ElectrochemistryChapter 20

Brown, LeMay, and Bursten

Definition The study of the relationships between

electricity and chemistry Review redox reactions Review balancing redox reactions in acid and

base

Voltaic Cell (also called Galvanic Cell) Device in

which the transfer of electrons takes place through an external pathway.

Electrons used to do work

Summary of Cell Each side is a half-cell Electrons flow from oxidation side to reduction side

– determine which is which Salt bridge allows ions to move to each terminal so

that a charge build up does not occur. Assignment of sign is this:

Negative terminal = oxidation (anode) Positive terminal = reduction (cathode)

Salt bridge allows ions to move to each terminal so that a charge build up does not occur. This completes the circuit.

Cell EMF Flow is spontaneous Caused by potential difference of two half

cells. (Higher PE in anode.) Measured in volts (V) 1 volt = 1 Joule/coulomb This is the electromotive force EMF (force

causing motion of electrons through the circuit.

Ecell

Also called the cell potential, or Ecell

Determined by reactant types, concentrations, temperature

Under standard conditions, this is E°cell

25° C, 1 M or 1 atm pressure This is 1.10 V for Zn-Cu Shorthand: Zn/Zn2+//Cu2+/Cu

Reduction Potentials Compare all half cells to a standard (like sea

level) 2H+ + 2e- → H2(g) = 0 volts (SHE) The greater the E°red, the greater the driving

force for reduction (better the oxidizing agent) In a sense, this causes the reaction at the anode

to run in reverse, as an oxidation. Use this equation:

E°cell = E°red (cathode) - E°red (anode)

Trends

Spontaneity Positive E value indicates that the process is

spontaneous as written. Activity series of Metals – listed as oxidation

reactions Reduction potentials in reverse Example, Ag is below Ni because solid Ni can

replace Ag in a compound. Actually, Ni is losing electrons and thus being oxidized by Ag+. Ag is listed very high as a reduction potential.

Relationship to ΔG ΔG = -nFE

n = number of electrons transferred F = Faraday constant = 96,500 C/mol or 96,500

J/V-mol Why negative? Spontaneous reactions have

+E and – ΔG. Volts cancel, units for ΔG are J/mol Standard conditions: ΔG° = -nFE°

Nernst Equation Nonstandard conditions – during the life of

the cell this is most common Derivation E = E ° - (RT/nF)lnQ Consider Zn(s) + Cu2+ → Zn2+ + Cu(s) What is Q? What is E when the ions are both 1M? What happens as Cu2+ decreases?

Concentration Cells Same electrodes and solutions, different

molarities. How will this generate a voltage? Look at

Nernst Equation. E = E ° - (RT/nF)lnQ When will it stop? Basis for a pH meter and regulation of

heartbeat in mammals

EMF and equilibrium When cell continues to discharge, E

eventually reaches 0. At this point, because ΔG = -nFE, it follows that ΔG = 0.

Equilibrium! Therefore, Q = Keq Derivation logKeq = nE°/0.0592

Batteries Portable, self-contained

electrochemical power source

Batteries in series, voltage is added.

Things to consider Size (car vs. heart) Amount of substances

before it reaches equilibrium

Toxicity (car vs. heart) A lot a voltage or a little

(car vs. heart) Example – alkaline camera

battery Dry – no water

Fuel Cells Not exactly a battery, because it is open to

the atmosphere How does the combustion of fuel generate

electricity? – heats water to steam which mechanically powers a turbine that drives a generator – 40% efficient

Voltaic cells are much more efficient http://www.fueleconomy.gov/feg/fuelcell8.s

wf

Corrosion Undesirable

spontaneous redox reactions

Thin coating can protect some metals (like aluminum) – forms a hydrated oxide)

Iron - $$$$$

Protection Higher pH Paint surface Galvanize (zinc

coating) – why? Zinc is a better anode Called cathodic

protection – sacrificial metal

More dramatic

Electrolysis Cells that use a battery or outside power

source to drive an electrochemical reaction in reverse

Example NaCl → Na+ + Cl-

Reduction at the cathode, oxidation at the anode

Voltage source pumps electrons to cathode.

Diagram

Solutions High temperatures necessary for previous

electrolysis (ionic solids have high MP) Easier for solutions, but water must be considered Example: NaF Possible reductions are:

Na+ + e- → Na(s) (Ered = -2.71 V) 2H2O + 2 e- → H2(g) + 2 OH- (Ered = -.83 V)

Far easier to reduce water! continue

Continued Look at possible oxidations:

2F- → F2(g) (Ered = 2.87 volts) 2H2O → O2(g) + 4H+ + 4e- (Ered = 1.23 volts) Far easier to oxidize water, or even OH-!

So for NaF, neither electrode would produce anything useful, and doesn’t by experiment

With NaCL, neither electrode is favored over water. However, the oxidation of Cl- is kinetically favored, and thus occurs upon experimentation!

Use Ered values of two products to find Ecell (minimum amount of energy that must be provided to force cell to work)

Active electrodes If electrode is not inert, it

can be coated with a thin layer of the metal being reduced, if its reduction potential is greater than that of water.

This is called electroplating Ecell = 0, so a small

voltage is needed to push the reaction.

Quantitative relationship