Slide 1: Electrochemistry - Introduction

• Electrochemistry is the branch of chemistry that deals with the relationship between electricity and chemical reactions.
• It involves the study of oxidation-reduction (redox) reactions and the use of electrodes to transfer and convert chemical energy into electrical energy.
• Electrochemical processes are important in various applications such as batteries, electrolysis, corrosion, and fuel cells.

Slide 2: Redox Reactions

• Redox reactions involve the transfer of electrons between species.
• Oxidation is the loss of electrons, while reduction is the gain of electrons.
• An oxidizing agent causes another substance to be oxidized, while a reducing agent causes another substance to be reduced.
• The overall reaction is always a combination of both oxidation and reduction.

Slide 3: Oxidation Number

• The oxidation number refers to the charge that an atom would have if electrons were transferred completely.
• Rules for assigning oxidation numbers include:
  • In pure elements, the oxidation number is always zero.
  • The sum of oxidation numbers in a neutral compound is zero.
  • The sum of oxidation numbers in a polyatomic ion equals the ionic charge.
• Example: In H2SO4, the oxidation number of hydrogen is +1, sulfur is +6, and oxygen is -2.

Slide 4: Balancing Redox Reactions - Half-Reaction Method

• The half-reaction method is used to balance redox reactions.
• Split the reaction into two half-reactions, one for the oxidation and one for the reduction.
• Balance the elements other than hydrogen and oxygen in each half-reaction.
• Balance the oxygen atoms by adding water molecules and balance the hydrogen atoms by adding hydrogen ions.
• Balance the charges by adding electrons.
• Finally, multiply the half-reactions by appropriate coefficients to equalize the number of electrons transferred.

Slide 5: Electrochemical Cells

• An electrochemical cell is a device that converts chemical energy into electrical energy through redox reactions.
• The two types of electrochemical cells are galvanic (voltaic) cells and electrolytic cells.
• In a galvanic cell, a spontaneous redox reaction occurs to produce electrical energy.
• In an electrolytic cell, an external source of electrical energy is used to drive a non-spontaneous redox reaction.

Slide 6: Galvanic (Voltaic) Cells

• A galvanic cell consists of two half-cells connected by a salt bridge or porous barrier.
• Each half-cell contains an electrode and an electrolyte solution.
• The anode is the electrode where oxidation occurs, and the cathode is the electrode where reduction occurs.
• Electrons flow from the anode to the cathode through an external circuit, generating a current.

Slide 7: Standard Electrode Potentials

• Standard electrode potential is the measure of the tendency of a half-reaction to occur as a reduction.
• The standard hydrogen electrode (SHE) is used as a reference electrode with a potential of 0 volts.
• Reduction potentials of other half-reactions are measured relative to the SHE.
• Positive standard electrode potentials indicate that the reaction is more likely to occur as reduction.

Slide 8: Nernst Equation

• The Nernst equation relates the electrode potential with the concentration of species involved in the half-reaction.
• It can be used to calculate the cell potential under non-standard conditions.
• The Nernst equation is given as: Ecell = E°cell - (0.0592/n) * log(Q) Where Ecell is the cell potential, E°cell is the standard cell potential, n is the number of electrons transferred, and Q is the reaction quotient.

Slide 9: Fuel Cells

• A fuel cell is an electrochemical cell that converts the chemical energy of a fuel (usually hydrogen) and an oxidizing agent (usually oxygen) into electrical energy.
• The fuel cell operates continuously as long as fuel and oxidant are supplied.
• It provides a clean and efficient method for the production of electricity.

Slide 10: Electrolysis

• Electrolysis is the process of using electrical energy to drive a non-spontaneous redox reaction.
• It involves the decomposition of a compound into its constituent elements or ions.
• The substance undergoing electrolysis is called the electrolyte, and the ends of the electrolytic cell where the current enters and leaves are called electrodes.
• During electrolysis, positive ions move towards the cathode (reduction) and negative ions move towards the anode (oxidation).

Slide 11: Electrochemistry - More Points on Molar Conductance

• Molar conductance is the conductance of all the ions present in one mole of an electrolyte.
• It is determined by measuring the resistance of the solution and using the equation: Molar conductance (Λm) = $ \frac{{K \times 1000}}{{c}} $ , where K is the conductance in siemens and c is the concentration in moles per liter (mol/L).
• Molar conductance can be calculated for both strong and weak electrolytes.
• For strong electrolytes, molar conductance is the sum of the conductances of all ions.
• For weak electrolytes, molar conductance is lower than for strong electrolytes due to incomplete dissociation.

