Gang Wang Instrumental_Week12&15_CH22_25-2010

Gang Wang Instrumental_Week12&15_CH22_25-2010 -...

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Chapter 22: Introduction to Electroanalytical Chemistry Two general categories: 1) Potentiometric Systems – measure voltage (i.e., potential) of a galvanic cell (produces electricity spontaneously) 2) Voltammetric Systems – control potential & usually measure current in an electrolytic cell (consumes power to cause an electrochemical reaction to occur) Many different electroanalytical methods: •fast • inexpensive • in situ • information about oxidation states stoichiometry rates charge transfer equilibrium constants
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Potentiometry Determine concentrations by measuring the potential (i.e., voltage) of an electrochemical cell (galvanic cell) Two electrodes are required 1) Indicator Electrode–potential responds to activity of species of interest 2) Reference Electrode–chosen so that its potential is independent of solution composition. Representation of Electrochemical Cell Most contain • external wires (electrons carry current) ion solutions (ions carry current) interfaces or junctions All contain • complete electrical circuit • conducting electrodes (metal, carbon)
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- + ionic conductor electronic conductor electronic conductor electrode electrode V Δ E Representation of Electrochemical Cell c abd Cathode or Indicator Anode or Reference V Represents device to measure potential (voltage) without drawing significant current i.e potentiometer or electrometer (high input impedance > 100 M Ω (mega ohms) Electrons transferred at electrode surface at liquid/solid interface Potential difference (voltage) is measure of tendency to move to equilibrium Galvanic cell - cell develops spontaneous potential difference Overall: Zn(s) + Cu2 + (aq) Zn2 + (aq) + Cu(s) Half reactions: Zn(s) Zn2 + + 2e Oxidation Cu2 + + 2e −→ Cu(s) Reduction Convention: Reduction at Cathode Oxidation at Anode Galvanic cell - Zn anode ( negative ), Cu cathode ( positive )
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Electrolytic cells - require potential difference greater than galvanic potential difference (to drive away from equilibrium) Galvanic cell Zn(s) Zn2 + + 2e Oxidation Cu 2 + + 2e Cu(s) Reduction Electrolytic cell Zn 2 + + 2e Zn(s) Reduction Cu(s) Cu 2 + + 2e Oxidation Electrolytic cell - Zn cathode ( positive ), Cu anode ( negative ) Short-Hand Cell notation: Convention: Anode on Left Zn|ZnSO4 (0.01 M)||CuSO4 (0.01 M)|Cu liquid-liquid interface Not all cells have liquid-liquid junctions (Fig 22-3) AgCl (s) Ag + (aq) + Cl (aq) H 2 (g) H 2 (aq) Cathode: Ag + (aq) + e Ag(s) Anode: H 2 (aq) 2H + (aq) + 2e Overall: 2AgCl(s) + H 2 (g) 2Ag(s) + 2H + + 2Cl Pt,H 2(p = 1atm)|H + (0.01 M),Cl (0.01 M),AgCl (sat'd)|Ag
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Electrode Potentials: • Cell potential is difference between anode and cathode potential Ecell = Ecathode Eanode when half-reactions written as reductions Example: 2AgCl(s) + H 2 (g) 2Ag(s) + 2H + + 2Cl 2AgCl(s) + 2e 2Ag(s) + 2Cl 2H + + 2e H2 (g) electrons on left Galvanic cell E cell =E cathode -E anode =+0.46 V Can't measure potential on each electrode independently – only differences E cell = E ind –E ref (+ E J ) E J = junction potential, a non ideal potential which develops across the interface between two dissimilar solutions
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This note was uploaded on 02/13/2012 for the course UML 84.314 taught by Professor Ryan during the Fall '11 term at UMass Lowell.

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Gang Wang Instrumental_Week12&15_CH22_25-2010 -...

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