Chapter 21a Electrochemistry: The Electrolytic Cell

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1 Electrochemistry Chapter 21a Electrochemistry: The Electrolytic Cell Electrochemical reactions are oxidation-reduction reactions. The two parts of the reaction are physically separated. The oxidation reaction occurs in one cell. The reduction reaction occurs in the other cell. 1 2 Electrochemistry Electrical Conduction There are two kinds electrochemical cells. 1. Electrochemical cells containin spontaneous chemical reactions are called voltaic or alvanic cells. The eneration of electric current from a chemical reaction. 2. Electrochemical cells containin in nonspontaneous chemical reactions are called electrolytic cells. The use of electric current to produce a chemical chane. Metals conduct electric currents well in a process called metallic conduction. In metallic conduction there is electron flow with no atomic motion. Metal atoms chanin oxidation states without movin. E.. Oxidative phosphorylation 3 4 Electrical Conduction Electrodes In ionic or electrolytic conduction ionic motion transports the electrons. Positively chared ions, cations, move toward the neative electrode. Neatively chared ions, anions, move toward the positive electrode. The followin convention for electrodes is correct for either electrolytic or voltaic cells: The cathode is the electrode at which reduction occurs. The cathode is neative in electrolytic cells and positive in voltaic cells. The anode is the electrode at which oxidation occurs. The anode is positive in electrolytic cells and neative in voltaic cells. 5 6

2 Electrodes Electrolytic Cells Inert electrodes do not react with the liquids or products of the electrochemical reaction. Two examples of common inert electrodes are raphite and platinum. Electrical enery is used to force nonspontaneous chemical reactions to occur. The process is called electrolysis. Two examples of commercial electrolytic reactions are: 1. The electroplatin of jewelry and auto parts. 2. The electrolysis of chemical compounds. 7 8 Electrolytic Cells Electrolytic cells consist of: 1. A container for the reaction mixture. 2. Two electrodes immersed in the reaction mixture. 3. A source of direct current. Electrolytic cells uses electrical enery to produce a chemical chane. The electrical enery forces a current throuh a cell that has a neative potential. The electrical enery forces a chemical chane to occur. 9 Fiure 11.19: Voltaic Cell Electrolytic Cell (a) A standard alvanic cell (b) A standard electrolytic cell The cell in (b) has a power source that forces the electrons in the opposite direction from the voltaic cell in (a). 10 Countin Electrons: Coulometry and Faraday s Law of The stoichiometry of electrolysis processes can quantify how much chemical chane occurs with the flow of a iven current for a specific time. Countin Electrons: Coulometry and Faraday s Law of Faraday s Law - The amount of substance underoin chemical reaction at each electrode durin electrolysis is directly proportional to the amount of electricity that passes throuh the electrolytic cell. A faraday is the amount of electricity that reduces one equivalent of a species at the cathode and oxidizes one equivalent of a species at the anode. 1 faraday of electricity e