Slide 12: Factors Affecting Molar Conductance

• Molar conductance is affected by various factors, including:
  • Concentration: Molar conductance decreases with decreasing concentration due to fewer ions present.
  • Temperature: Molar conductance generally increases with increasing temperature due to increased ion mobility.
  • Size and charge of ions: Molar conductance increases with increasing charge and decreasing size of ions.
  • Nature of solvent: Molar conductance varies with different solvents, as some solvents can solvate ions and hinder their movement.

Slide 13: Kohlrausch’s Law

• Kohlrausch’s law states that the molar conductance of an electrolyte at infinite dilution is equal to the sum of the molar conductances of its constituent ions.
• This law helps in determining the degree of dissociation of weak electrolytes.
• It is expressed as: Λ∞ = Λ+ + Λ-, where Λ∞ is the molar conductance at infinite dilution and Λ+ and Λ- are the molar conductances of the cation and anion, respectively.

Slide 14: Faraday’s Laws of Electrolysis

• Faraday’s laws of electrolysis describe the relationship between the quantity of electricity passed through an electrolyte and the chemical reactions occurring during electrolysis.
• Faraday’s First Law: The amount of any substance deposited or liberated during electrolysis is directly proportional to the quantity of electricity passed through the electrolyte.
• Faraday’s Second Law: The amounts of different substances deposited or liberated by the same quantity of electricity are directly proportional to their equivalent weights.

Slide 15: Corrosion

• Corrosion is the gradual destruction of a metal due to chemical reactions with its environment.
• It occurs through oxidation-reduction reactions in the presence of moisture, oxygen, and other corrosive substances.
• Factors affecting corrosion include humidity, temperature, pH, and the presence of salts or acids.
• Corrosion can be prevented by using sacrificial anodes, protective coatings, and controlled environments such as cathodic protection.

Slide 16: Batteries

• A battery is an electrochemical device that converts chemical energy into electrical energy.
• It consists of two or more electrochemical cells connected in series or parallel.
• In a cell, two electrodes (anode and cathode) are immersed in an electrolyte solution.
• During discharge, a spontaneous redox reaction occurs, producing electrical energy.
• Common types of batteries include lead-acid, alkaline, lithium-ion, and nickel-metal hydride batteries.

Slide 17: Primary and Secondary Cells

• Primary cells are non-rechargeable cells that are used until their chemical reactants are consumed.
• Examples of primary cells include dry cells (used in flashlights) and button cells (used in watches).
• Secondary cells are rechargeable cells that can be used multiple times by recharging them with electrical energy.
• Examples of secondary cells include automotive batteries and lithium-ion batteries used in smartphones and laptops.

Slide 18: Standard Hydrogen Electrode

• The standard hydrogen electrode (SHE) is commonly used as a reference electrode for measuring electrode potentials.
• It consists of a platinum electrode immersed in a solution of 1M HCl and 1M H2 gas maintained at 1 atm pressure.
• The electrode potential of the SHE is defined as 0 volts by convention.
• All other electrode potentials are measured relative to the SHE.

Slide 19: Electroplating

• Electroplating is the process of depositing a metal coating onto a metallic surface by means of electrolysis.
• It is commonly used to improve the appearance, corrosion resistance, and durability of objects.
• The object to be plated acts as the cathode, and a metal salt solution is used as the electrolyte.
• The metal cations from the electrolyte are reduced at the cathode, forming a metal coating.

Slide 20: Industrial Applications of Electrochemistry

• Electrochemical processes have various industrial applications, such as:
  • Electrorefining: Purification of metals by electroplating.
  • Electrolysis of water: Production of hydrogen and oxygen gases.
  • Electrosynthesis: Production of valuable chemicals through electrochemical reactions.
  • Electrowinning: Extraction of metals from ores using electrolysis.
  • Electrochemical machining: Precise removal of metal by controlled dissolution.

Slide 21: Electrochemistry - More points on Molar Conductance

• Molar conductance is the measure of the conductance of all ions in one mole of an electrolyte.
• It is influenced by the degree of dissociation and the mobility of ions in the solution.
• For strong electrolytes, which completely dissociate in solution:
  • Molar conductance is higher due to a larger number of ions present.
• For weak electrolytes, which only partially dissociate in solution:
  • Molar conductance is lower due to a smaller number of ions present.
  • The degree of dissociation can be determined by comparing the molar conductance at a certain concentration with the molar conductance at infinite dilution.

Slide 22: Factors Affecting Molar Conductance (Continued)

• Effect of concentration:
  • Molar conductance decreases with increasing concentration due to the increase in ionic interactions and decrease in ion mobility.
  • At infinite dilution, ions are far apart, and conductance is at its maximum.
• Effect of temperature:
  • Molar conductance generally increases with increasing temperature due to increased ion mobility.
  • Higher temperatures provide more kinetic energy, allowing ions to move more rapidly.
• Effect of size and charge of ions:
  • Molar conductance increases as the charge on an ion increases because higher charges lead to higher conductivity.
  • Molar conductance decreases as the size of ions increases due to increased resistance to motion.

