Chapter 17Electrochemistry

Redox review (4.9) 17.1-17.2 17.4-17.5 17.6-17.7

Review Oxidation-Reduction

Involves transfer of electrons from reducing agent to oxidizing agent

Oxidation= loss of e- (increase in oxid #)

Reduction= gain of e- (decrease in oxid#)

GER and LEO

REVIEW

1. atom in element = 02. monatomic ion = charge3. fluorine = -14. oxygen = -25. hydrogen = +1

6. sum of oxid. # in compound = 0

7. sum of oxid. # in polyatomic ion = charge on ion

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4–4

The Half-Reaction Method (Acidic Solution)

Review- Balancing Oxidation-Reduction Reactions

1. Separate in ½ reactions2. Intermediate steps

a. balance all elements other than H and Ob. balance O with H2O

c. balance H with H+

d. balance charge with (e-)3. Multiply ½ rxn. so that the number of electrons

is same4. Add ½ rxns.

Capture the Energy

MnO4- + 5Fe2+ Mn2+ 5Fe3+

MnO4- and Fe2+ will react directly in solution.

Electrons will be transferred and energy will be released as heat.

No useful work will result.

Capture the Energy! Zn + Cu2+ - Zn2+ + Cu

Separate ½ reactions Connect metals w/ wire (to transfer

electrons)

Connect soln w/ bridge (keeps solns separate but allows ions to move)

Converts Chemical Energy to Electrical Energy!!- A Battery!!

Galvanic Cell

Capture the energy

You have separated the oxidizing agent from the reducing agent

Requires electron transfer through wire

Attach a motor, light bulb, bell etc-the current produced in the wire by e- flow provides work!!

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17–10

Figure 17.6 A Galvanic Cell involving the Half-Reactions

Cell potential is…..

The pressure of a Galvanic cell to “push” the e- “driving force”

Electromotive Force, emf Symbol E Units: Joule/coulomb (=1Volt, V) Coulomb = unit of charge

Specifies # of e-

E cell = E anode +E cathode

(oxidation) (reduction) pushing e- pulling e-

(black wire) (red wire)

A spontaneous rxn in a Galvanic cell must be positive.

E > 0

E 1/2 reactions

P. 796 table Standard Reduction potentials

1M solutions 1atm gases 25 C

Hydrogen ½ rxn = 0.00V

Table 17.1 Standard Reduction Potentials at 25°C (298K) for Many Common Half-Reactions

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17–14

Helpful Info

Need balanced oxidation-reduction rxns from the reduction potentials.

One reduction ½ rxn must be reversed.* The ½ rxn with largest positive potential will

run as written (reduction).The other ½ rxn will run in reverse (oxidation).

Reversing Direction Changes Sign of E

Because:E oxidation = -E reduction

Then: E cell = E cathode – E anode

Examples:

Standard Reduction Potential Math Rules

# of e- lost must equal # e- gained

½ rxns must be multiplied by integers to balance equations

Value of E is not changed when ½ rxn multiplied by an integer.

Potential is NOT multiplied by integer.Example….

Line Notation

Anode listed on leftCathode listed on right

Mg(s) l Mg2+ ll Al3+l Al(s)

Anode Mg0(s) - Mg2+

Cathode Al3+ - Al0(s)

Cell Potential & Free Energy

A galvanic cell will run in the direction that gives a positive value for E

+E corresponds to - G

+E and - G indicates a spontaneous reaction.

G = -n FE

G = nFE

n = # of e- (exchanged in overall rxn)

F = 96,485(c/mol e-) (Faraday’s constant)

Examples:

Effects of Concentration on E

So far the cells have been under standard conditions….

Le Chatelier’s principle applies if not std. conditions..

Determine if E cell > or < E cell ??

To summarize:

If E cell not at standard conditions:

[Reactant] > 1mol/L E cell > E *cell

[Product] < 1mol/L E cell> E *cell

Reverse is also true

Concentration cell

Same components in cells, but different concentrations.

Equilibrium wants these concentrations to be Equal.

Examples:

Nernst Equation

Establishes relationship b/t cell potential and concentration of cell components.

For cells not at 1M Concentration:E = E * - RT/nF ln (Q)

E * is std cell potentialRT/nF ln (Q) is correction factor

Common form:

E = E * - RT/nF ln (Q)

Commonly written :

E = E * - 0.0591/n log (Q)

Examples:

A Battery @ Equilibrium

At Equilibrium:Ecell = 0 (completely discharged)

Q = K and delta G = 0

Using the Nernst Equation:@Equilibrium: 0 =E * - 0.0591/n log(K)Or log K = nE */0.0591

Corrosion Process of returning metals to their natural state.

Metals oxidize readily resulting in corrosion.

Metal ½ rxn is reversed for oxidation. Combined with Oxygen ½ rxn. to give (+) Ecell

Electrolysis

Involves forcing current through a cell to produce a chemical change resulting in (-) cell potential.

Example:

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17–29

Figure 17.19 a-b (a) A Standard Galvanic Cell Based on the Spontaneous Reaction Zn + Cu2+ - Zn2+ + Cu (b) A Standard Electrolytic Cell. A Power Source Forces the Opposite Reaction Cu + Zn2+ - Cu2+ + Zn.