🔌Electrochemistry Unit 2 – Redox Reactions and 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 relationship between electrical energy and chemical change
  • Includes the study of batteries, fuel cells, and electrolysis
• Electrodes are conductors where reduction and oxidation half-reactions occur
  • Anode is the site of oxidation and cathode is the site of reduction
• Electrolytes are substances that conduct electricity when dissolved in water (acids, bases, salts)
• Standard reduction potentials ($E^0$) measure the tendency of a species to gain electrons and be reduced under standard conditions (1 M, 1 atm, 25°C)

Oxidation and Reduction Basics

• Oxidation numbers (oxidation states) represent the degree of oxidation of an atom in a compound
  • Oxidation involves an increase in oxidation number
  • Reduction involves a decrease in oxidation number
• Rules for assigning oxidation numbers:
  • Free elements have an oxidation number of 0 (Na, H2, O2)
  • Monatomic ions have an oxidation number equal to their charge (Na+ is +1, Cl- is -1)
  • Oxygen is usually -2, except in peroxides (H2O2) where it is -1 and in compounds with fluorine
  • Hydrogen is usually +1, except in metal hydrides (NaH) where it is -1
  • The sum of oxidation numbers in a neutral compound is 0
• Oxidizing agents are chemical species that oxidize other substances and are reduced in the process (gain electrons)
  • Examples include O2, H2O2, halogens (F2, Cl2), and metal ions with high oxidation states (MnO4-, Cr2O72-)
• Reducing agents are chemical species that reduce other substances and are oxidized in the process (lose electrons)
  • Examples include metals (Na, Mg, Zn), H2, and compounds with low oxidation states (H2S, HI)

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 (e-)
  6. Multiply half-reactions by appropriate factors to equalize electrons transferred
  7. Add half-reactions and cancel common terms
  8. For basic solutions, add OH- to neutralize H+ and simplify
• Oxidation number method for balancing redox equations:
  1. Assign oxidation numbers to all atoms
  2. Identify atoms that change oxidation number
  3. Write skeleton equation showing species that are oxidized and reduced
  4. Balance atoms that do not change oxidation number
  5. Balance atoms that change oxidation number by adding coefficients
  6. Verify that the equation is balanced for both mass and charge

Electrochemical Cells

• Electrochemical cells convert chemical energy into electrical energy (voltaic cells) or vice versa (electrolytic cells)
• Voltaic (galvanic) cells spontaneously generate electricity from redox reactions
  • Consist of two half-cells connected by a salt bridge or porous membrane
  • Anode is the site of oxidation and cathode is the site of reduction
  • Electrons flow from anode to cathode through an external circuit
  • Examples include batteries (dry cells, lead-acid) and fuel cells
• Electrolytic cells use an external power source to drive nonspontaneous redox reactions
  • Reduction occurs at the cathode and oxidation at the anode
  • Used in electroplating, electrolysis, and electrochemical synthesis
• Salt bridge maintains electrical neutrality by allowing ions to migrate between half-cells
  • Prevents mixing of solutions and enables electron transfer
• Cell notation (cell diagram) represents the components of an electrochemical cell
  • Anode | Anode solution || Cathode solution | Cathode
  • Example: Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

Electrode Potentials and Standard Reduction Potentials

• Electrode potential is the measure of the tendency of a half-reaction to occur at an electrode
• Standard electrode potentials ($E^0$) are measured under standard conditions (1 M, 1 atm, 25°C)
  • Defined relative to the standard hydrogen electrode (SHE) which has an assigned potential of 0.00 V
• Standard reduction potentials are tabulated for various half-reactions
  • More positive $E^0$ values indicate a greater tendency to be reduced (stronger oxidizing agents)
  • More negative $E^0$ values indicate a greater tendency to be oxidized (stronger reducing agents)
• Nernst equation relates the cell potential ($E_{cell}$) to the standard cell potential ($E^0_{cell}$) and the concentrations of reactants and products:
  • $E_{cell} = E^0_{cell} - \frac{RT}{nF} \ln Q$
  • $R$ is the gas constant, $T$ is temperature, $n$ is the number of electrons transferred, $F$ is Faraday's constant, and $Q$ is the reaction quotient
• Cell potential is related to the Gibbs free energy change ($\Delta G$) and the equilibrium constant ($K$):
  • $\Delta G = -nFE_{cell}$
  • $\Delta G^0 = -nFE^0_{cell} = -RT \ln K$

