Chapter 19 Electrochemistry Insert picture from First page of chapter Copyright McGraw-Hill 2009

19.1 Balancing Redox Reactions Half-reaction method – acid solution Can add H2O Can add H+ Steps Separate the unbalanced reaction into half-reactions. Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Balance each of the half-reactions with regard to atoms other than O and H. Balance both half-reactions for O by adding H2O. Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Balance both half-reactions for H by adding H+. Balance both half-reactions for charge by adding electrons. Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 multiply one or both of the half-reactions by the number(s) required to make the number of electrons the same in both. add the balanced half-reactions together and cancel any identical terms that appear on both sides. Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Additional steps needed in basic solution For each H+ ion in the final equation, add one OH ion to each side of the equation, combining the H+ and OH ions to produce H2O. Make any additional cancellations made necessary by the new H2O molecules. Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 19.2 Galvanic Cells Galvanic cell - the experimental apparatus for generating electricity through the use of a spontaneous reaction Electrodes Anode (oxidation) Cathode (reduction) Half-cell - combination of container, electrode and solution Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Salt bridge - conducting medium through which the cations and anions can move from one half-cell to the other. Ion migration Cations – migrate toward the cathode Anions – migrate toward the anode Cell potential (Ecell) – difference in electrical potential between the anode and cathode Concentration dependent Temperature dependent Determined by nature of reactants Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 cell diagram convention anode cathode salt bridge phase boundary Copyright McGraw-Hill 2009

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19.3 Standard Reduction Potentials Designated Eo Measured relative to the standard hydrogen electrode (SHE) Standard conditions Assigned a value of 0 V Reaction Copyright McGraw-Hill 2009

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Copyright McGraw-Hill 2009 Determine the overall cell reaction and E°cell (at 25°C) of a galvanic cell made of an Al electrode in a 1.0 M Al(NO3)3 solution and a Cu electrode in a 1.0 M Cu(NO3)2 solution. Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Cell Potential Anode Cathode Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Cell Reaction Copyright McGraw-Hill 2009

19.4 Spontaneity of Redox Reactions Under Standard-State Conditions Based on electric charge and work n is the number of moles of electrons F is the Faraday constant 96,500 J/V . mol e Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 At standard conditions Relation to the equilibrium constant Using the values of R, T (298 K) and F Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Determine the value of a) DGo and b) K for the following reaction. Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 a) Standard free energy change, DGo Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 b) Equilibrium constant, K Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 19.5 Spontaneity of Redox Reactions Under Conditions Other than Standard-State Conditions The Nernst Equation –derived from thermodynamics or Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Will the following reaction occur spontaneously at 298 K if [Fe2+] = 0.60 M and [Cd2+] = 0.010 M? Copyright McGraw-Hill 2009

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Copyright McGraw-Hill 2009 Spontaneous since E is a positive even though very small. Copyright McGraw-Hill 2009

Concentration cells – galvanic cells composed of the same material but differing in ion concentrations 30

Copyright McGraw-Hill 2009 19.6 Batteries A battery is a galvanic cell, or a series of cells connected that can be used to deliver a self-contained source of direct electric current. Dry Cells and Alkaline Batteries no fluid components Zn container in contact with MnO2 and an electrotyte Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Dry Cell Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Alkaline Cell Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Lead Storage Batteries Six identical cells in series Lead anode and PbO2 cathode Immersed in H2SO4 Each cell delivers ~ 2 V Rechargeable Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Lead Storage Battery Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Lithium-Ion Batteries The overall cell potential is 3.4 V, which is a relatively large potential. Lithium is also the lightest metal—only 6.941 g of Li (its molar mass) are needed to produce 1 mole of electrons. recharged hundreds of times. Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Fuel Cells Direct production of electricity by electrochemical means Increased efficiency of power production In its simplest form Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Simple Form of Fuel Cell Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 19.7 Electrolysis electrolysis - the use of electric energy to drive a nonspontaneous chemical reaction electrolytic cell – the cell used to carry out electrolysis same principles apply to both galvanic and electrolytic cells in aqueous solutions you must also consider the oxidation or reduction of water Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Examples Molten sodium chloride Carried out in a Downs cell Laboratory Commercial Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Electrolysis of water Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Electrolysis of aqueous sodium chloride solution Possible anode reactions Possible cathode reactions Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Consider overvoltage – difference between estimated and actual voltage Overvoltage for O2 is higher than Cl2 Cell reaction Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Quantitative Applications- measure the current (in amperes) that passes through an electrolytic cell in a given period of time. Mass of product formed or reactant consumed Proportional to the amount of electricity transferred Molar mass Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Defining equation Coulomb (C) Second (s) Ampere (A) (coulomb/second) Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 A constant current of 0.912 A is passed through an electrolytic cell containing molten MgCl2 for 18 h. What mass of Mg is produced? Copyright McGraw-Hill 2009

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Copyright McGraw-Hill 2009 19.8 Corrosion Corrosion - generally refers to the deterioration of a metal by an electrochemical process. Many metals undergo corrosion For example, corrosion of Fe, oxidation of Al Can be enhanced by atmospheric conditions ( e.g. acidic medium) Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Electrochemical Process for the Formation of Rust Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Electochemical processes can be used to prevent corrosion Passivation – formation of a thin oxide layer by treating with an oxidizing agent Formation of an alloy Coating with a layer of a less active metal Tin cans Galvanization (zinc-plating) Zinc oxide coating constitutes the protective coating Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Key Points Balancing redox reactions Half-reaction method Galvanic cells Electrodes (anode and cathode) Cell potential (Ecell) Cell diagrams Standard Reduction potentials Relative to the standard hydrogen electrode Used to determine cell potential Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Spontaneity – standard conditions Free energy Equilibrium Constant Spontaneity – nonstandard conditions Nernst Equation Copyright McGraw-Hill 2009

Copyright McGraw-Hill 2009 Batteries Dry cell and alkaline batteries Lead storage battery Lithium-ion batteries Fuel cells Electrolysis Molten salts Aqueous solutions Quantitative applications Corrosion Metal deterioration Prevention Copyright McGraw-Hill 2009