Electrolytic Cells PDF

Electrolytic Cells CHEM131 Prepared by: Engr. Kirsten Gail T. Miaga Electrochemistry study of the relationships between electricity and chemical reactions. In electrochemistry, electricity can be generated by movements of electrons from one element to another in a reaction known as r...

Electrolytic Cells CHEM131 Prepared by: Engr. Kirsten Gail T. Miaga Electrochemistry study of the relationships between electricity and chemical reactions. In electrochemistry, electricity can be generated by movements of electrons from one element to another in a reaction known as redox or oxidation- reduction reaction. The conversion takes place in an electrochemical cell, of which there are two main types: Voltaic cells Electrolytic cells Review Mnemonic Oxidation Oxidation and anode both begin Loss of electrons with vowels Occurs at electrode called the anode An ox Reduction Reduction and cathode both Gain of electrons begin with consonants Occurs at electrode called the cathode Red cat Redox reactions Oxidation and reduction occur together Electrolytic Cells An electrolytic cell is an electrochemical cell wherein a non spontaneous redox reaction is made to occur by pumping electrical energy into the system. It is composed of a voltage source battery) connected by wires to two electrodes dipped in an electrolyte. Electrolytic Cells The battery pulls the electrons from the anode, making it positively charged, and pushes it to the cathode, making it negatively charged The two electrodes are dipped in an electrolyte which is either molten or dissolved in a suitable solvent.  To maintain electrical neutrality: the anions in the electrolyte will go to the positively charged anode to give up electrons, undergo oxidation to produce a neutral substance. On the other hand, the cations in the electrolyte will go to the negatively charged cathode to receive electrons, undergo reduction to produce a neutral substance. This process taking place in an electrolytic cell wherein electricity is used to pass through an electrolyte causing the separation of materials is called electrolysis. Electrolysis of Molten Electrolyte When a molten electrolyte is made to undergo electrolysis: There is reduction of a cation at the cathode to a corresponding metal, and Oxidation of an anion at the anion to a corresponding nonmetal Example: The industrial production of sodium metal is done through the electrolysis of molten NaCl. Electrolysis of Aqueous Electrolyte The electrolysis of an aqueous electrolyte is more complicated because of the presence of water which can also undergo electrolysis to produce hydrogen gas and oxygen gas. Decomposition of Aqueous NaF Decomposition of Aqueous KI Electrolysis of Aqueous NaCl Applications of Electrolysis Production of Chemicals Electrolytic Reduction of Metals from their Compounds Extraction and Refining of Metals Electrolytic Refining of Copper Electroplating Quantitative Aspect of Electrolysis The chemical change that occurs in an electrolytic cell is related to the number of moles of electrons that pass through the cell. For every 1 mol of electron passing through the cell, 96480 C is needed. In addition, a quantity of electricity supplied to a cell produces a corresponding amount of substance electrolyzed. This relation was first formulated by Michael Faraday in 1832. Faraday’s Law The mass of the substances formed at the electrodes during electrolysis are directly proportional to the quantity of electricity that passes through the electrolyte. 1 F = 96480 C/mol e− Solving Electrolysis Problems: To solve stoichiometric problems in electrolysis, we must be familiar with the following quantities and units: Q - charge, Coulomb (C) F - farad: 1F = 96480 I - current, ampere(A): I = Q/t number of moles of electrons (ne): This can be obtained from the half reaction equation. For any half reaction, the amount of substance reduced or oxidized in an electrolytic cell is directly proportional to the number of electrons passed into the cell. Example: 1 mol of electron plates out 1 mol of Na metal: 𝑁𝑎+ + 𝑒−→𝑁𝑎 2 mol of electron plates out 1 mol of Cu metal: 𝐶𝑢2+ + 2𝑒−→𝐶𝑢 E - energy, Joule (J): E = QV P - power, watt (W): P = E/t Example: Find the amount of copper formed when 1.25 A of current is made to flow through an aqueous solution of cupric sulfate for 25 minutes and the energy consumed if the process uses a cell voltage of 3.0 V. Example: Chromium metal can be electroplated from an aqueous solution of potassium dichromate. The reduction half reaction is: 𝐶𝑟2𝑂7(𝑎𝑞)2− + 14𝐻(𝑎𝑞)+ + 12𝑒−→ 𝐶𝑟(𝑠) + 7𝐻2𝑂(𝑙) A current of 6.00A and a voltage of 4. 5 V are used in electroplating. How many grams of chromium can be plated if the current is run for 48 minutes? Example: How long will it take to deposit 20g of aluminum ore from an electrolytic cell containing 𝐴𝑙2𝑂3 at a current of 125A? Example: How many grams of sodium and chlorine can be produced in 1.00 hr by the electrolysis of molten sodium chloride in a Downs cell that operates at 5.0 V and 75000 A? How many kilowatt hours of energy are consumed to produce this amount of sodium and chlorine? Anode: 2𝐶𝑙−→𝐶𝑙2+2𝑒− Cathode: 𝑁𝑎+ + 𝑒−→𝑁𝑎