3 Countin Electrons: Coulometry and Faraday s Law of A coulomb is the amount of chare that passes a iven point when a current of one ampere (A) flows for one second. 1 ampere (amp) = 1 coulomb/second 23-1 faraday e 1 faraday 1.0 mol e 1.0 mol e 96, 485 coulombs Countin Electrons: Coulometry and Faraday s Law of Faraday s Law states that durin electrolysis, one faraday of electricity (96,485 coulombs) reduces and oxidizes, respectively, one equivalent of the oxidizin aent and the reducin aent. This corresponds to the passae of one mole of electrons throuh the electrolytic cell equivalent of oxidizin aent ain of e 23 1 equivalent of reducin aent loss of e Countin Electrons: Coulometry and Faraday s Law of Example: Calculate the mass of palladium produced by the reduction of palladium (II) ions durin the passae of 3.20 amperes of current throuh a solution of palladium (II) sulfate for 30.0 minutes Cathode : Pd + 2 e Pd 1 mol 2 mol 1 mol 106 2(96,485) 106 C 3.20 amp = 3.20 s 1 mol e = 96,485C 60 s 3.20 C 1mole mol Pd 106 Pd? = ( 30.0 min) - = 3.16 Pd min s 96,485C 2mol e mol Pd 15 The of Water Hydroen and oxyen combine spontaneously to form water. The decrease in free enery that accompanies this spontaneous reaction can be used to run fuel cells to produce electricity. The reverse process, which is not spontaneous, requires enery to occur. The formation of oxyen and hydroen ases from water can be forced by electrolysis. 16 The of Water Countin Electrons: Coulometry and Faraday s Law of + Anode reaction 2 H O O + 4 H + 4 e 2 2( ) Cathode reaction 2(2 H O + 2 e H + 2 OH ) 2 2( ) + Cell reaction 6 H2O 2 H 2( ) + O2( ) + 4 H + 4 OH The overall reaction is 2 H O 2 H + O 2 2( ) 2( ) 4 HO 2 17 Example: Calculate the volume of oxyen (measured at STP) produced by the oxidation of water durin the passae of 3.20 amperes of current for 30.0 minutes. + - Anode : 2 H2O O2( ) + 4 H + 4e 2 mol 1 mol 4 mol 4 mol 22.4 LSTP 4( 96,500 C)? LSTP C O2 = 3.20 amp = 3.20 s 1mol e = 96,485C 1.0 mol = 22.4 LSTP 60 s 3.20 C 1mole molo LSTP O 2 ( 30.0 min) - = min s 96,485C 4mole molo2 = L O or 334 ml O STP 2 STP 2 18

4 The of Molten Sodium Chloride Liquid sodium is produced at one electrode. Indicates that the reaction Na + ( ) + e- Na ( ) occurs at this electrode. Is this electrode the anode or cathode? Reduction occurs at the cathode. Gaseous chlorine is produced at the other electrode. Indicates that the reaction 2 Cl - Cl 2() + 2 e - occurs at this electrode. Is this electrode the anode or cathode? Oxidation occurs at the anode. The of Molten Sodium Chloride In all electrolytic cells, electrons are forced to flow from the positive electrode (anode) to the neative electrode (cathode) electrode Na + ains e eneratin liquid Na 0. The of Molten Sodium Chloride Diaram of this electrolytic cell. Na ( ) Na + + e e e e - Cl 2 molten NaCl e - Generator-source of DC + electrode e chloride loses e 2e 2Cl eneratin Cl 2 as The of Molten Sodium Chloride The nonspontaneous redox reaction that occurs is: Anode reaction Cl Cl2( ) + 2 e Cathode reaction ( Na + e Na( ) ) Cell reaction Cl + 2 Na Cl2( ) + 2 Na ( ) Na + + e - Na ( ) cathode reaction Porous barrier 2Cl - Cl 2() + 2e - anode reaction Fiure 11.25: The Downs cell for the electrolysis of molten sodium chloride. Sodium metal is produced by the electrolysis of molten sodium chloride. NaCl is mixed with CaCl 2 to lower the meltin point (from 800 o C to 600 o C). The liquid sodium is drained, cast into blocks and stored in inert solvents. The of Aqueous Sodium Chloride In this electrolytic cell, hydroen as is produced at one electrode. The aqueous solution becomes basic near this electrode. What reaction is occurrin at this electrode? Gaseous chlorine is produced at the other electrode

