Basics-of-Electrochemistry.pdf

Basics of Electrochemistry Schedule 1. 2. 3. 4. 5. Standard electrode potential Electrolysis Laws and equations Cells Literature Standard electrode potential In an electrochemical cell, an electric potential is created between two dissimilar metals. This potential is a measure of the energy per unit...

Basics of Electrochemistry Schedule 1. 2. 3. 4. 5. Standard electrode potential Electrolysis Laws and equations Cells Literature Standard electrode potential In an electrochemical cell, an electric potential is created between two dissimilar metals. This potential is a measure of the energy per unit charge which is available from the oxidation/reduction reactions to drive the reaction. It is customary to visualize the cell reaction in terms of two half-reactions, an oxidation half-reaction and a reduction halfreaction. Electrode Potential 1.The electrode potential cannot be determined in isolation, but in a reaction with some other electrode. 2.The electrode potential depends upon the concentrations of the substances, the temperature, and the pressure in the case of a gas electrode. Thermodynamic conditions 1. Measured against standard hydrogen electrode. 2. Concentration 1 Molar 3. Pressure 1 atmosphere 4. Temperature 25°C Table of some standard potentials Tab. 1 Standard Potentials Redox reactions (reduction & oxidation) Zn (E0 = -0,76 V), Cu (E0 = +0,35 V) Zn0 + Cu2+ = Zn2+ + Cu0 Mg (E0 = -2,34 V), Pb (E0 = -0,13 V) Mg0 + Pb2+ = Mg2+ + Pb0 Redox reactions (reduction & oxidation) Electrolysis In chemistry and manufacturing, electrolysis is a method of using a direct electric current (DC) to drive an otherwise nonspontaneous chemical reaction. Electrolysis is commercially highly important as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell. Electrolysis Pic. 1 Electrolysis of water Faraday’s Laws 1. Faraday’s Law The mass of a substance altered at an electrode during electrolysis is directly proportional to the quantity of electricity transferred at that electrode. Quantity of electricity refers to the quantity of electrical charge, typically measured in coulomb. m=k·I·t=k·Q Faraday’s Laws 2. Faraday’s Law For a given quantity of D.C electricity (electric charge), the mass of an elemental material altered at an electrode is directly proportional to the element's equivalent weight. Faraday’s Laws Faraday's laws can be summarized by where: m is the mass of the substance liberated at an electrode in grams Q is the total electric charge passed through the substance F = 96,485 C mol−1 is the Faraday constant M is the molar mass of the substance z is the valency number of ions of the substance (electrons transferred per ion) Nernst Equation The cell potential can be written: Ecell = oxidation potential + reduction potential In general, a real voltaic cell will differ from the standard conditions, so we need to be able to adjust the calculated cell potential to account for the differences. This can be done with the application of the Nernst equation. Galvaniv cell A galvanic cell, or voltaic cell, named after Luigi Galvani, or Alessandro Volta respectively, is an electrochemical cell that derives electrical energy from spontaneous redox reaction taking place within the cell. It generally consists of two different metals connected by a salt bridge, or individual half-cells separated by a porous membrane. Galvanic & electrolytic cell Lead cell Lead cell reactions Discharging: Cathode (+): PbO2 + 4H+ + 2e- → 2H2O Anode (-): Pb → Pb2+ + 2eCharging: Cathode (-): Pb2+ + 2e- → Pb Anode (+): Pb2+ + 2H2O → PbO2 + 4H+ +2e- Voltage of this cell is near 2V Other Cells Daniell Cell Leclanché Cell Weston Cell Fuel cell Reactions Anode Reaction: 2H2 + 2O2− → 2H2O + 4e− Cathode Reaction: O2 + 4e– → 2O2− Overall Cell Reaction: 2H2 + O2 → 2H2O Literatur - www.wikipedia.org - http://hyperphysics.phyastr.gsu.edu/