Electrochemistry – part 2 Chapter 17 Electrochemistry – part 2

Led-acid storage batteries Consists of six cells wired in series. Each cell contains a porous lead anode and a lead oxide cathode, both immersed in sulfuric acid. An electric current is drawn from the battery, both the anode and cathode become coated with PbSO4(s) Can be recharged by running electric current through it in reverse direction

Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Batteries Lead Storage Battery Anode: Pb(s) + HSO4(aq) PbSO4(s) + H+(aq) + 2e Cathode: PbO2(s) + 3H+(aq) + HSO4(aq) + 2e PbSO4(s) + 2H2O(l) Overall: Pb(s) + PbO2(s) + 2H+(aq) + 2HSO41(aq) 2PbSO4(s) + 2H2O(l) Copyright © 2011 Pearson Prentice Hall, Inc.

Dry-Cell Batteries Zinc acts as the anode and a graphite rod immersed in a moist, slightly acidic pasted of MnO2 and NH4Cl acts a cathode.

Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Batteries Dry-Cell Batteries Alkaline cell Anode: Zn(s) + 2OH(aq) ZnO(s) + H2O(l) + 2e Cathode: 2MnO2(s) + H2O(l) + 2e Mn2O3(s) + 2OH(aq) Copyright © 2011 Pearson Prentice Hall, Inc.

Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Batteries Dry-Cell Batteries Leclanché cell Anode: Zn(s) Zn2+(aq) + 2e Cathode: 2MnO2(s) + 2NH4+(aq) + 2e Mn2O3(s) + 2NH3(aq)+ H2O(l) Copyright © 2011 Pearson Prentice Hall, Inc.

Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Batteries Nickel-Cadmium (“ni-cad”) Batteries Anode: Cd(s) + 2OH(aq) Cd(OH)2(s) + 2e Cathode: NiO(OH)(s) + H2O(l) + e Ni(OH)2(s) + OH(aq) Nickel-Metal Hydride (“NiMH”) Batteries Anode: MHab(s) + OH(aq) M(s) + H2O(l) + e Cathode: NiO(OH)(s) + H2O(l) + e Ni(OH)2(s) + OH(aq) Overall: MHab(s) + NiO(OH)(s) M(s) + Ni(OH)2(s) Copyright © 2011 Pearson Prentice Hall, Inc.

Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Batteries Lithium and Lithium Ion Batteries Lithium Anode: xLi(s) xLi+(soln) + xe Cathode: MnO2(s) + xLi+(soln) + xe LixMnO2(s) Lithium Ion Notice where lithium is on the standard reduction table. Anode: LixC6(s) xLi+(soln) + 6C(s) + xe Cathode: Li1-xCoO2(s) + xLi+(soln) + xe LiCoO2(s) Copyright © 2011 Pearson Prentice Hall, Inc.

Fuel cells Like batteries, but the reactants must be constanly replenished. Normal batteries los their ability to generate voltage with use because the reactants become depleted as electric current is drawn from the battery. In fuel cell, the reactant – the fuel-constanly flow through the battery, generating electric current as they undergo redox reaction.

Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Fuel Cells Hydrogen-Oxygen Fuel Cell Copyright © 2011 Pearson Prentice Hall, Inc.

Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry Corrosion 4/24/2017 2:41:42 PM Corrosion: The oxidative deterioration of a metal. Copyright © 2011 Pearson Prentice Hall, Inc.

Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Corrosion Moisture must be present for rusting to occur Additional electrolytes promote more rusting Such as NaCl, on the surface of iron because it enhances current flow The presence of acid promotes rusting. (H+ ions are involved in the reduction of oxygen, lower pH enhances the cathodic reaction and leads to faster rusting. Iron rust doesn’t protect iron since the rust is too porous. Copyright © 2011 Pearson Prentice Hall, Inc.

Preventing Corrosion Keep dry Coat the iron with a substance that is impervious to water Painting Placing a sacrificial electrode in electrical contact with the iron. For some metals, oxidation protects the metal (aluminum, chromium, magnesium, titanium, zinc, and others).

Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Corrosion Prevention of Corrosion Galvanization: The coating of iron with zinc. Even if the zinc coating is scratched, the iron is still protected since the zinc will be selectively oxidized. Copyright © 2011 Pearson Prentice Hall, Inc.

Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Corrosion Prevention of Corrosion Galvanization: The coating of iron with zinc. When some of the iron is oxidized (rust), the process is reversed since zinc will reduce Fe2+ to Fe: Fe(s) Fe2+(aq) + 2e E° = 0.45 V Fe2+ is higher up in the reduction table so it will be selectively reduced when in contact with zinc. Zn(s) Zn2+(aq) + 2e E° = 0.76 V Copyright © 2011 Pearson Prentice Hall, Inc.

Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Corrosion Prevention of Corrosion Cathodic Protection: Instead of coating the entire surface of the first metal with a second metal, the second metal is placed in electrical contact with the first metal: Anode: Mg(s) Mg2+(aq) + 2e E° = 2.37 V Cathode: O2(g) + 4H+(aq) + 4e 2H2O(l) Magnesium is farther down in the reduction table (on the right) so it is more likely to be oxidized than zinc. Zinc is used as a sacrificial anode for iron sea walls for hurricane protection (an iron wall placed into salt water for a long period of time is not a good idea from a corrosion perspective!). Blocks of solid zinc (sacrificial anodes) are attached to the iron and spaced at regular intervals. The sacrificial anodes are monitored over time and replaced when needed. E° = 1.23 V Attaching a magnesium stake to iron will corrode the magnesium instead of the iron. Magnesium acts as a sacrificial anode. Copyright © 2011 Pearson Prentice Hall, Inc.

Electrolysis and Electrolytic Cells Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Electrolysis and Electrolytic Cells Electrolysis: The process of using an electric current to bring about chemical change. Copyright © 2011 Pearson Prentice Hall, Inc.

Electrolysis and Electrolytic Cells Chapter 17: Electrochemistry 4/24/2017 Electrolysis and Electrolytic Cells Electrolysis: The process of using an electric current to bring about chemical change. Process occurring in galvanic cell and electrolytic cells are the reverse of each other In an electrolytic cell, two inert electrodes are dipped into an aqueous solution Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Electrolysis and Electrolytic Cells Anode: where oxidation takes place Anions are oxidized at this electrode labeled positive to reflect anions attraction to anode Cathode: where reduction takes places Cations are reduced at this electrode Labeled negative to reflect the cations attraction to cathode

Molten salt- mixture of cations and anions In general: The cation that is most easily reduced (the one with least negative, or most positive, reduction-half cell potential) is reduced first The anion is most easily oxidize ( the one has the least negative, or most positive, oxidation half-cell potential) is oxidized first

Electrolysis and Electrolytic Cells Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry Electrolysis and Electrolytic Cells 4/24/2017 2:41:42 PM Electrolysis of Molten Sodium Chloride Anode: 2Cl(l) Cl2(g) + 2e Cathode: 2Na+(l) + 2e 2Na(l) Overall: 2Na+(l) + 2Cl(l) 2Na(l) + Cl2(g) Copyright © 2011 Pearson Prentice Hall, Inc.

In aqueous solutions The cations of active metals-those that are not easily reduced, such as Li+, K+, Na+, Mg2+, Ca2+, and Al3+ - Cannot be reduced from aqueous solution by electrolysis because water is reduced at lower voltage.

Electrolysis and Electrolytic Cells Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Electrolysis and Electrolytic Cells Electrolysis of Aqueous Sodium Chloride Anode: 2Cl(aq) Cl2(g) + 2e Cathode: 2H2O(l) + 2e H2(g) + 2OH(aq) Overall: 2Cl(l) + 2H2O(l) Cl2(g) + H2(g) + 2OH(aq) Copyright © 2011 Pearson Prentice Hall, Inc.

Examples Predict the half-reaction occurring at the anode and the cathode for electrolysis of the following: a mixture of molten AlBr3 and MgBr2 An aqueous solution of LiI

Commercial Applications of Electrolysis Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Commercial Applications of Electrolysis Down’s Cell for the Production of Sodium Metal Copyright © 2011 Pearson Prentice Hall, Inc.

Commercial Applications of Electrolysis Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Commercial Applications of Electrolysis A Membrane Cell for Electrolytic Production of Cl2 and NaOH Copyright © 2011 Pearson Prentice Hall, Inc.

Commercial Applications of Electrolysis Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Commercial Applications of Electrolysis Hall-Heroult Process for the Production of Aluminum © 2012 Pearson Education, Inc. Copyright © 2011 Pearson Prentice Hall, Inc.

Commercial Applications of Electrolysis Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Commercial Applications of Electrolysis Electrorefining of Copper Metal © 2012 Pearson Education, Inc. Copyright © 2011 Pearson Prentice Hall, Inc.

Quantitative Aspects of Electrolysis Chemistry: McMurry and Fay, 6th Edition Chapter 17: Electrochemistry 4/24/2017 2:41:42 PM Quantitative Aspects of Electrolysis Charge(C) = Current(A) × Time(s) 1 mol e Moles of e = Charge(C) × 96,500 C Faraday constant Copyright © 2011 Pearson Prentice Hall, Inc.

Example God can be plated out of a solution containing Au3+ according to the following half-reaction: Au3+(aq) + 3e-  Au(s) What mass of gold (in grams) will be plated by the follow of 5.5A of current for 25 minutes?

Example Silver can be plated out of a solution containing Ag+ according to the following half-reaction: Ag+(aq) + e-  Ag(s) How much time (in minutes) would it takes to plate 12.0 g of silver using a current of 3.0A?