• Reduction at Cathode: Simultaneously, at the cathode, reduction occurs. For example, in the same copper-zinc cell, copper ions (Cu²⁺) may gain electrons to form copper metal.

Cu²⁺(aq)    +   2e⁻  →    Cu(s)

• Flow of Electrons: Electrons released at the anode flow through the external circuit to the cathode.
• Ion Flow: Ions move through the electrolyte, and the salt bridge maintains charge balance.
• Cell Potential: The potential difference between the anode and cathode generates electrical energy. The higher the potential difference, the greater the cell's ability to produce electricity.

Example of Galvanic Cell:

Consider the Daniell cell, which consists of a zinc electrode in a ZnSO₄ solution as the anode and a copper electrode in a CuSO₄ solution as the cathode. The overall cell reaction is:

Zn(s)  +  Cu²⁺(aq)    →   Zn²⁺(aq)   +   Cu(s)

This cell produces a potential difference, and electrons flow from zinc to copper, generating electrical energy.

Salt Bridge: A salt bridge is a component of an electrochemical cell that serves to maintain electrical neutrality within the system. It consists of an inert electrolyte (usually a concentrated KNO₃ or NH₄NO₃ solution) housed in a tube or bridge connecting the two half-cells of the electrochemical cell. The primary functions of the salt bridge include:

Ion Flow: The salt bridge allows the flow of ions between the two half-cells, enabling the completion of the electrical circuit.

Preventing Charge Build-Up: As electrons flow from the anode to the cathode in a galvanic cell, a negative charge builds up at the anode, and a positive charge builds up at the cathode. The salt bridge helps prevent the excessive accumulation of charges by allowing ions to move, maintaining overall electrical neutrality.

Maintaining Electrolyte Balance: In electrolytic cells and some galvanic cells, where chemical reactions occur that involve the consumption or production of ions, the salt bridge helps maintain a balance of ions between the two half-cells.

Preventing Mixing of Solutions: The salt bridge physically separates the two half-cells, preventing the direct mixing of the electrolytes in the two compartments, which could disrupt the electrochemical processes.

Electrode Potential: Electrode potential refers to the potential difference between an electrode and its surrounding solution in an electrochemical cell. It is a measure of the tendency of a half-reaction to occur at a particular electrode. The standard electrode potential (E°) is measured under standard conditions (1 M concentration, 1 atm pressure, 25°C temperature).

Reduction Potential: For a half-reaction, the reduction potential (E°red) indicates the tendency of a species to gain electrons and undergo reduction. A higher E° red value corresponds to a stronger tendency to be reduced.

Oxidation Potential: The oxidation potential (E°ox) is related to the reverse of the reduction potential. It indicates the tendency of a species to lose electrons and undergo oxidation. The more positive the E°ox value, the stronger the tendency to be oxidized.

Cell Potential: In a complete electrochemical cell, the overall cell potential (E°cell) is the difference between the reduction potentials of the cathode and anode. It provides information about the spontaneity of the overall redox reaction.

The Nernst equation allows for the calculation of electrode potential under non-standard conditions, considering factors such as concentration changes.

3.0Nernst equation 

The Nernst equation is a mathematical expression that relates the electrode potential of an electrochemical cell to the concentrations of reactants and products at non-standard conditions. It helps determine the cell potential under conditions where the concentrations of species involved in the cell reactions are not at their standard state (1 M concentration).

The Nernst equation is given by:

E_Cell = E° – $\frac{RT}{nF}$ ln Q  =  E° – $\frac{2.303RT}{nF}$ log Q

Where,

R = Gas constant

T = Temperature (in K)

F = Faraday (96500 coulomb mol^–1)

n = No. of e^⊝ gained or lost in the balanced equation.

Q = Reaction quotient 

$\frac{2.303\times 8.314\times 298}{96500}$ = 0.059 volt (At 298 K)

Note:

 (i)   For writing the Nernst equation, first write a balanced cell reaction.

