Chapter 20: Electrochemistry

Electrochemistry , titled Electrochemistry, provides a detailed analysis of both non-spontaneous chemical reactions driven by electricity and the spontaneous generation of electrical potential from chemical reactions. A foundational concept is electrolysis, the decomposition of a compound using an electric current, typically employed for purifying metals or extracting highly reactive elements. Key components include the electrolyte (a molten ionic compound or concentrated aqueous solution), the anode (positive electrode, site of oxidation or electron loss), and the cathode (negative electrode, site of reduction or electron gain). When predicting the products of electrolysis in aqueous solutions, the outcome depends on the electrolyte's state (molten or aqueous), the relative positions of competing ions in the redox series (electrode potential), and the concentration of the ions. Quantitatively, the chapter establishes that the amount of substance produced is proportional to the quantity of electricity (charge) passed, calculated using the relationship Q=It (charge = current times time). This ties into the Faraday constant (F), which represents the charge carried by one mole of electrons, related to the Avogadro constant (L) and the charge of a single electron (e) by the equation F=Le. This principle enables the determination of L through electrolytic methods. The discussion moves into electrode potentials (E), which arise from the redox equilibrium established between an element and its ions in solution. These potentials are measured relative to the Standard Hydrogen Electrode (SHE), which is arbitrarily set at 0.00 V under standard conditions (1.00 mol dm⁻³ ion concentration, 25 degrees Celsius, 101 kPa gas pressure). When two half-cells are connected via a wire and a salt bridge (which maintains ionic balance) to form an electrochemical cell, the standard cell potential ($E^\text{\textcurrency}$ cell) can be calculated as the difference between the $E^\text{\textcurrency}$ values of the two half-cells. Electrode potentials are a powerful tool for deducing the relative strength of species as oxidizing or reducing agents, since a more positive $E^\text{\textcurrency}$ indicates a greater tendency for reduction. The electron flow in the external circuit moves from the electrode with the more negative $E^\text{\textcurrency}$ (negative pole) to the one with the more positive $E^\text{\textcurrency}$ (positive pole). A reaction is predicted to be feasible if the overall cell potential ($E^\text{\textcurrency}$ cell) is positive. Finally, the Nernst equation is introduced to quantitatively predict how electrode potential varies when ion concentrations deviate from standard conditions. It is important to remember that $E^\text{\textcurrency}$ values predict thermodynamic feasibility but not the reaction rate, which can sometimes be impractically slow.