ELECTROCHEMISTRY Chap 20

Types of Electrochemical Cells Voltaic (galvanic) ⇒ spontaneously produces electricity Electrolytic ⇒ consumes electricity Reversible ⇒ reversing current reverses reaction e.g., auto lead storage battery, hydrogen fuel cell Irreversible ⇒ cannot be reversed due to different ½-rxn occurring at one or both electrodes e.g., dry cell, alkaline cell, Hg cell

Fig 20.3 A spontaneous oxidation-reduction reaction

Fig 20.4 a Voltaic Cell based on Equation 20.7 Zn (s) + Cu2+ (aq, 1 M) → Zn2+ (aq, 1 M) + Cu (s) 1.10

Components of a Cell All cells contain 2 electrodes and an electrolyte Anode – electrode where oxidation occurs e.g.: Zn (s) ⇌ Zn2+ (aq) + 2e− Cathode – electrode where reduction occurs e.g.: Cu2+ (aq) + 2e− ⇌ Cu (s) Liquid junction - maintains balance of ion charges e.g.: salt bridge or porous disk

Fig 20. 5 A voltaic cell that uses a salt bridge to complete the Fig 20.5 A voltaic cell that uses a salt bridge to complete the electrical circuit

Fig 20.6 A summary of the terminology used to describe voltaic cells

Fig 20.7 Atomic-level depiction of the Zn (s) / Cu2+ (aq) rxn

Fig 20.8 Atomic-level depiction of the voltaic cell in Fig 20.5 2+ 2+

Cell EMF under standard conditions Fig 20.9 Water analogy for electron flow Volt (V) - potential difference (E) between two points 1 V = 1 J/C

Also called the cell potential, Ecell Electromotive force (emf) ≡ potential difference between the anode and cathode in a cell Also called the cell potential, Ecell Fig 20.5

Standard Reduction (Half-Cell) Potentials, E° Only differences in potential measurable: ∴ measure E° in combination with a standard half-reaction in a cell Standard Hydrogen Electrode (SHE)

Standard Hydrogen Electrode (SHE) By convention, reaction written as a reduction: 2H+(aq; 1.00 M) + 2e− ⇌ H2 (g, 1.00 atm) By definition: E°SHE ≡ 0.00V as a cathode When used as an anode: H2 (g, 1.00 atm) ⇌ 2H+(aq; 1.00 M) + 2e− By definition: E°SHE ≡ 0.00V as a anode

Standard Hydrogen Electrode (SHE)

Standard Electrode Potential E° ≡ potential of a cell with electrode of interest acting as a cathode and SHE as the anode e.g., Cu/Cu2+ couple E° = 0.34 V e.g., Zn/Zn2+ couple E° = − 0.76 V E° describes ½-rxn written as a reduction with respect to SHE E° > 0: spontaneous w.r.t. SHE E° < 0: nonspontaneous w.r.t. SHE

Fig 20.11 A voltaic cell using a SHE to measure the Eo of a Zn/Zn2+ electrode

Standard Cell Potentials Ecell  = Ered (cathode) − Ered (anode) In the previous example: Ecell  = Ered (SHE) − Ered (Zn2+/Zn) Ecell  = 0.00 V − 0.76 V = − 0.76 V Ered (Zn2+/Zn) = − 0.76 V 

Table 20.1 Standard Reduction Potentials in Water at 25 oC Very good oxidizing agent (poor reducing agent) Very poor oxidizing agent (very good reducing agent)

Calculating Standard Cell Potentials Ecell  = Ered (cathode) − Ered (anode) Fig 20.5 For the oxidation: For the reduction: Ered = −0.76 V  Ered = +0.34 V  Ecell  = Ered (cathode) − (anode) = +0.34 V − (−0.76 V) = +1.10 V © 2009, Prentice-Hall, Inc.

  Ecell = Ered (cathode) − Ered (anode) Sample Exercise 20.6 Calculating E°cell from E°red Using the standard reduction potentials listed in Table 20.1 (p 857), calculate the standard emf for the voltaic cell, which is based on the reaction: Ecell  = Ered (cathode) − Ered (anode) Cathode: Cr2O72−(aq) + 14 H+(aq) + 6 e− → 2 Cr3+(aq) + 7 H2O(l) Eo = +1.33 V Anode: 3 I2(s) + 6 e− → 6 I− (aq) Eo = +0.54 V Ecell  = Ered (cathode) − Ered (anode) = 1.33 V − 0.54 V = 0.79 V