Electrochemistry is all about the movement of electrons in chemical reactions. It's the science behind batteries, fuel cells, and even some industrial processes. Understanding how electrons flow between substances helps us harness electrical energy from chemical reactions.

Electrode potentials are key to predicting how chemicals will behave in these reactions. By comparing potentials, we can figure out which substances will give up electrons and which will grab them. This knowledge is crucial for designing better batteries and more efficient chemical processes.

Electrochemistry Fundamentals

Electrode and cell potentials

• Electrode potential measures the tendency of a chemical species to gain electrons and be reduced
  • Determined relative to the standard hydrogen electrode (SHE) with a potential of 0 V
  • Reduction potentials are more positive and oxidation potentials are more negative
• Cell potential ($E_{cell}$) represents the difference between the reduction potential of the cathode and the oxidation potential of the anode
  • Calculated using the equation: $E_{cell} = E_{cathode} - E_{anode}$
  • Positive $E_{cell}$ indicates a spontaneous redox reaction will occur
  • Cell potential is also referred to as voltage in electrochemical systems
• Electrochemical cells consist of two half-cells, each containing an electrode and an electrolyte
  • Electrons flow from the anode (oxidation) to the cathode (reduction) through an external circuit
  • Salt bridge enables ion transfer to maintain charge balance between the half-cells
  • A galvanic cell (also known as a voltaic cell) is a type of electrochemical cell that produces electricity from spontaneous redox reactions

Oxidant vs reductant strengths

• Standard reduction potentials measure the strength of an oxidizing agent
  • More positive reduction potentials signify stronger oxidizing agents (fluorine, chlorine)
• Standard oxidation potentials measure the strength of a reducing agent
  • More negative oxidation potentials signify stronger reducing agents (alkali metals)
• Comparing electrode potentials reveals that a species with a higher (more positive) reduction potential will oxidize a species with a lower (more negative) reduction potential
  • The species with the lower reduction potential acts as the reducing agent (anode)
• The electrochemical series is a list of elements arranged according to their standard electrode potentials, helping predict the relative strengths of oxidizing and reducing agents

Calculating Cell Potentials and Predicting Spontaneity

Calculations with electrode potentials

• Standard cell potential ($E^°_{cell}$) is calculated using standard reduction potentials ($E^°$) under standard conditions (25°C, 1 M concentrations, 1 atm pressure)
  • $E^°{cell} = E^°{cathode} - E^°_{anode}$
• Spontaneity of redox reactions is determined by the sign of $E^°_{cell}$
  1. If $E^°_{cell} > 0$, the redox reaction is spontaneous under standard conditions
  2. If $E^°_{cell} < 0$, the redox reaction is non-spontaneous under standard conditions
  3. If $E^°_{cell} = 0$, the system is at equilibrium
• Nernst equation relates cell potential to standard cell potential and concentrations of reactants and products
  • $E_{cell} = E^°_{cell} - \frac{RT}{nF} \ln Q$, where $Q$ is the reaction quotient, $R$ is the gas constant, $T$ is the temperature, $n$ is the number of electrons transferred, and $F$ is Faraday's constant
  • Allows for the calculation of cell potentials under non-standard conditions (varying concentrations, temperatures)

Electrochemistry Applications

• Electrochemistry is the study of chemical processes that cause electrons to move, resulting in the interconversion between chemical and electrical energy
• Applications include batteries, fuel cells, and electroplating processes
• Understanding electrode and cell potentials is crucial for designing and optimizing electrochemical systems