Slide 23: Kohlrausch’s Law - Application

• Kohlrausch’s law can be used to determine the degree of dissociation of weak electrolytes.
• By comparing the molar conductance of a weak electrolyte at a certain concentration with its molar conductance at infinite dilution, the degree of dissociation can be calculated.
• Example:
  • The molar conductance of acetic acid at 0.1 M concentration is 283 S cm^2 mol^-1, while its molar conductance at infinite dilution is 393 S cm^2 mol^-1.
  • Using Kohlrausch’s law, we can calculate the degree of dissociation as: α = (molar conductance at given concentration / molar conductance at infinite dilution) * 100 α = (283 S cm^2 mol^-1 / 393 S cm^2 mol^-1) * 100 α = 71.9%

Slide 24: Faraday’s Laws of Electrolysis - First Law

• Faraday’s First Law states that the amount of any substance liberated (in electrolysis) or deposited (in electroplating) is directly proportional to the quantity of electricity passed through the electrolyte.
• The relationship can be expressed as: M = $ \frac{{Q}}{{n \cdot F}} $ Where M is the amount of substance in grams, Q is the quantity of electricity in coulombs, n is the number of electrons transferred, and F is Faraday’s constant (96485 C/mol).
• Example: To deposit 1 gram of copper, 96485 C of electricity is required (since n = 2 for copper).
• The Stoichiometry law is also applicable here.

Slide 25: Faraday’s Laws of Electrolysis - Second Law

• Faraday’s Second Law states that the amounts of different substances deposited or liberated by the same quantity of electricity are directly proportional to their equivalent weights.
• The relationship can be expressed as: $ \frac{{M_1}}{{m_1}} = \frac{{M_2}}{{m_2}} = \frac{{M_3}}{{m_3}} = \ldots $ Where M is the molar mass of the substance and m is its equivalent mass. The equivalent weight is equal to the molar mass divided by the valency.
• Example: To produce 1 gram of hydrogen gas (H2), 96500 C of electricity is required. To produce 31 grams of zinc (Zn), 96500 C of electricity is required (since the equivalent mass of zinc is 31 g/mol).

Slide 26: Corrosion - Introduction

• Corrosion is the gradual destruction of a metal due to chemical reactions with its environment.
• It is a common problem for metals used in construction, transportation, and many other applications.
• Corrosion occurs through electrochemical reactions in the presence of moisture, oxygen, and other corrosive substances.
• The most well-known form of corrosion is rusting, which happens when iron reacts with oxygen and water to form hydrated iron(III) oxide.

Slide 27: Factors Affecting Corrosion

• Corrosion is influenced by various factors:
  • Humidity: Higher humidity increases the likelihood of corrosion due to the presence of moisture.
  • Temperature: Higher temperatures can accelerate corrosion reactions.
  • pH: Corrosion rates can increase in acidic or alkaline environments.
  • Presence of salts or acids: These compounds can promote corrosion by facilitating electrolytic reactions.
  • Dissimilar metals: When two different metals are in contact, corrosion can occur due to galvanic interactions.

Slide 28: Corrosion Prevention - Sacrificial Anodes

• Sacrificial anodes are a common method used to prevent corrosion of a metal.
• A more reactive metal is electrically connected to the metal to be protected (cathode).
• The sacrificial metal will corrode instead of the protected metal.
• Examples of sacrificial anodes include zinc or magnesium attached to the hull of a ship or the underground pipelines.

Slide 29: Corrosion Prevention - Protective Coatings

• Protective coatings are applied to metal surfaces to act as a barrier against corrosive agents.
• Examples of protective coatings include paint, varnish, enamel, and galvanizing (zinc coating).
• These coatings prevent direct contact between the metal surface and the corrosive environment, inhibiting corrosion reactions.
• Regular inspection and maintenance of these coatings are essential to ensure their effectiveness.

Slide 30: Corrosion Prevention - Cathodic Protection

• Cathodic protection is a technique used to prevent corrosion by making the metal to be protected the cathode of an electrochemical cell.
• Two common methods of cathodic protection are impressed current cathodic protection (ICCP) and sacrificial anode cathodic protection.
• ICCP involves the use of an external power source (e.g., impressed current) to drive the corrosion reaction towards the cathodic direction.
• Sacrificial anode cathodic protection involves using sacrificial anodes, similar to the sacrificial anode method discussed earlier.