Electrochemistry - What We Have Learnt

Electrochemistry - What We Have Learned

Slide 1:

Electrochemistry is the branch of chemistry that deals with the interaction between electrical energy and chemical reactions.
It involves the study of redox reactions, electrolysis, and the generation of electricity through chemical reactions.

Slide 2:

Redox reactions involve the transfer of electrons from one species to another.
Oxidation is the loss of electrons, while reduction is the gain of electrons.
The species that gets oxidized is called the reducing agent, and the species that gets reduced is called the oxidizing agent.

Slide 3:

Electrolysis is a process that uses an electric current to drive a non-spontaneous chemical reaction.
During electrolysis, positive ions migrate to the cathode (negative electrode) and gain electrons, while negative ions migrate to the anode (positive electrode) and lose electrons.

Slide 4:

Faraday’s laws of electrolysis describe the relationship between the amount of substance consumed or produced during electrolysis and the amount of electric current passed.
First law: The amount of substance consumed or produced at the electrodes is directly proportional to the quantity of electricity passed through the electrolyte.

Slide 5:

Faraday’s second law states that the mass of different substances deposited or liberated by the same quantity of electricity is directly proportional to their chemical equivalent weights.
The chemical equivalent weight is the mass of a substance that reacts with or is released by one mole of electrons.

Slide 6:

Electrochemical cells are devices that convert chemical energy into electrical energy or vice versa.
They consist of two half-cells connected by a salt bridge or porous membrane, allowing the flow of ions and completing the circuit.

Slide 7:

A half-cell consists of an electrode immersed in an electrolyte solution.
The electrode can be classified as either an anode (where oxidation occurs) or a cathode (where reduction occurs).

Slide 8:

The standard hydrogen electrode (SHE) is used as a reference electrode in electrochemical measurements.
It consists of a platinum electrode immersed in a solution of 1M H+ ions at a pressure of 1 atm.

Slide 9:

The electromotive force (EMF) of a cell is the difference in electrical potential between its two half-cells.
It is measured in volts (V) and determines the direction and magnitude of electron flow in the cell.

Slide 10:

The Nernst equation relates the EMF of a cell to the concentrations of reactants and products involved in the cell reaction.
It can be used to calculate the cell potential under non-standard conditions.

Slide 11:

Galvanic cells, also known as voltaic cells, are electrochemical cells that convert chemical energy into electrical energy.
They are composed of two half-cells, a metal electrode, and an electrolyte solution.
The metal electrode in one half-cell oxidizes, and the metal electrode in the other half-cell reduces.

Slide 12:

The cell potential, or voltage, can be calculated by subtracting the reduction potential of the anode from the reduction potential of the cathode.
The more positive the cell potential, the more spontaneous the reaction and the greater the amount of electrical energy produced.

Slide 13:

The standard cell potential, E°cell, is the cell potential when all species in the reaction are in their standard states.
It can be calculated using the standard reduction potentials for each half-cell reaction.

Slide 14:

The Nernst equation can be used to calculate the cell potential under non-standard conditions using the concentrations of the reactants and products.
It is given by: Ecell = E°cell - (RT/nF) * ln(Q), where R is the gas constant, T is the temperature, n is the number of moles of electrons transferred, F is Faraday’s constant, and Q is the reaction quotient.

Slide 15:

Concentration cells are electrochemical cells where the only difference between the two half-cells is the concentration of the same species.
The cell potential of a concentration cell arises from the difference in concentration of the same species across the two half-cells.

Slide 16:

Batteries are examples of electrochemical cells used to provide portable sources of electrical energy.
They can be classified into primary batteries (non-rechargeable) and secondary batteries (rechargeable).

Slide 17:

An example of a primary battery is the alkaline battery, which uses a zinc anode and a manganese dioxide cathode.
The electrolyte is a potassium hydroxide solution, and the reaction that occurs is: Zn + 2MnO2 + H2O → Zn(OH)2 + Mn2O3.

