Electro Chemistry - Study Notes
Chapter Summary
This chapter investigates the dynamic relationship between electrical energy and chemical transformations. It provides a comprehensive look at how chemical energy is converted into electrical work and how external current can drive non-spontaneous reactions. The content covers the principles of ionic conductivity, the mechanics of galvanic and electrolytic cells, and the quantitative relationships defined by laws from Faraday and Nernst. It also addresses practical applications such as commercial batteries, fuel cells, and the electrochemical nature of metal corrosion.
Learning Objectives
- Explain the conductivity and resistivity of electrolytic solutions.
- Define terms like molar and equivalent conductivity and understand their dependence on concentration.
- Apply Kohlrausch's law to determine limiting molar conductivity.
- Represent electrochemical cells using proper IUPAC notation.
- Derive and utilize the Nernst equation to calculate potential.
- State and apply Faraday's laws of electrolysis in various scenarios.
- Describe the construction and function of primary and secondary batteries.
- Explain the mechanism of corrosion and methods for its prevention.
Key Concepts and Definitions
- Conductivity (κ): The measure of a solution's ability to conduct electricity, defined as the reciprocal of resistivity.
- Molar Conductivity (Λm): The conducting power of all ions produced by dissolving one mole of electrolyte in a given volume of solution.
- Galvanic Cell: A device that produces electrical energy from spontaneous chemical redox reactions, also known as a voltaic cell.
- Electrolytic Cell: A device where electrical energy is consumed to drive a non-spontaneous chemical change.
- Standard Hydrogen Electrode (SHE): A reference electrode with an assigned potential of zero volts used to measure other electrode potentials.
- Electrochemical Series: A sequence of metals and non-metals arranged according to their standard reduction potentials.
Worked Methods
Calculating Standard Cell Potential
Determine the standard potential of a cell by subtracting the standard reduction potential of the anode from that of the cathode. Ensure all potentials used are reduction potentials. A positive result indicates a spontaneous reaction.
Applying the Nernst Equation
Use the Nernst equation to find the cell potential under non-standard conditions. At 298K, apply the formula E = E° - (0.0591 / n) * log Q. Calculate the reaction quotient (Q) based on the concentrations of products and reactants and use the number of electrons (n) exchanged in the balanced reaction.
Common Exam Traps
- Incorrectly assigning signs to electrodes: In a galvanic cell, the anode is negative, while in an electrolytic cell, it is positive.
- Confusing specific conductance with molar conductivity; remember that while specific conductance decreases with dilution, molar conductivity increases.
- Neglecting to convert time from minutes or hours into seconds when using Faraday's laws of electrolysis.
Exam Tips
- Memorize the electrochemical series to quickly predict which metals will act as anodes when paired with others.
- Understand that for the Nernst equation, pure solids and liquids have an activity (concentration) of one and are excluded from the reaction quotient.
- Practice re-writing the recharging equations for secondary batteries like the lead storage cell, as these are the exact reverse of the discharge reactions.