CHM 320 - Lecture 23 Chapt 14Chapter 14 – Fundamentals of ElectrochemistrycontinuedVocabulary:Oxidation - Loss of electrons (increase oxidation state)Reduction – Gain of electrons (decrease oxidation state)Oxidizing agent – Substance that takes electrons (Standard reduction potential is more positive.)Reducing agent – Substance that gives up electrons (Standard reduction potential is more negative.)Anode – Electrode that oxidation takes place (positive polarity)Cathode – Electrode that reduction takes place (negative polarity)Coulombs – unit of chargeVolt – unit of potentialAmpere – unit of current (coulomb/sec)Joule – unit of workWatt – unit of power (work/sec)CHM 320 - Lecture 23 Chapt 14Review of Important Electrochem Equationsq (coulombs) = n (moles) x F (Faraday constant, coulombs/mole)Work (joules) = E (volts) x q (coulombs)E (volts) = I (Amp) x R (ohms, Ω)Power (joules/sec) = E (volts) x I (Amp)ΔG (Gibb’s free energy) = - n (moles) x F (Faraday constant) x E (volts)(This equation shows the thermodynamic importance of the cell potential.)CHM 320 - Lecture 23 Chapt 14Cell PotentialFor any cell, the measured potential between the anode and the cathode can be calculated:cell)left theandter potentiome theof sideleft he(usually treaction) (oxidationter potentiome theof terminal(-) the toconnected)(electrode cell-half for the potential theis E___________________________________________________usually) ter,potentiome theof sideright theand cellright (the reaction) (reductionter potentiome theof terminal )( the toconnected )(electrode cell-half for thepotential calculated theis E whereE E E--cell+−=++NOTE: E+and E-are both reduction potentialsCHM 320 - Lecture 23 Chapt 14What is happening at the electrode(s) and how do we describe the cell?CHM 320 - Lecture 23 Chapt 14What is happening at the electrode(s) and how do we describe the cell?(s)4(aq)4(aq)(s)(s)0(aq)2(aq)2(s)0(s)-2(aq)-(aq)20(s)Cu |CuSO || ZnSO |Znsolutions theabout something us tells ions-counter the Includingright the on cathode left, the on Anodebridgesalt the marks line vertical double adifference phase the marks line vertical single asymbols! use weshorthand, inCu Zn Cu Zn :reaction Cell CompleteCu e 2 Cu :rxn-half Cathodee 2 Zn Zn :rxn-half Anode+⇔+⇔++⇔++++CHM 320 - Lecture 23 Chapt 14The Standard Hydrogen Electrode (SHE)• The basis by which all other measurements are made.• Assigned a potential of zero by definition!• Not practical for regular useHydrogen Half-CellH2(g)= 2 H+(aq)+ 2 e-reversible reactionSHE consists of a platinum electrode covered with a fine powder of platinum around which H2(g)is bubbled. Its potential is defined as zero volts.CHM 320 - Lecture 23 Chapt 14CHM 320 - Lecture 23 Chapt 14CHM 320 - Lecture 23 Chapt 14Standard Potentials• Standardized potentials (Eo), listed as reductions, for all half-reactions• Measured versus the S.H.E (0)• Used in predicting the action in either a galvanic cell or how much energy would be needed to force a specific reaction in a non-spontaneous cell• Assumes an activity of one for the species of interest (usually a fair approximation) at a known temperature in a cell with the S.H.E.• Assumes that the cell of interest is connected to the (+) terminal of the potentiometer (voltmeter) and the S.H.E. is connected to the (-) terminalCHM 320 - Lecture 23 Chapt 14Better Oxidizing Agents in upper left hand corner.Better ReducingAgents in lowerRight hand cornerCHM 320 - Lecture 23 Chapt 14Calculating Ecell• Determine E+and E-(Even though E-is an oxidation reaction, the E-is determined for the reduction reaction then subtracted.)• Calculate Ecell• Write a balanced cell reaction, by adding the two half-reactions– Write out the right cell half reaction– Write out the left cell half reaction and reverse it– Add the two reactions together to get a net, balanced cell reaction.• If, you use the conventions described here,then:– If Ecell>0, the reaction is spontaneous to the right– If Ecell<0, the reaction is spontaneous to the leftCHM 320 - Lecture 23 Chapt 14Relationship between E° and the Equilibrium ConstantRecall:ΔG (Gibb’s free energy) = - n (moles) x F (Faraday constant) x E (volts)K (equilbrium constant) = e- ΔG/RTK = e- (-nFE°/RT)At 25° C K = 10nE°/0.05916 or E° = (0.05916/n) log KCHM 320 - Lecture 23 Chapt 14Problem - Calculate the E° and K for the following reaction:Cr2++ Fe(s) ⇔ Fe2++ Cr(s)CHM 320 - Lecture 23 Chapt 14Nernst EquationThe Nernst equation allows you to determine the cell potential when the activities of the species involved ≠ 1(i.e. non-standard conditions, more typical to real-life)For aA + ne-⇔ bBE = E° - (RT/nF) ln (Abb/ Aaa)At 25° C:E = E° - (0.05916/n) log (Abb/ Aaa)CHM 320 - Lecture 23 Chapt 14Nernst Equation• Accounts for potentials of cells where the reagents are not at an activity of 1– Remember that standard potentials are at A=1• Accounts for the number of electrons transferred in a reaction, the temperature of the reaction, LeChatelier’s Principle and a variety of other factors• Used to calculate E+and E-under non-standard conditions – Most real cases!CHM 320 - Lecture 23 Chapt 14The Nernst Equation for Complete Reactions…..• Setup two Nernst equations– One for E+– One for E-• Solve each Nernst equation to get E+and E-• Solve for Ecell(Ecell= E+-E-)CHM 320 - Lecture 23 Chapt 14A galvanic cell is assembled in which the left cell is the anodewhere cadmium metal electrode is oxidized to cadmium ion in 0.010 M cadmium nitrate. In the right cell, the cathode, silver ion is reduced to silver metal on a silver metal electrode in 0.50 M silver nitrate.1. Draw the cell (both a picture and a schematic diagram)2. Write the half and net cell reactions3. Calculate the net cell voltage4. Indicate in which direction the cell is
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