Lab 9: Electrochemistry Spring 2018 Lab 9: Electrochemistry

Goals Measure voltage of a galvanic cell Determine how concentration affects voltage Graphically determine value of Faraday constant

Oxidation-Reduction Reactions A.k.a. Redox reactions Involve electron transfer Oxidation = Loss of electrons Reduction = Gain of electrons Balanced using half-reactions Representations of oxidation and reduction

Half-Reactions Show how substances lose or gain electrons Electrons written into reactions Balance charge between reactants/products Example: Fe(s) + Cd2+(aq) → Fe2+(aq) + Cd(s) Fe(s) → Fe+2(aq) + 2e- (oxidation) Cd2+(aq) + 2e- → Cd(s) (reduction) 0 charge = 0 charge

Balancing Redox Reactions Balanced with half-reactions Need same # of electrons in both so they cancel Example: Unbalanced Reaction: Fe(s) + Cl2(aq) → Fe3+(aq) + Cl-(aq) Half-Reactions: Fe(s) → Fe3+(aq) + 3e- Cl2(aq) + 2e- → 2Cl-(aq) Balanced Half-Reactions: 2Fe(s) → 2Fe3+(aq) + 6e- 3Cl2(aq) + 6e- → 6Cl-(aq) Balanced Reaction: 2Fe(s) + 3Cl2(aq) → 2Fe3+(aq) +6Cl-(aq)

Galvanic Cells Produce voltage by separating half-reactions into half-cells Necessary components: Electrodes in matching solutions Salt bridge External circuit between electrodes to allow electron flow Electron flow Salt Bridge Cu Zn Cu2+ Zn2+

Electrodes Solid surface for redox reaction Two types: Anode – site of Oxidation (-) Cathode – site of Reduction (+) Current carried from anode to cathode “The RED CAT gets FAT”

The Experiment The reaction: Cu(s) + 2Ag+(aq) → Cu2+(aq) + 2Ag(s) The set-up: Half-cells in beral pipets Agar plug salt bridge Voltmeter attached to electrodes measures cell potential (Ecell)

Cell Potential Symbol: Ecell Units: Volts (V) Related directly to Gibbs Free Energy: n = # moles of transferred electrons (see half reactions) F = Faraday’s constant (96,485 C/mol) Potential in a galvanic cell is positive (spontaneous)

Standard Cell Potential Symbol: E°cell Units: Volts (V) Voltage under standard conditions 1 M and 1 bar Can be calculated for reaction using standard reduction potentials

Standard Reduction Potential (E°red) Assigned to reduction half reactions (e.g.) Cu2+(aq) + 2e- → Cu(s) E°red = 0.342 V See next slide for other values Magnitude unaffected by balancing coefficients I.e. Don’t multiply value! Same number for reverse direction (oxidation) Reversing direction leads to sign change.

More (+), more likely to be reduced More (-), more likely to be oxidized

Calculating E°cell Combination of potentials for both half-cells: E cell ° = E red ° − E ox ° Example: Fe(s) + Cl2(aq) → Fe3+(aq) + Cl-(aq) E red ° Oxidation: 2Fe(s) → 2Fe3+(aq) + 6e- -0.037 V Reduction: 3Cl2(aq) + 6e- → 6Cl-(aq) +1.358 V E cell ° = +1.358 V – (-0.037 V) = +1.395 V Note: Eox is the same as Ered for a particular half-reaction (subtraction accounts for sign change associated with opposite direction).

Non-Standard Conditions Potential (Ecell) affected by concentration Concentrations used to calculate Ecell using the Nernst Equation: R = gas constant (8.314 J/(K*mol) T = temperature (K) n = # moles of transferred electrons (see half reactions) F = Faraday’s constant (96,485 C/mol) Q = Reaction quotient that incorporates the concentrations

Reaction Quotient (Q) Same set-up as equilibrium constant expression: Products as numerator raised to coefficient Reactants as denominator raised to coefficient Concentrations NOT at equilibrium Can be initial concentrations Example: N2O4(g) ⇌ 2NO2(g)

Determining F Graphically Nernst equation can be graphed as a straight-line equation: Slope used to calculate Faraday’s constant y = b + m * x Experimentally—measured potential for the five trials Natural log of ratio of solution concentrations for the five trials