 (ii) Nernst equation can be applied on half-cell as well as complete Galvanic cell reaction.

The Nernst equation provides a more accurate prediction of cell potential under various conditions, taking into account the impact of concentration changes on the spontaneity of the redox reaction. When

Q = K (equilibrium conditions), the Nernst equation reduces to the standard cell potential (E=0), and the system is at equilibrium.

4.0Electrochemical Series

The electrochemical series is a list of elements and their standard electrode potentials, arranged in decreasing order of their reduction potentials. It provides insights into the tendency of each element to undergo reduction or oxidation reactions.

Note - Here is a Mnemonic to memorize the above electrochemical series in a simpler way.

“Please stop calling Me A Zebra I Truly Like Her Calling Me Super”

The sentence represents the order of metals from most reactive to least reactive, using the first letter of each metal.

"Please: Potassium

Stop: Sodium

Calling: Calcium

Me: Magnesium

A: Aluminium

Zebra: Zinc

I: Iron

Truly: Tin

Like: Lead

Her: Hydrogen

Calling: Copper

Me: Mercury

Super: Silver"

5.0Electrolytic Cells

Electrolytic cells are devices that use an external electric current to drive a non-spontaneous chemical reaction, causing a breakdown of compounds into their constituent elements or ions. These cells are essential for various applications, including metal extraction, electroplating, and the production of gases.

Components of Electrolytic Cells:

Electrolyte:

• The electrolyte is the substance undergoing electrolysis.
• It can be either a molten ionic compound or a concentrated aqueous solution of ions.
• The choice of electrolyte depends on the specific reaction and the desired products.

Electrodes:

• Electrodes are conductive rods usually made of metal or carbon (often graphite).
• They facilitate the flow of electricity to and from the electrolyte.
• There are two types of electrodes: anode and cathode.

Anode:

• The anode is the positive electrode in the electrolytic cell.
• It is where oxidation occurs during electrolysis.
• Electrons are released at the anode, contributing to the overall current flow.

Cathode:

• The cathode is the negative electrode in the electrolytic cell.
• It is where reduction takes place during electrolysis.
• Electrons are accepted at the cathode, completing the overall circuit.

External Power Source:

• An external power source, such as a battery or a direct current (DC) power supply, provides the electric current needed for electrolysis.
• It drives the non-spontaneous reaction by supplying energy.

6.0Electrolysis

Electrolysis is a chemical process that involves the decomposition of a compound through the passage of an electric current. It typically occurs in an electrolytic cell, where two electrodes are immersed in an electrolyte solution. When an electric current is passed through the solution, ions in the electrolyte migrate towards the electrodes. 

At the electrodes, the ions undergo oxidation or reduction reactions, leading to the formation of new substances. This process allows for the separation of elements or the production of new compounds.

Examples of electrolysis include:

1. Electrolysis of water: 

• Electrolysis can be used to decompose water into hydrogen and oxygen gases. When an electric current is passed through water, hydrogen ions (H⁺) are attracted to the cathode (negative electrode) and undergo reduction to form hydrogen gas (H₂), while hydroxide ions (OH^-) are attracted to the anode (positive electrode) and undergo oxidation to form oxygen gas (O₂).

2. Electrolysis of Brine: 

• Brine, a solution of sodium chloride (NaCl) in water, can be electrolyzed to produce chlorine gas (Cl₂), hydrogen gas (H₂), and sodium hydroxide (NaOH). At the cathode, sodium ions (Na⁺) are reduced to form sodium metal (Na) and hydrogen gas, while at the anode, chloride ions (Cl^-) are oxidized to form chlorine gas.

The half-reactions and cell reaction for the electrolysis of aqueous sodium chloride to chlorine and hydroxide ion are as follows:

Because the electrolysis started with sodium chloride, the cation in the electrolyte solution is Na⁺. On evaporation of the electrolyte solution, sodium hydroxide NaOH is obtained.

