Inorganic Chemistry I Unit 7 Review: Redox Reactions in Electrochemistry

Key Concepts and Definitions

• Redox reactions involve the transfer of electrons between species
• Oxidation is the loss of electrons and reduction is the gain of electrons (OIL RIG mnemonic)
• Oxidizing agents are reduced and reducing agents are oxidized in a redox reaction
• Electrochemistry studies the interconversion of electrical and chemical energy in redox reactions
• Electrochemical cells convert chemical energy into electrical energy (galvanic cells) or vice versa (electrolytic cells)
• Electrodes are conductors where redox reactions occur in an electrochemical cell
  • Anode is the electrode where oxidation occurs
  • Cathode is the electrode where reduction occurs
• Electrolytes are ionic compounds that conduct electricity when dissolved in a solvent (water)

Oxidation and Reduction Basics

• Oxidation states (oxidation numbers) help keep track of electron transfer in redox reactions
• Oxidation involves an increase in the oxidation state of an atom or ion
• Reduction involves a decrease in the oxidation state of an atom or ion
• Rules for assigning oxidation states:
  • Free elements have an oxidation state of zero
  • Monatomic ions have an oxidation state equal to their charge
  • Hydrogen has an oxidation state of +1 (except in metal hydrides, where it is -1)
  • Oxygen has an oxidation state of -2 (except in peroxides, where it is -1, and in compounds with fluorine)
  • The sum of oxidation states in a neutral compound is zero
  • The sum of oxidation states in a polyatomic ion equals the charge of the ion
• Redox reactions can be classified as combination, decomposition, displacement, or disproportionation reactions

Electrochemical Cells and Their Components

• Galvanic (voltaic) cells spontaneously convert chemical energy into electrical energy
  • Examples include batteries (dry cells, lead-acid batteries) and fuel cells
• Electrolytic cells use electrical energy to drive non-spontaneous redox reactions
  • Examples include electroplating, electrolysis of water, and the Hall-Héroult process for aluminum production
• Half-cells are the oxidation and reduction compartments of an electrochemical cell
• Salt bridges or porous membranes maintain electrical neutrality and allow ion flow between half-cells
• Electrodes can be active (participate in the redox reaction) or passive (provide a surface for electron transfer)
• The anode is the site of oxidation and the cathode is the site of reduction in an electrochemical cell
• The direction of electron flow in the external circuit is from anode to cathode
• The direction of conventional current is opposite to the flow of electrons (cathode to anode)

Electrochemical Potentials and Standard Reduction Potentials

• Electrochemical potential (cell potential) is the driving force for the redox reaction in an electrochemical cell
• Standard reduction potentials ($E^0$) are measured relative to the standard hydrogen electrode (SHE) under standard conditions (1 M concentrations, 1 atm pressure, 25°C)
• The standard cell potential ($E^0_\text{cell}$) is the difference between the standard reduction potentials of the half-reactions
  • $E^0_\text{cell} = E^0_\text{cathode} - E^0_\text{anode}$
• A positive $E^0_\text{cell}$ indicates a spontaneous redox reaction under standard conditions
• The Nernst equation relates the cell potential to the concentrations (or partial pressures) of the reactants and products
  • $E_\text{cell} = E^0_\text{cell} - \frac{RT}{nF} \ln Q$, where $Q$ is the reaction quotient
• Concentration cells have the same species in both half-cells but at different concentrations

Balancing Redox Equations

• Redox equations must be balanced for both mass and charge
• Half-reaction method for balancing redox equations in acidic or basic solution:
  1. Write separate half-reactions for oxidation and reduction
  2. Balance all atoms except H and O
  3. Balance O atoms by adding H2O
  4. Balance H atoms by adding H+
  5. Balance charge by adding electrons
  6. Multiply the half-reactions by appropriate factors to equalize the number of electrons transferred
  7. Add the half-reactions and cancel common terms
  8. For basic solutions, add OH- to neutralize H+ and simplify the equation
• Oxidation number method for balancing redox equations:
  1. Assign oxidation numbers to all atoms
  2. Identify the atoms that change oxidation number
  3. Write skeleton equations for the oxidation and reduction half-reactions
  4. Balance the half-reactions and combine them as in the half-reaction method

Electrochemical Series and Predicting Reactions

• The electrochemical series arranges elements in order of their standard reduction potentials
• 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
• To predict the spontaneity of a redox reaction:
  • Identify the possible oxidation and reduction half-reactions
  • Look up the standard reduction potentials in a table
  • Calculate the standard cell potential ($E^0_\text{cell}$)
  • If $E^0_\text{cell} > 0$, the reaction is spontaneous under standard conditions
• The electrochemical series can also predict the reactivity of metals (corrosion) and the feasibility of metal displacement reactions

Applications in Industry and Technology

• Batteries and fuel cells:
  • Primary (non-rechargeable) batteries (alkaline, lithium, silver oxide)
  • Secondary (rechargeable) batteries (lead-acid, nickel-cadmium, lithium-ion)
  • Fuel cells (hydrogen, methanol, solid oxide)
• Corrosion and its prevention:
  • Galvanic corrosion occurs when dissimilar metals are in electrical contact in the presence of an electrolyte
  • Cathodic protection uses sacrificial anodes (zinc, magnesium) or impressed current to prevent corrosion of steel structures
  • Passivation forms a protective oxide layer on metal surfaces (stainless steel, aluminum)
• Electroplating and electrorefining:
  • Electroplating deposits a thin layer of metal onto a conductive surface for decoration, protection, or improved properties
  • Electrorefining purifies metals by selective electrodeposition (copper, gold, silver)
• Electrolysis and electrochemical synthesis:
  • Electrolysis of water produces hydrogen and oxygen gases
  • Electrolysis of brine (NaCl solution) produces chlorine gas and sodium hydroxide (chlor-alkali process)
  • Electrochemical synthesis of organic compounds (adiponitrile, sebacic acid) and inorganic compounds (persulfates, perchlorates)

Common Mistakes and How to Avoid Them

• Confusing oxidation and reduction:
  • Remember: OIL RIG (Oxidation Is Loss, Reduction Is Gain) of electrons
  • Oxidation increases the oxidation state, reduction decreases the oxidation state
• Incorrectly assigning oxidation states:
  • Follow the rules for assigning oxidation states consistently
  • Check that the sum of oxidation states equals the charge of the species
• Misidentifying the anode and cathode:
  • The anode is always the site of oxidation, and the cathode is always the site of reduction
  • In galvanic cells, the anode is the negative electrode, and the cathode is the positive electrode
  • In electrolytic cells, the anode is the positive electrode, and the cathode is the negative electrode
• Misbalancing redox equations:
  • Balance both mass and charge in the half-reactions
  • Equalize the number of electrons transferred between the half-reactions
• Misinterpreting cell notation:
  • The anode is always written on the left, and the cathode is always written on the right
  • The single vertical line represents a phase boundary (solid/liquid or liquid/liquid)
  • The double vertical line represents a salt bridge or porous membrane
• Misusing the Nernst equation:
  • Use the correct sign convention for the cell potential and the reaction quotient
  • Remember that the Nernst equation assumes activities are equal to concentrations (or partial pressures for gases)
• Neglecting the effects of concentration, temperature, and pressure on cell potentials:
  • Use the Nernst equation to account for non-standard conditions
  • Consider the effects of temperature on the standard cell potential and the reaction quotient
  • For gas-phase reactions, use partial pressures instead of concentrations in the Nernst equation