Refer Back to Chapter 3.7 for Oxidation State/Number;

Ch 12 Electrochemistry

Refer back to chapter 3.7 for oxidation state/number;
Recall: oxidation as loss of electrons; reduction as gain of electrons; oxidation - reduction (or, redox) reactions; balancing redox equations; redox titrations (introduced in chapter 6.4, pp. 163-170)

12.1 Electrochemical Cells

energy of a spontaneous redox reaction can be used to do electrical work in a voltaic (= galvanic) cell by forcing electron transfer through an external path

eg., first with direct transfer, then indirect/external (alternative to Zn/Cu example in Introduction):

direct reaction: Cu strip in a beaker of Ag+ (aq.)
blue colour of Cu2+gradually builds up
Cu begins to dissolve
Ag begins to deposit (black) on Cu strip

reaction via external path: Voltaic Cell
2 metal strips (= electrodes) in separate compartments or beakers
Cu2+ and Ag+ solutions as nitrates, one in each
2 means of connection of solution
salt bridge, Fig. 12.1 and 12.2
porous glass disk

two electrodes:

anode (oxidation): Cu(s) to Cu2+(aq)

cathode (reduction): Ag+(aq) to Ag(s)

electron flow is from anode (removed from Cu) to cathode (supplied to Ag+)
this is general
need for solution connection: maintain charge neutrality (eg. Na+ and NO3- move through bridge)
special aspects of this example, not necessarily in all voltaic cells:
electrodes can be simple conductors, not chemical participants (eg. Pt foil)
need not have deposition of solid at either electrode

Galvanic & Electrolytic Cells

external path as above, Example: Cd(s) + Ni2+(aq)  Ni(s) + Cd2+(aq)
at anode Cd is oxidized, Cd2+ goes into solution, balanced by NO3- coming from salt bridge
at cathode, Ni2+ is reduced, comes out of solution as Ni, balanced by NO3- going into bridge
negative charge as NO3- flows through bridge from cathode (Ni) to anode (Cd)
negative charge as e- flows through wire from anode (Cd) to cathode (Ni)
conversely, Na+ moves through bridge to Ni side (Ni2+ consumed) from Cd side (Cd2+ produced)

Galvanic Cell: spontaneous reaction; current flows between electrodes; potential difference (see below) measured as a positive voltage; electrical energy produced; examples above

Electrolytic Cell: non-spontaneous reaction is forced by applying a potential between electrodes; electrical energy consumed; example, reverse of Cu/Ag reaction above:

12.2 Free Energy & Cell Voltage

in Galvanic cells above, spontaneous direction of e- flow from Zn to Cu, from Cu to Ag and from Cd to Ni, ie.- to lower (potential) energy, from anode to cathode
potential difference measured in volts, V:

in Zn/Cu eg., , potential is 1.10 V

= driving force or electron pressure

= electromotive force, emf

= cell potential, Ecell

Standard Cell Voltage

define: Eocell (or,Eo), standard emf or standard cell potential with 1 M in all solution components (1 bar for gases) at 25oC; cell potentials positive for the spontaneous (product-favoured) direction
note: reverse the reaction, change the sign of Eo

Eo and Go (Emf & Free-Energy Change)

both are measures of a reaction’s tendency to proceed to products

n = number of moles of electrons transferred

F = the Faraday, molar equivalent of e- charge

= 96,500 C/(mol e-)

= 96,500 J/(V.mol e-)

Example 12.4: for Zn/Cu2+ reaction, Eo = +1.10 V, calculate Go = - 212 kJ

Calculating the Cell Potential, Eo

Half-Cell Potentials (Voltages)

cell potential comprised of two half-cell potentials, standard oxidation potential (for one couple) and standard reduction potential (for other couple)



scaled according to a reference half-cell, standard hydrogen electrode, at 0 V (Fig. 12.3):

eg. voltaic cell with Zn anode and standard hydrogen electrode as cathode:

this determines the standard oxidation half-cell potential for the Zn/Zn2+ couple

opposite reaction, reduction of Zn2+ to Zn has a standard reduction potential of same magnitude, opposite sign
similarly for Cu/Cu2+, but here Cu is the cathode and the hydrogen electrode is the anode
cell potential = + 0.34V = standard oxidation potential for Cu/Cu2+
return to Zn + Cu2+ reaction:

Eocell = Eoox + Eored

= 0.76 + 0.34 = 1.10 V (Fig. 12.4)

can also determine half-reaction (half-cell) potentials by measurement against a known couple other than the standard hydrogen electrode
Example, Ni/Ni2+ is done vs. Zn/Zn2+
do Example 12.5

for simplicity of tabulation, all standard half-cell potentials are listed as reduction potentials, eg. Appendix E

note: half-cell potential an intensive property, hence not multiplied by stoichiometry numbers

Using Standard Electrode Potentials

Oxidizing and Reducing Agents

trends in Appendix E
the more positive an E, the greater the tendency for the half-reaction to occur as written
F2 most easily reduced, strongest oxidizing agent
Li+poorest oxidizing agent
Li strongest reducing agent
reaction between any substance on the left column and any one lower on the right column is spontaneous (see below)

Predicting Eo & Spontaneity of Redox Reactions

positive emf indicates spontaneous process
predict from half-cell potentials:

half-cells:

overall: E = 1.24 V  spontaneous
note: complication if adding or subtracting half-cell reactions to get a new half-cell reaction; then have to account for stoichiometry

(aside: Electrical Work)

voltaic cell: wmax = -nFE (spontaneous)

electrolytic cell: wmin = -nFE(non-spontaneous)

12.3 Concentration Effects & the Nernst Equation

from: (NB: implicit in ch. 9.7)

used to calculate emf under non-standard conditions, Example 12.7
or, to calculate a concentration, if cell emf measured. Example 12.9
or, to calculate an equilibrium constant, Example 12.8
application: the pH meter (read for interest)
read 12.4: Batteries & Fuel Cells for interest

12.5 Corrosion & Its Prevention

nearly all metals undergo thermodynamically favoured oxidation in air
"skin" of oxide frequently protects against further oxidation
cathodic and anodic areas on the metal, Fig. 12.14
for iron, if limited O2 (slow) :

for iron, if O2 and water available (fast) :

and Fe2+ oxidized further, near cathode:

can protect iron from corrosion:
anodic inhibition: coating with tin; good until surface broken; more recently, promote Fe2O3 formation as “skin”
cathodic inhibition: coating with Zn; Zn is sacrificial anode in preference to Fe; "cathodic protection", i.e.- force the metal to become a cathode

12.7 Electrolysis

reverse of foregoing: non-spontaneous redox reactions "driven" by electrical energy
eg. molten NaCl
note sign convention opposite to voltaic cell
used in some metals production

electrolysis of aqueous solutions
eg. brine:
at cathode, H2 production preferred:

while at anode, little thermodynamic preference:

(if use NaI, where I- I2, Eo = -0.535 V; more clearcut)

and, due to "overvoltage", Cl- oxidation preferred
net:

all products are industrial commodities
minimum emf required: 0.83 + 1.36 = 2.19 V

Electrolysis with Active Electrodes

electroplating: make anode of metal to be deposited onto an object; make object the cathode
eg. nickel

(in preference to water electrolysis)