Cell spontaneous redox reactions take place in a

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Unformatted text preview: ous redox reactions take place in a voltaic cell. aka galvanic cell. • Nonspontaneous redox reactions can be forced to occur in an electrolytic cell by the addition of electrical energy. 11 Voltaic Cell Oxidation and reduction reactions kept separate: half-cells. Electric circuit: e- flow via wire, ion flow via solution & salt bridge Electrodes (anode & cathode): conductive solid allowing transfer of e-. Resulting voltage is called electromotive force, emf, in units of volts oxidation occurs, cations released Anode Zn(s) Cathode cations attracted, reduction occurs Cu(s) Cell Potential & Notation • The difference in potential energy between anode and cathode in a voltaic cell is called the cell potential or emf (in Vs). • The cell potential under standard conditions is called the standard emf, E°. 25°C, 1 bar for gases, 1 M concentration for solutions sum of the cell potentials for the half-reactions: E° = E°red2 + E°ox1 • Electrode(anode) | electrolyte || electrolyte | Electrode(cathode) oxidation half-cell on left, reduction half-cell on the right single “|” denote the phase barrier − if multiple electrolytes in same phase, a comma is used rather than | − often use an inert electrode double line “||” denote salt bridge Some times simply write as Red1,Ox1|| Red2,Ox2 Redox couple: Ox1/Red1 and Ox2/Red2 Write the cell notation & sketch the voltaic cell for Fe(s) + MnO4−(aq) → Mn2+(aq) + Fe2+(aq) Fe(s) | Fe2+(aq) || MnO4−(aq), Mn2+(aq), H+(aq) | Pt(s) Redox couples & Half-reactions Standard Reduction Potential • A half-reaction with a strong tendency to be reduced has a large positive half-cell potential when two half-cells are connected, the electrons will flow toward that half-reaction. • We cannot measure the absolute tendency of a halfreaction, we can only measure it relative to another half-reaction, i.e. relative potential/emf. • We select as a standard half-reaction the reduction of H+ to H2 under standard conditions, which we assign an absolute potential = 0 V. standard hydrogen electrode, SHE 2H+(aq,1M)+2e- → H2(g,1bar) Eo=Eored=Eoox=0V H2(g,1bar) → 2H+(aq,1M)+2e- Eo re d= 0 Zn(s) → Zn2+(aq,1M) +2e- Eoox = 0.76V Zn2+(aq,1M) +2e- → Zn(s) Eored = –0.76V Standard Reduction Potential Half-Cell Potentials • • The SHE reduction potential is defined to be exactly 0 V. half-reactions with a stronger tendency toward reduction than the SHE have a positive value for E°red half-reactions with a stronger tendency toward oxidation than the SHE have a negative value for E°red It is convenient to always write the half-cell potential based on its reduction process. Red1 + Ox2 Half reactions: Red1 Ox2 + ν e- Ox1 + Red2 Cell potentials refer to reduction potentials. Ox1 + ν eRed2 Red1,Ox1|| Red2,Ox2 Spontaneous only if Eo>0 EoL = Eored1 = -Eoox1 EoR = Eored2 Eo = Eored2 + Eoox1 = Eored,R – Eored,L = EoR – EoL Predicting whether a metal dissolves in acid Exception: HNO3...
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This note was uploaded on 09/25/2013 for the course CHE 301 taught by Professor Raineri,f during the Fall '08 term at SUNY Stony Brook.

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