Physical Chemistry - Electrochemistry

Explore the basics of electrochemistry, the interconversion of electrical and chemical energy, and electrochemical changes, including electronic and electrolytic conductors, ion flow, and temperature effects.

Explore how the electrochemical series ranks elements by their tendency to gain or lose electrons, using the standard hydrogen electrode as a reference for reduction potentials.

Explore the electrochemical series, using hydrogen as zero reference to interpret standard reduction potentials, identify oxidizing and reducing agents, and show reactivity decreases down the series.

Explore the electrochemical series to determine oxidizing and reducing strengths, predict metal displacement, assess reaction spontaneity using reduction potentials, and apply zinc-copper and hydrogen liberation examples.

Explain how electrolytes dissociate to increase conductivity, contrast with non-electrolytes, and classify electrolytes as strong or weak with examples like strong acids, strong bases, and water-soluble salts.

Represent galvanic cells using standard electrochemical notation. Place oxidation on the left, reduction on the right, and separate phases with inert electrodes and vertical lines.

Study galvanic cells that spontaneously convert chemical energy to electrical energy and electrolytic cells that use external energy to drive chemical energy, via redox at the anode and cathode.

Explore how a salt bridge maintains electrical neutrality between half-cells in electrochemical cells, preventing mixing while enabling ion transport through inert, nonreactive solutions with matching transport numbers and ionic mobility.

Explore current efficiency in electrochemistry, define its purpose, and present a simple percentage formula using the actual current to mitigate meter reading faults.

emf is the maximum potential difference of a cell, and this lecture derives its temperature dependence using Gibbs free energy under constant pressure, linking emf to delta G and nF.

Explore electrode potential and how concentration, partial pressure, and temperature influence its magnitude, including oxidation and reduction and the standard electrode potential.

Learn how a reference electrode fixes the cell potential, often exactly zero, to measure electrochemical cell potential; primary and secondary references, hydrogen electrodes, and future colonial electrode study.

understand how the standard hydrogen electrode acts as the primary reference electrode, using one molar ion solutions, 1 atm hydrogen, and a platinum electrode to determine other potentials.

Explore the secondary reference calomel electrode, its glass body with a platinum wire and mercury and mercury chloride in potassium chloride, and its fixed potential for zinc-calomel cells.

this lecture presents the Nernst equation, linking electrode potential to concentration and partial pressure at fixed temperature, with e = e° − (RT/nF) ln Q, and 298 K simplification.

Apply the Nernst equation to calculate cell potential in chemical cells, using zinc–copper half-reactions, standard electrode potentials, and ion concentrations to determine Ecell.

Learn to apply the Nernst equation to compute equilibrium constants, Gibbs energy, and cell potentials, and explore standard electrode potentials using the hydrogen electrode example.

Explore the electrical conductance of a solution through ohms law, deriving resistance, conductance, resistivity, and conductivity, and understand their units and relationships in electrochemistry.

measure a solution's conductivity with a Wheatstone bridge and platinum electrodes, determine the cell constant from a known resistance, then compute conductivity as cell constant divided by the unknown resistance.

The lecture explains molar conductivity as a measure that normalizes conductivity by concentration to compare solutions. It shows the formula and connects molar conductivity with normality and equivalent conductivity.

Explores how dilution and concentration affect molar conductivity, using a graph to contrast strong and weak electrolytes and explain limiting molar conductivity and the linear relationship for strong electrolytes.

Identify how ionic attraction, solvent polarity, and viscosity govern electrolyte conductivity, and explain how temperature, dilution, and pressure influence conductivity.

Explore Kohlrausch's law of independent migration of ions, its application to finite dilution, and how to calculate limiting molar conductivity for weak electrolytes using cation and anion contributions.

Calculate the degree of dissociation and the dissociation constant for a weak electrolyte using molar conductivity at finite and zero concentration, and derive their relationship with alpha and concentration.

Learn to compute molar conductivity for weak electrolytes using Kohlrausch's law, relate to strong electrolytes, and derive solubility and Ksp of sparingly soluble salts.

Explore how electrolysis decomposes electrolytes by passing electric current, converting electrical energy into chemical energy at two electrodes, with ions undergoing oxidation and reduction guided by potentials.

Describe how the amount of substance undergoing oxidation or reduction at each electrode scales with the electricity. Define z as the electrochemical equivalent—the amount produced per one coulomb of electricity.

Explains Faraday's second law of electrolysis: when the same amount of electricity flows, deposited substance depends on chemical equivalent mass, as copper deposits with two electrons and silver with one.

Examine how electrolyte type and electrode choice govern products of electrolysis, comparing inert and reactive electrodes, and illustrating with copper sulfate where hydrogen, oxygen, or copper deposit occur.

Explore the dry, non-rechargeable primary voltaic cell (Leclanche cell), where zinc anode and graphite cathode with ammonium chloride and manganese dioxide convert chemical energy to electrical energy via redox reactions.

Explore how the mercury cell uses zinc and mercury amalgam with potassium hydroxide electrolyte to generate about 1.35 volts, powering low-voltage devices like watches and hearing aids.

Explore electrolysis of molten sodium chloride in an electrolytic cell, where Na+ deposits as sodium at the cathode and chlorine gas forms at the anode, driven by electrical energy.

Discover how a lead storage battery—the rechargeable galvanic cell—converts chemical energy to electrical energy through Pb and PbO2 redox reactions in sulfuric acid, and how charging reverses the process.

Explore hydrogen–oxygen fuel cells, a galvanic redox system converting chemical energy to electricity with water as a byproduct, emphasizing continuous reactant supply and their advantages and disadvantages for future vehicles.

Explore how corrosion arises from redox reactions between iron and oxygen in the presence of moisture, forming iron oxide (rust) and triggering galvanic cell processes.

Explore methods to protect metals from corrosion, including painting, galvanizing with zinc, and catholic protection using sacrificial magnesium or chromium alloying to prevent iron rust in moist soil.