5 The of Aqueous Potassium Chloride What reaction is occurrin at this electrode? These experimental facts lead us to the followin nonspontaneous electrode reactions: Anode reaction 2 Cl Cl + 2 e 2( ) Cathode reaction 2 H 2O + 2 e H 2( ) + 2 OH Cell reaction 2 Cl + 2 H 2O H 2( ) + Cl2( ) + 2 OH + Na is a spectator ion. Note that water is electrolyzed! 25 The of Aqueous Potassium Chloride Cell diaram e - flow H 2 as pole of battery + pole of battery Battery, a source of direct current e - flow -electrode 2 H 2 O + 2e - H 2 () + 2 OH - cathode reaction aqueous NaCl + electrode Cl 2 as 2Cl - Cl 2 () + 2e - anode reaction 26 Electrolytic Cells Commercial Applications of Electrolytic Cells In all electrolytic cells the most easily reduced species is reduced and the most easily oxidized species is oxidized. Electrolytic Refinin and Electroplatin of Metals Impure metallic copper can be purified electrolytically to 100% pure Cu. The impurities commonly include some active metals plus less active metals such as: A, Au, and Pt. The cathode is a thin sheet of copper metal connected to the neative terminal of a direct current source. The anode is lare impure bars of copper Commercial Applications of Electrolytic Cells Commercial Applications of Electrolytic Cells The electrolytic solution is CuSO 4 and H 2 SO 4 The impure Cu dissolves to form Cu 2+. The Cu 2+ ions are reduced to Cu at the cathode. 0 2 ( ) ( s) ( aq) 2+ 0 ( ) ( ) + aq ( s) Anode impure Cu Cu + 2e Cathode very pure Cu 2e Cu Net rxn. No net rxn. + Any active metal impurities are oxidized to cations that are more difficult to reduce than Cu 2+. This effectively removes them from the Cu metal Zn Zn + 2e 0 2+ Fe Fe + 2e And so forth for other active metals 29 30

6 Commercial Applications of Electrolytic Cells Metal Platin Objects can be plated by makin a particular object a cathode in a tank with ions of the platin metal. Fiure 11.24: Schematic of the electroplatin of a spoon. The spoon is the cathode and is plated out by the A + ions that are released from the solid silver bar that is the anode. A salt bride is not required because A + are at actin at both electrodes Copper Platin Commercial Applications of Electrolytic Cells The less active metals are not oxidized and precipitate to the bottom of the cell. These metal impurities can be isolated and separated after the cell is disconnected. Some common metals that precipitate include: A, Au, Pt, Pd ( Se, Te) Corrosion Corrosion Metallic corrosion is the oxidation-reduction reactions of a metal with atmospheric components such as CO 2, O 2, and H 2 O. Metals corrode because they oxidize easily. Many common metals that are used for structural and decorative purposed have standard reduction potentials that are more neative and oxyen. Corrosion of iron The importance of steel in many of our structures, controllin corrosion is a very important issue. The corrosion mechanism involves electrochemical processes Fe + 3 O2 2 Fe2O3 ( overall reaction) The reaction occurs rapidly at exposed points

7 Corrosion steel The surface of steel is not uniform. The chemical composition of steel is not a homoeneous mixture. Stress points are produced on the surface due to physical strains. At these stress areas, iron can be more easily oxidized in some reions than in other reions. o Oxidized areas act as anodes o The other areas act as cathodes Fe 2+ ions travel throuh the surface moisture to the reion actin as a cathode. In the reion of the cathode, the Fe 2+ ions react with O 2 to form rust. The moisture acts as a salt bride in the process of corrosion. Without moisture, steel does not rust Some examples of corrosion protection. 1. Plate a metal with a thin layer of a less active (less easily oxidized) metal. " Tin plate" or " chromium plate" for steel. 2. Galvanizin, the coatin of steel with zinc, provides a more active metal on the exterior. The thin coat of Zn must be oxidized before Fe beins to rust. Zinc Steel Connect the metal to a sacrificial anode, a piece of a more active metal. Soil pipes and ship hulls have M and Zn on the exterior as sacrificial anodes. Manesium is easily oxidized; protectin the iron from oxidation

8 4. Allow a protective film to form naturally Al + 3 O 2 Al O Al O forms a hard, transparent film on exterior of aluminum foil. 5. Paint or coat with a polymeric material such as plastic or ceramic. Steel bathtubs are coated with ceramic End of Chapter 11b Electrochemistry is an important part of the electronics industry. 45

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