7.1 Principles of electrochemistry

Electrochemistry is all about the dance between electricity and chemical reactions. It's like a party where electrons are the guests, moving between different chemical compounds and creating energy in the process.

In this section, we'll explore how this electron exchange works, what drives it, and how we can use it to measure stuff. From batteries to sensors, electrochemistry is the hidden hero in many everyday technologies.

Fundamentals of Electrochemistry

Redox Reactions and Electron Transfer

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• Electrochemistry studies the interconversion of electrical and chemical energy through oxidation-reduction (redox) reactions
• Redox reactions involve the transfer of electrons between chemical species
  • Oxidation is the loss of electrons
  • Reduction is the gain of electrons
• The anode is the electrode where oxidation occurs, while the cathode is the electrode where reduction takes place

Electrodes and Electrolytes

• Electrodes are conductors that allow the flow of electrons to and from the electrolyte in an electrochemical cell
  • Electrodes are typically made of metals (platinum, gold, silver) or carbon (graphite, glassy carbon)
  • The choice of electrode material depends on the specific application and the potential window required
• Electrolytes are substances that dissociate into ions when dissolved in a solvent, enabling the conduction of electricity through the solution
  • Common electrolytes include aqueous solutions of salts (NaCl, KCl), acids (HCl, H2SO4), and bases (NaOH, KOH)
  • The concentration and composition of the electrolyte affect the conductivity and the electrochemical reactions occurring at the electrodes

Electrical Potential and Chemical Reactions

Standard Electrode Potential and the Nernst Equation

• Electrical potential, or voltage, is the driving force for the flow of electrons in an electrochemical cell
• The standard electrode potential (E°) measures the tendency of a chemical species to be reduced
  • More positive E° values indicate a greater tendency to be reduced
  • Examples of standard electrode potentials: Cu^2+/Cu (+0.34 V), Ag^+/Ag (+0.80 V), Zn^2+/Zn (-0.76 V)
• The Nernst equation relates the electrode potential to the standard electrode potential and the concentrations of the oxidized and reduced species in the cell
  • $E = E° - \frac{RT}{nF} \ln \frac{[Red]}{[Ox]}$, where R is the gas constant, T is the temperature, n is the number of electrons transferred, and F is Faraday's constant

Cell Potential and Gibbs Free Energy

• The difference in electrode potentials between the anode and cathode determines the cell potential (E_cell) and the direction of electron flow
  • $E_{cell} = E_{cathode} - E_{anode}$
  • A positive E_cell indicates a spontaneous reaction, while a negative E_cell indicates a non-spontaneous reaction
• The Gibbs free energy change (ΔG) of a redox reaction is directly related to the cell potential
  • $ΔG = -nFE_{cell}$, where n is the number of electrons transferred and F is Faraday's constant
  • Negative ΔG values indicate a spontaneous reaction, while positive ΔG values indicate a non-spontaneous reaction

Electrochemical Cells and Applications

Components and Principles of Electrochemical Cells

• An electrochemical cell consists of two half-cells, each containing an electrode and an electrolyte, connected by a salt bridge or a porous membrane
  • The salt bridge or porous membrane allows the flow of ions between the half-cells to maintain charge balance while preventing the mixing of the electrolytes
  • Examples of salt bridge materials include agar gel with KCl or KNO3, and porous glass frits
• The half-cell reactions and the overall cell reaction can be written to describe the electrochemical processes occurring in the cell
  • Example: In a Daniell cell, the half-cell reactions are Cu^2+ + 2e^- → Cu (cathode) and Zn → Zn^2+ + 2e^- (anode), with the overall cell reaction being Cu^2+ + Zn → Cu + Zn^2+

Applications in Analytical Chemistry

• Electrochemical cells can be used to measure the concentration of analytes in solution through potentiometric or amperometric methods
• Potentiometric methods measure the potential difference between a reference electrode and an indicator electrode to determine the concentration of a specific ion
  • Ion-selective electrodes (ISEs) are commonly used in potentiometric methods (pH electrode, fluoride ISE)
  • The potential difference is related to the analyte concentration through the Nernst equation
• Amperometric methods measure the current produced by the oxidation or reduction of an analyte at a working electrode
  • Voltammetry and coulometry are examples of amperometric methods
  • The current is proportional to the analyte concentration and can be used for quantitative analysis

Galvanic vs Electrolytic Cells

Galvanic Cells

• Galvanic cells, also known as voltaic cells, convert chemical energy into electrical energy through spontaneous redox reactions
• In a galvanic cell, the anode undergoes oxidation, and the cathode undergoes reduction, with electrons flowing from the anode to the cathode through an external circuit
  • Example: In a lead-acid battery, the anode is lead (Pb), and the cathode is lead dioxide (PbO2), with the overall reaction being Pb + PbO2 + 2H2SO4 → 2PbSO4 + 2H2O
• Galvanic cells are used in potentiometric methods, such as the determination of pH using a glass electrode and a reference electrode
  • The glass electrode acts as the indicator electrode, with a potential that depends on the pH of the solution
  • The reference electrode (Ag/AgCl or calomel) provides a stable and known potential for comparison

Electrolytic Cells

• Electrolytic cells convert electrical energy into chemical energy by using an external power source to drive a non-spontaneous redox reaction
• In an electrolytic cell, the anode undergoes oxidation, and the cathode undergoes reduction, with electrons flowing from the external power source to the cathode and then to the anode through the electrolyte
  • Example: In the electrolysis of water, the anode reaction is 2H2O → O2 + 4H^+ + 4e^-, and the cathode reaction is 2H^+ + 2e^- → H2, with the overall reaction being 2H2O → 2H2 + O2
• Electrolytic cells are used in amperometric methods, such as the electrolysis of water to produce hydrogen and oxygen gases or the electrodeposition of metals
  • In the electrodeposition of copper, Cu^2+ ions are reduced to Cu metal at the cathode, while an inert anode (platinum) is used to complete the circuit
  • The amount of Cu deposited can be calculated using Faraday's laws of electrolysis, which relate the mass of the substance deposited or liberated to the quantity of electricity passed