Slide 18:

An example of a secondary battery is the lead-acid battery, commonly used in cars.
It consists of lead and lead dioxide electrodes in an electrolyte of sulfuric acid.
The reaction that occurs during discharge is: Pb + PbO2 + 2H2SO4 → 2PbSO4 + 2H2O.

Slide 19:

Fuel cells are devices that convert the chemical energy of a fuel, such as hydrogen, directly into electrical energy.
They operate on the principles of electrochemical cells, with the reactants being continuously supplied from an external source.

Slide 20:

A commonly used fuel cell is the hydrogen fuel cell, which uses hydrogen gas as the fuel and oxygen gas as the oxidant.
The overall reaction that occurs is: 2H2 + O2 → 2H2O.
The hydrogen fuel cell is known for its high efficiency and environmentally friendly nature. ``

Slide 21:

Corrosion is the process of the gradual destruction of a metal due to chemical reactions with its environment.
The most common type of corrosion is the oxidation of metals in the presence of oxygen and water.
Example: Rusting of iron: Fe + O2 + H2O → Fe2O3·xH2O

Slide 22:

Corrosion can be prevented or reduced through various methods:
- Coating: Applying a protective layer, such as paint or enamel, to prevent direct contact with the corrosive environment.
- Galvanization: Plating the metal with a more reactive metal, such as zinc, which acts as a sacrificial anode.
- Cathodic protection: Connecting the metal to be protected to a more easily corroded metal, such as magnesium, which becomes the sacrificial anode.

Slide 23:

Galvanization is a common method used to prevent corrosion of iron or steel.
The metal is coated with a layer of zinc, which acts as a sacrificial anode and protects the iron or steel underneath.
The zinc prevents the iron from oxidizing by undergoing oxidation itself.

Slide 24:

Electroplating is a process of coating an object with a thin layer of metal using electrolysis.
It is commonly used to provide a decorative or protective coating to objects.
Example: Electroplating a silver spoon with gold:
- The spoon is connected to the cathode and a gold electrode is connected to the anode.
- A solution containing a gold salt is used as the electrolyte.
- When a current is passed through the circuit, gold ions are reduced at the spoon, forming a thin layer of gold on its surface.

Slide 25:

Electrochemical cells find applications in various industries and devices:
- Batteries: Used as a portable source of electrical energy in various devices such as mobile phones, laptops, and electric vehicles.
- Electrolysis: Used in the production of metals, extraction of reactive metals, and electroplating.
- Fuel cells: Used in power generation for stationary and portable applications.

Slide 26:

The measurement of pH is important in various chemical processes and industries.
pH is a measure of the acidity or alkalinity of a solution and is determined by the concentration of hydrogen ions (H+).
The pH scale ranges from 0 to 14, with pH 7 being neutral, pH less than 7 being acidic, and pH greater than 7 being alkaline or basic.

Slide 27:

The pH of a solution can be measured using indicators or pH meters.
Indicators are substances that change color depending on the pH of the solution.
Example: Litmus paper turns red in acidic solutions and blue in basic solutions.

Slide 28:

Acid-base titrations are commonly used to determine the concentration of an unknown acid or base.
The process involves adding a standardized solution of known concentration (titrant) to the solution of the analyte until the reaction reaches equivalence point.
An indicator is used to signal the completion of the reaction, usually by changing color.

Slide 29:

The equivalence point in an acid-base titration is the point at which the stoichiometrically equivalent amounts of acid and base have reacted.
At this point, the moles of acid are equal to the moles of base.
The pH at the equivalence point depends on the strength of the acid and base being titrated.

Slide 30:

Acid rain is a type of deposition that occurs when sulfur dioxide (SO2) and nitrogen oxides (NOx) emitted from industrial sources react with water in the atmosphere.
Sulfur dioxide and nitrogen oxides dissolve in water to form weak acids, such as sulfuric acid and nitric acid.
Acid rain can have harmful effects on the environment, including damage to plants, aquatic life, and infrastructure. ``