Electrochemistry

Electrochemistry - Recap

Slide 1

Electrochemistry is the branch of chemistry that deals with the relationship between electricity and chemical reactions.
It involves the study of redox reactions, where one species loses electrons (oxidation) and another gains electrons (reduction).
Redox reactions take place in electrochemical cells, where chemical energy is converted into electrical energy.

Slide 2

In an electrochemical cell, there are two half-cells: an anode (where oxidation occurs) and a cathode (where reduction occurs).
The anode is the electrode where oxidation takes place, while the cathode is where reduction takes place.
Electrons flow from the anode to the cathode through an external circuit, generating electricity.

Slide 3

The movement of ions in the electrolyte solution completes the circuit and balances the charges.
The electrolyte can be an aqueous solution or a molten salt bridge.
The flow of positive ions from the anode to the cathode is called the cationic flow, while the flow of negative ions is called anionic flow.

Slide 4

The standard hydrogen electrode (SHE) is used as a reference electrode in electrochemical cells.
It consists of a platinum electrode in contact with a solution containing dissolved hydrogen gas at 1 atmosphere pressure.
The standard reduction potential of the SHE is arbitrarily defined as 0 volts.

Slide 5

The standard reduction potential (E°) is a measure of the tendency of a species to gain electrons and undergo reduction.
It is measured relative to the standard hydrogen electrode (SHE) and can be used to predict the direction of redox reactions.
An E° value more positive than 0 indicates a strong tendency to be reduced, while a value less than 0 indicates a strong tendency to be oxidized.

Slide 6

The Nernst equation relates the standard reduction potential (E°), actual cell potential (E), and concentrations of reactants and products.
It is given by: E = E° - (RT / nF) * ln(Q)
Where E is the actual cell potential, R is the gas constant, T is the temperature in Kelvin, n is the number of electrons transferred, F is Faraday’s constant, and Q is the reaction quotient.

Slide 7

Galvanic cells, also known as voltaic cells, produce electrical energy through spontaneous redox reactions.
The cell potential (E°cell) is positive for galvanic cells, indicating that the reaction is thermodynamically favorable.
The electrons flow from the anode (where oxidation occurs) to the cathode (where reduction occurs) through an external circuit.

Slide 8

Electrolytic cells use electrical energy to drive non-spontaneous redox reactions in the opposite direction.
The cell potential (E°cell) is negative for electrolytic cells, indicating that the reaction is not thermodynamically favorable.
The electrons flow from the external power source to the cathode (where reduction occurs) and from the anode (where oxidation occurs) back to the power source.

Slide 9

Faraday’s laws of electrolysis describe the relationship between the amount of substance undergoing electrolysis and the quantity of electricity passed through the cell.
Faraday’s first law states that the amount of chemical change is directly proportional to the quantity of electricity passed.
Faraday’s second law states that the amount of chemical change is directly proportional to the equivalent weight of the substance.

Slide 10

Electroplating is a common application of electrolysis, where a metal is deposited onto another metal surface.
It is used to improve the appearance, corrosion resistance, and wear resistance of objects.
The metal to be plated acts as the cathode, while the metal to be deposited is usually an anode made of that metal.

The Nernst equation is a fundamental equation in electrochemistry that relates the cell potential to the concentrations of reactants and products.
It is given by: E = E° - (RT / nF) * ln(Q), where E is the cell potential, E° is the standard reduction potential, R is the gas constant, T is the temperature, n is the number of electrons transferred, F is Faraday’s constant, and Q is the reaction quotient.
The Nernst equation allows us to calculate the cell potential under non-standard conditions and understand how changes in concentration affect the reaction.

Concentration cells are electrochemical cells where the half-cells have the same components, but different concentrations.
The cell potential in a concentration cell arises from the difference in concentration of the species involved.
For example, a concentration cell can be created using two silver/silver chloride electrodes with different chloride ion concentrations. The higher chloride concentration will result in a lower cell potential.

Batteries are portable electrochemical cells that provide electrical energy for various devices.
Primary batteries are non-rechargeable and have a limited lifespan, while secondary batteries are rechargeable.
Common types of batteries include lead-acid batteries, nickel-cadmium batteries, and lithium-ion batteries.

Lead-acid batteries are commonly used in automotive applications. They consist of lead and lead dioxide electrodes immersed in a sulfuric acid electrolyte.
During discharge, lead and lead dioxide react with sulfuric acid to form lead sulfate and water.
During charging, the reaction is reversed, converting lead sulfate back into lead and lead dioxide.

Nickel-cadmium (NiCd) batteries have a nickel hydroxide cathode, a cadmium anode, and a potassium hydroxide electrolyte.
During discharge, the cadmium is oxidized to cadmium hydroxide at the anode, while the nickel hydroxide is reduced to nickel oxyhydroxide at the cathode.
During charging, the reaction is reversed, converting cadmium hydroxide back into cadmium and nickel oxyhydroxide back into nickel hydroxide.

