Honors Chemistry Unit 11 Review: Redox Reactions & Electrochemistry

Key Concepts

• Redox reactions involve the transfer of electrons between species
• Oxidation is the loss of electrons, while reduction is the gain of electrons
• Oxidation numbers help keep track of electron transfer in redox reactions
• Electrochemical cells convert chemical energy into electrical energy through redox reactions
• Half-reactions represent the oxidation and reduction processes occurring at each electrode
• Standard reduction potentials indicate the relative tendency of species to gain electrons
• Cell potentials can be calculated using standard reduction potentials to determine the spontaneity and direction of redox reactions
• Redox reactions and electrochemistry have numerous real-world applications (batteries, corrosion, electrolysis)

Oxidation and Reduction Basics

• Oxidation involves an increase in the oxidation number of an atom or ion
• Reduction involves a decrease in the oxidation number of an atom or ion
• In a redox reaction, the species being oxidized loses electrons, while the species being reduced gains electrons
• The oxidizing agent is the species that accepts electrons and becomes reduced
• The reducing agent is the species that donates electrons and becomes oxidized
• Oxidation numbers are assigned to atoms in compounds to track electron transfer
  • The sum of oxidation numbers in a neutral compound is always zero
  • The sum of oxidation numbers in an ion is equal to the charge of the ion

Balancing Redox Equations

• Redox equations must be balanced to ensure the conservation of mass and charge
• The half-reaction method is commonly used to balance redox equations
  1. Split the equation into two half-reactions: oxidation and reduction
  2. Balance all atoms except hydrogen and oxygen in each half-reaction
  3. Balance oxygen atoms by adding $H_2O$ to the appropriate side
  4. Balance hydrogen atoms by adding $H^+$ to the appropriate side
  5. Balance the charge in each half-reaction by adding electrons
  6. Multiply the half-reactions by appropriate factors to equalize the number of electrons transferred
  7. Add the balanced half-reactions and cancel out common terms
• Redox equations can also be balanced using the oxidation number method
  • Assign oxidation numbers to each atom in the reactants and products
  • Identify the atoms that change oxidation numbers and determine the number of electrons transferred

Electrochemical Cells

• Electrochemical cells convert chemical energy into electrical energy through spontaneous redox reactions
• A galvanic (voltaic) cell consists of two half-cells connected by a salt bridge or porous membrane
• Each half-cell contains an electrode immersed in an electrolyte solution
• The anode is the electrode where oxidation occurs, while the cathode is the electrode where reduction occurs
• Electrons flow from the anode to the cathode through an external circuit, generating an electric current
• The salt bridge maintains electrical neutrality by allowing the flow of ions between the half-cells
• The direction of electron flow determines the polarity of the electrodes (anode is negative, cathode is positive)

Electrodes and Half-Reactions

• Electrodes are conductors (metals or graphite) that allow the transfer of electrons between the external circuit and the electrolyte
• Half-reactions represent the oxidation and reduction processes occurring at each electrode
• The oxidation half-reaction occurs at the anode, where the species loses electrons
• The reduction half-reaction occurs at the cathode, where the species gains electrons
• The overall cell reaction is the sum of the two half-reactions
• The electrode potential is a measure of the tendency of the half-reaction to occur relative to the standard hydrogen electrode (SHE)
  • The SHE is assigned a potential of 0.00 V by convention
  • Half-reactions with more positive potentials have a greater tendency to occur as reductions

Standard Reduction Potentials

• Standard reduction potentials ($E^0$) are measured under standard conditions (1 M concentrations, 1 atm pressure, 25°C)
• The standard reduction potential table lists half-reactions in order of decreasing tendency to gain electrons
• Species with more positive $E^0$ values are stronger oxidizing agents and are more easily reduced
• Species with more negative $E^0$ values are stronger reducing agents and are more easily oxidized
• The sign of the standard reduction potential indicates the direction of the half-reaction under standard conditions
  • Positive $E^0$ values indicate a spontaneous reduction
  • Negative $E^0$ values indicate a spontaneous oxidation

Calculating Cell Potentials

• The standard cell potential ($E^0_{cell}$) is the difference between the standard reduction potentials of the cathode and anode half-reactions
  • $E^0_{cell} = E^0_{cathode} - E^0_{anode}$
• A positive $E^0_{cell}$ indicates a spontaneous redox reaction under standard conditions
• The Nernst equation relates the cell potential to the concentrations of reactants and products
  • $E_{cell} = E^0_{cell} - \frac{RT}{nF} \ln Q$
  • $R$ 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 cell potential can be used to determine the spontaneity and direction of a redox reaction under non-standard conditions
  • $E_{cell} > 0$ indicates a spontaneous reaction
  • $E_{cell} < 0$ indicates a non-spontaneous reaction

Real-World Applications

• Batteries are portable electrochemical cells that convert chemical energy into electrical energy
  • Primary batteries (disposable) undergo irreversible redox reactions (alkaline, lithium)
  • Secondary batteries (rechargeable) can be recharged by reversing the redox reaction (lithium-ion, nickel-cadmium)
• Fuel cells generate electricity through the oxidation of a fuel (hydrogen, methanol) and the reduction of an oxidant (oxygen)
• Corrosion is an electrochemical process involving the oxidation of a metal (rusting of iron)
  • Cathodic protection prevents corrosion by making the metal the cathode in an electrochemical cell (sacrificial anodes, impressed current)
• Electrolysis is the process of using electrical energy to drive a non-spontaneous redox reaction
  • Used in the production of pure metals (aluminum, copper) and the electroplating of metals
• Electrochemical sensors detect the presence and concentration of specific analytes through redox reactions (glucose sensors, breathalyzers)