3. Electroplating: 

Electroplating is a process where a metal object is coated with a thin layer of another metal using electrolysis. For example, silverware can be electroplated with a thin layer of gold to improve its appearance and corrosion resistance. In this process, the metal object to be plated (cathode) and a metal source (anode) are immersed in an electrolyte solution containing ions of the metal to be plated. When an electric current is passed through the solution, metal ions are reduced at the cathode, depositing a layer of metal onto the object.

Faraday’s laws of Electrolysis

1. First law of electrolysis : The amount of substance deposited or liberated at an electrode is directly proportional to the amount of charge passed through the solution.

W ∝ Q

W = amount of substance deposited, Q = charge in coulomb

W = ZQ

Z = electrochemical equivalent

when Q = 1 coulomb, then W = Z

Thus, the amount of substance deposited or liberated by 1 coulomb charge is called electrochemical equivalent.

Let ‘I’ ampere current is passed till ‘t’ seconds.

Then, Q = I t     

W  = Z I t  

1 Faraday = 96500 coulomb = Charge on one mole electrons.

2. Second law of electrolysis:

When the same amount of charge is passed through different electrolyte solutions connected in series then weight of substances deposited or liberated at electrodes are in ratio of their respective equivalent weights.

7.0Batteries

Batteries are electrochemical devices that store and provide electrical energy through reversible chemical reactions. They are classified into primary and secondary batteries based on their ability to be recharged.

1. Primary Batteries:

Primary batteries, also known as non-rechargeable batteries, are designed for single use and cannot be recharged. Once their chemical reactions are exhausted, they must be replaced. Primary batteries are commonly used in devices where long-term reliability and low maintenance are essential.

Examples of Primary Batteries are Zinc-carbon batteries (e.g., 9-volt batteries) and Lithium batteries.

2. Secondary Batteries:

Secondary batteries, also known as rechargeable batteries, can be recharged by passing an electric current through them in the opposite direction of discharge. They are more cost-effective and environmentally friendly than primary batteries, as they can be reused multiple times.

Examples of secondary batteries include: Lead-acid batteries (e.g., car batteries), Nickel-cadmium (NiCd) batteries.

Below is a comparison table outlining the main differences between primary and secondary batteries:

8.0Fuel Cell

A fuel cell is an electrochemical device that converts the chemical energy of a fuel and an oxidizing agent directly into electricity. It operates much like a battery, but it does not require recharging and can continuously generate electricity as long as fuel and oxidizer are supplied. Here are some key points about fuel cells:

Operation: Fuel cells generate electricity through an electrochemical reaction between a fuel, such as hydrogen, and an oxidizing agent, typically oxygen or air.

Types: There are several types of fuel cells, including proton exchange membrane (PEM) fuel cells, alkaline fuel cells (AFCs), solid oxide fuel cells (SOFCs), and molten carbonate fuel cells (MCFCs), among others.

Efficiency: Fuel cells can achieve high efficiencies, especially when used in combined heat and power (CHP) applications where waste heat is captured and used for heating or cooling.

Clean Energy: They produce electricity with high efficiency and virtually no harmful emissions, making them a clean and environmentally friendly energy source.

Applications: Fuel cells are used in various applications, including stationary power generation, transportation (e.g., fuel cell vehicles), portable power systems, and backup power for critical infrastructure.

Challenges: Despite their advantages, fuel cells face challenges such as high cost, limited durability, and the need for infrastructure to support widespread adoption, particularly in the transportation sector.

Reactions:

In a hydrogen fuel cell, the following reactions occur:

• Anode (Oxidation):

2H₂ → 4H⁺+ 4e⁻

• Cathode (Reduction):

O₂  +  4H⁺  + 4e⁻ → 2 H₂O

• Overall Reaction:

2H₂  +  O₂  →  2H₂O

• 9.0Corrosion

  Corrosion is the gradual deterioration of metals due to chemical reactions with the environment. It weakens structures, affects functionality, and poses safety risks. Prevention methods include protective coatings, corrosion-resistant alloys, and maintenance practices.