Electrochemical Series and Predicting Reactions

• The electrochemical series arranges elements in order of their standard reduction potentials
  • Elements at the top (more positive $E^0$) are strong oxidizing agents and are easily reduced
  • Elements at the bottom (more negative $E^0$) are strong reducing agents and are easily oxidized
• The electrochemical series can be used to predict the spontaneity and direction of redox reactions
  • A species will spontaneously reduce another species with a lower (more positive) reduction potential
  • The species with the higher (more negative) reduction potential will be oxidized
• To determine the cell potential and spontaneity of a redox reaction:
  1. Identify the half-reactions and write them as reductions
  2. Look up the standard reduction potentials for each half-reaction
  3. Reverse the half-reaction with the more negative $E^0$ (this will be the oxidation)
  4. Calculate the cell potential: $E^0_{cell} = E^0_{cathode} - E^0_{anode}$
  5. If $E^0_{cell} > 0$, the reaction is spontaneous; if $E^0_{cell} < 0$, the reaction is nonspontaneous
• The electrochemical series can also predict the reactivity of metals
  • Metals with more negative reduction potentials are stronger reducing agents and more reactive
  • Metals can displace less reactive metals from their compounds in solution

Applications in Industry and Technology

• Batteries are portable electrochemical cells that convert chemical energy into electrical energy
  • Primary batteries are single-use and irreversible (alkaline, lithium)
  • Secondary batteries are rechargeable and reversible (lead-acid, lithium-ion)
• Fuel cells generate electricity from the oxidation of fuels (hydrogen, methanol, ethanol)
  • Efficient and environmentally friendly alternative to combustion engines
  • Applications in transportation, stationary power generation, and portable electronics
• Electroplating uses electrolytic cells to deposit a thin layer of metal onto a surface
  • Improves appearance, corrosion resistance, and wear resistance
  • Examples include chrome plating, gold plating, and electroless nickel plating
• Electrolysis is the decomposition of a substance using an electric current
  • Used in the production of metals (aluminum, copper, sodium)
  • Electrolysis of water produces hydrogen and oxygen gases
• Corrosion is an electrochemical process that involves the oxidation of metals
  • Rusting of iron is a common example: $4Fe + 3O_2 + 6H_2O \rightarrow 4Fe(OH)_3$
  • Prevention methods include cathodic protection, sacrificial anodes, and protective coatings
• Electrochemical sensors and biosensors detect specific analytes using redox reactions
  • Examples include glucose sensors, oxygen sensors, and pH meters
  • Applications in medical diagnostics, environmental monitoring, and food safety

Common Mistakes and How to Avoid Them

• Confusing oxidation and reduction
  • Remember: OIL RIG (Oxidation Is Loss, Reduction Is Gain) of electrons
• Incorrectly assigning oxidation numbers
  • Follow the rules for assigning oxidation numbers and practice with many examples
• Forgetting to balance both mass and charge in redox equations
  • Use the half-reaction or oxidation number method to systematically balance equations
• Reversing the direction of electron flow in electrochemical cells
  • Electrons always flow from the anode (oxidation) to the cathode (reduction) in the external circuit
• Using incorrect signs for cell potentials or reduction potentials
  • Cell potentials are positive for spontaneous reactions and negative for nonspontaneous reactions
  • Reduction potentials are defined as reductions (gaining electrons) and become more positive for stronger oxidizing agents
• Neglecting to consider concentration effects on cell potentials
  • Use the Nernst equation to calculate cell potentials at non-standard conditions
• Misinterpreting the electrochemical series
  • Species at the top (more positive $E^0$) are strong oxidizing agents and are easily reduced
  • Species at the bottom (more negative $E^0$) are strong reducing agents and are easily oxidized
• Not recognizing the limitations of standard reduction potentials
  • Standard reduction potentials are measured under specific conditions (1 M, 1 atm, 25°C) and may not accurately predict reactivity under different conditions
  • Kinetic factors, such as activation energy and surface area, can also influence the rate and extent of redox reactions