Electrochemistry and Redox Reactions

• Electrochemistry is the study of chemical processes that cause electrons to move.
• Electricity is generated by movements of electrons from one element to another.
• These movements of electrons occur during a reaction known as an oxidation-reduction ("redox") reaction.

Redox Reactions

• Redox reactions involve electron transfer between chemical species.
• Oxidation involves the loss of electrons.
• Reduction involves the gain of electrons.
• Oxidizing agents cause oxidation and gain electrons and are thus reduced.
• Reducing agents cause reduction and lose electrons and are thus oxidized.
• Oxidation state/number indicates the degree of oxidation of an atom in a chemical compound.
• Oxidation states are assigned based on these rules:
  • Free elements have an oxidation state of 0.
  • Monatomic ions have an oxidation state equal to their charge.
  • Oxygen usually has an oxidation state of -2, except in peroxides (-1) or with fluorine.
  • Hydrogen usually has an oxidation state of +1, except in metal hydrides (-1).
  • Fluorine always has an oxidation state of -1.
  • The sum of oxidation states in a neutral molecule or ion equals its charge.
• Balancing redox reactions is achieved through the half-reaction method.
• The half-reaction method separates the oxidation and reduction processes into half-reactions.

Electrochemical Cells

• Electrochemical cells convert chemical energy into electrical energy or vice versa.
• Voltaic (galvanic) cells use spontaneous redox reactions to generate electricity.
• Electrolytic cells use electrical energy to drive non-spontaneous redox reactions (electrolysis).

Voltaic Cells

• A voltaic cell consists of two half-cells, each with an electrode in an electrolyte solution.
• Oxidation occurs at the anode.
• Reduction occurs at the cathode.
• A salt bridge connects the two half-cells and allows ion flow, maintaining charge neutrality.
• Electrons flow from the anode to the cathode through an external circuit, producing an electric current.
• Cell notation represents the voltaic cell, with the anode on the left and the cathode on the right.
• Double vertical lines separate the anode and cathode, indicating the salt bridge.
• Electromotive force (EMF), or cell potential (Ecell), is the potential difference between the cathode and anode.
• Cell potential measures the spontaneity of the redox reaction.
• Standard cell potential (E°cell) is measured under standard conditions (298 K, 1 atm for gases, 1 M for solutions).
• E°cell = E°(cathode) - E°(anode), where E° is the standard reduction potential.
• Standard reduction potentials are given for reduction half-reactions.
• The more positive the E° value, the greater the tendency for the species to be reduced.
• A positive E°cell indicates a spontaneous reaction (galvanic cell).
• A negative E°cell indicates a non-spontaneous reaction.
• Gibbs free energy change (ΔG) relates to the cell potential.
• ΔG = -nFE, where n is the number of moles of electrons transferred and F is the Faraday constant (96485 C/mol).
• Under standard conditions: ΔG° = -nFE°

Electrolytic Cells

• Electrolytic cells use an external voltage to drive non-spontaneous redox reactions.
• Electrolysis uses electrical energy to cause a non-spontaneous chemical reaction.
• Electrolytic cells consist of an anode and a cathode.
• The anode is positive and the cathode is negative - the opposite of voltaic cells.
• Electrolysis of water involves decomposition of water into hydrogen and oxygen gas.
• Electrolysis of molten salts and aqueous solutions are important industrial processes.
• These processes are used for producing metals and other chemicals.
• The amount of substance produced during electrolysis is related to the amount of charge passed through the cell (Faraday's laws of electrolysis).
• mass = (Q * M) / (n * F), where:
  • mass is the mass of substance produced
  • Q is the charge passed
  • M is the molar mass
  • n is the number of electrons transferred
  • F is the Faraday constant

Nernst Equation

• The Nernst equation relates cell potential (Ecell) to standard cell potential (E°cell) and the reaction quotient (Q).
• It is used to calculate cell potentials under non-standard conditions (non-standard concentrations and/or pressures).
• The Nernst equation is: Ecell = E°cell - (RT/nF)lnQ, where:
  • R is the ideal gas constant (8.314 J/mol·K)
  • T is the temperature in Kelvin
  • n is the number of moles of electron transferred
  • F is the Faraday constant
  • Q is the reaction quotient
• At 298 K, the Nernst equation simplifies to: Ecell = E°cell - (0.0592/n)logQ

Batteries

• Batteries apply voltaic cells, providing portable electrical energy.
• Primary batteries are non-rechargeable (e.g., alkaline batteries).
• Secondary batteries are rechargeable (e.g., lead-acid batteries, lithium-ion batteries).
• Fuel cells convert the chemical energy of a fuel (e.g., hydrogen) and an oxidant (e.g., oxygen) into electricity.

Corrosion

• Corrosion is the oxidation of metals, and is often accelerated by electrochemical processes.
• Rusting of iron is a common example of corrosion.
• Corrosion can be prevented by:
  • Protective coatings (e.g., painting, galvanizing)
  • Cathodic protection
  • Using more corrosion-resistant alloys

Explore electrochemistry, the study of chemical reactions causing electron movement. Learn about redox reactions, where oxidation involves electron loss and reduction involves electron gain. Understand oxidizing and reducing agents, and how to assign oxidation states.