Lithium-ion batteries are commonly used in portable electronic devices. They have a lithium cobalt oxide or lithium iron phosphate cathode, a graphite anode, and an organic electrolyte.
During discharge, lithium ions are oxidized at the cathode and intercalated into the graphite anode, releasing electrical energy.
During charging, the reaction is reversed, intercalating lithium ions back into the cathode and extracting electrical energy from an external source.

Fuel cells are devices that convert the chemical energy of a fuel and oxidant into electrical energy and thermal energy.
The most common type is the hydrogen fuel cell, where hydrogen acts as the fuel and oxygen acts as the oxidant.
In hydrogen fuel cells, hydrogen is oxidized at the anode to produce protons and electrons. The protons pass through an electrolyte to the cathode, while the electrons flow through an external circuit to the cathode, completing the circuit and generating electricity.

Corrosion is a natural electrochemical process that occurs when metals react with their surroundings.
It is an oxidation process where metals are oxidized to their respective metal ions.
Factors such as moisture, oxygen, impurities, and temperature influence the rate of corrosion.

Ways to prevent or control corrosion include using protective coatings, applying sacrificial anodes, employing cathodic protection, and using corrosion-resistant alloys.
Protective coatings, such as paint or varnish, create a barrier between the metal and its surroundings, preventing contact and oxidation.
Sacrificial anodes are more chemically active metals that are attached to the metal structure, sacrificing themselves to protect the metal from corrosion.

Cathodic protection involves making the metal to be protected the cathode in an electrochemical cell.
This can be achieved by connecting the metal to a more reactive metal, such as a zinc or magnesium anode, which becomes the anode in the cell.
The more reactive metal will corrode instead of the metal to be protected, offering cathodic protection.

Slide 21

pH is a measure of the acidity or basicity of a solution.
It is defined as the negative logarithm of the hydrogen ion concentration (pH = -log[H+]).
Acidic solutions have a pH less than 7, while basic solutions have a pH greater than 7.
A neutral solution, such as pure water, has a pH of 7.

Slide 22

The pH scale is logarithmic, meaning that each unit represents a tenfold difference in acidity or basicity.
For example, a solution with a pH of 3 is ten times more acidic than a solution with a pH of 4.
Similarly, a solution with a pH of 10 is ten times more basic than a solution with a pH of 9.

Slide 23

The pH of a solution can be measured using indicators, pH paper, or a pH meter.
Indicators are substances that change color depending on the pH of the solution they are in.
Common indicators include litmus paper, phenolphthalein, and universal indicator.

Slide 24

Buffer solutions are solutions that resist changes in pH when small amounts of acid or base are added to them.
They consist of a weak acid and its conjugate base (or a weak base and its conjugate acid) in roughly equal concentrations.
Buffer solutions are important in biological systems, where maintaining a stable pH is crucial for proper function.

Slide 25

The Henderson-Hasselbalch equation relates the pH of a buffer solution to the pKa of the weak acid (or base) and the ratio of the concentration of the conjugate base to the weak acid (or conjugate acid to the weak base).
It is given by: pH = pKa + log([conjugate base]/[weak acid])
The Henderson-Hasselbalch equation helps us understand how changes in the concentration of the conjugate base and weak acid affect the pH of a buffer solution.

Slide 26

Acid-Base titrations are used to determine the concentration of an acid or base in a solution.
A titration involves adding a solution of known concentration (the titrant) to a solution of unknown concentration (the analyte) until the reaction is complete.
The equivalence point is the point at which stoichiometrically equivalent amounts of the acid and base have reacted, and the solution is neutral.

Slide 27

The pH at the equivalence point of an acid-base titration depends on the strength of the acid and base being titrated.
For a strong acid and strong base, the pH at the equivalence point is 7, indicating a neutral solution.
However, for a weak acid and strong base titration, the pH at the equivalence point will be greater than 7, indicating a basic solution.
Similarly, for a strong acid and weak base titration, the pH at the equivalence point will be less than 7, indicating an acidic solution.

Slide 28

The pH of a solution can also be influenced by the presence of common ions.
Common ion effect occurs when a solution already contains one of the ions involved in the reaction.
The presence of the common ion suppresses the ionization of a weak acid or base, shifting the equilibrium and resulting in a lower pH (for an acid) or higher pH (for a base).

Slide 29

Solubility is the ability of a substance to dissolve in a solvent and form a homogeneous mixture.
The solubility of a solute depends on various factors, including temperature, pressure, and the nature of the solute and solvent.
The solubility of most solid solutes increases with increasing temperature, while the solubility of most gas solutes decreases with increasing temperature.

Slide 30

Solubility can be expressed in different ways, such as in grams per liter (g/L) or moles per liter (mol/L).
The solubility product constant (Ksp) is a measure of the extent to which a sparingly soluble salt dissociates into its constituent ions in a saturated solution.
It is the product of the concentration of the ions raised to the power of their respective stoichiometric coefficients in the balanced chemical equation.