ELECTROCHEMISTRY
A. Electrolysis cell (electrolytic cell) -- apparatus used to perform electrolysis
1. Anode -- positive electrode where oxidation occurs
2. Cathode -- negative electrode where reduction occurs
3. Represent each reaction by a redox half reaction
4. See diagram on page 846
B. Conduction of charge in Electrochemical Cells
1. Electrolytic conduction -- electrical charge is carried through the liquid by the movement of ions, not electrons
2. Negative ions move towards the anode
3. Positive ions move towards the cathode
C. Cell Reactions
1. Cell reaction -- overall reation that takes place in the electrolysis cell
2. Use Hess's Law to combine half reactions into one reaction
3. Write the word "electrolysis" above the arrow to show that electricity is the driving force
D. Electrolysis Reactions in Aqueous solutions
1. Reduction and Oxidation do not always happen with the solute particles
a. sometimes the solute particles (electrolyte) is to maintain electrical neutrality in the vicinity of the electrodes
b. The electrolyte cations are able to move toward the cathode and mingle with the anions that are being formed
c. The electrolyte anions are able to move toward the anode and mingle with cations that are being formed
2. Predicting the outcome of electrolysis reactions in aqueous solutions can be tricky
a. Use what we have learned about one electrolysis to predict what will happen in others
b. The behavior of a particular ion in an aqueous solution toward oxidation or reduction by electrolysis is the same regardless of the source of the ion
II. Stoichiometric Relationships in electrolysis
A. Amount of chemical change during electrolysis is directly proportional to the amount of electrical charge that is passed through an electrolysis cell
1. Ampere (A) -- SI unit for electric current
2. Columb (C) -- SI unit for charge (the amount of charge that passes by a given point in a wire when an electric current of one ampere flows for one second
a. 1 C = 1 A x 1 sec
B. 1 mole of electrons carries a charge of 96485 C
1. 1 mol of electrons is called 1 Faraday
2. Faraday constant -- number of Columbs per mole of electrons
a. 1 F = 9.65 x 10(4) C/mol e-
IV. Galvanic Cells (Voltaic Cells) -- a cell that provides electricity from spontaneous redox reactions in which the electron transfer is forced to take place through a wire
A. Setting Up a Galvanic Cell
1. To produce electrical energy, the two half-reactions involved in the net reaction must be made to occur in separate containers or compartments
a. Half-Cells
2. For a galvanic cell to work, the solutions in both half-cells must remain electrically neutral
3. Cathode (now is positive) -- still reduction occurs
4. Anode (now is negative) -- still oxidation occurs
5. Salt bridge -- a tube filled with an electrolyte solution, commonly KNO3 or KCl and fitted with porous plugs at each end
a. supplies ions to keep solutions neutral
b. electrolytic contact must be maintained for the cell to function
B. Charges on the Electrodes -- label of electrodes determined by the nature of the chemical change, not the electrical charge
1. Elecrolytic Cell
a. Cathode is negative (reduction)
b. Anode is positive (oxidation)
2. Galvanic Cell
a. Cathode is positive (reduction)
b. Anode is negative (oxidation)
3. Electrolytic cells are forced by electrical energy provided from the outside, galvani cells deliver electrical energy spontaneously to the outside
a. Cations always move away from anode toward cathode
b. Anions always move away from cathode toward anode
C. Cell Notation for Galvanic Cells
1. Example --> Cu(s)|Cu2+(aq)||Ag+(aq)|Ag(s)
2. Anode electrode, anode electrolyte, salt bridge, cathode electrolyte, cathode electrode
V. Cell Potentials and Reduction Potentials
A. Potential -- magnitude of the ability of a galvanic cell to push electrons through an external circuit
1. Volts (V) -- units of potential (1 V = 1J/C)
B. Cell Potentials of Galvanic Cells
1. Cell Potential {E(cell)} -- the maximum potential that a given cell can generate
a. depends on the composition of the electrodes, the concentrations of the ions and the temperature
b. Standard cell potential {Eo(cell)} -- potential when all ions are 1.00 M, the temperature is 25oC and any gases are at 1 atm
C. Reduction Potentials -- the magnitude of the tendency of a half-cell to acquire electrons
1. Standard Reduction Potentials
2. When two half-cells are connected, the one with the larger reduction potential acquires electrons from the half-cell with the lower reduction potential, which is therefore forced to undergo oxidation
a. Eo(cell) = (stand. red. pot. of the subs. reduced) - (stand. red. pot. of the subs. oxidized)
D. Assigning Standard Reduction Potentials
1. To assign values to the various standard reduction potentials, a reference has been arbitrarily chosen and its standard reduction potential has been assigned a value of exactly 0 V.
a. Standard hydrogen electrode
2. To obtain the cell reaction, we have to know what is oxidized and what is reduced and which is the cathode and which is the anode
a. determine by measuring the charges on the electrode (in galvanic cells, cathode is positive and anode is negative)
b. a positive reduction potential simply means that the substance is more easily reduced then H+
c. a negative reduction potential simply means that the substance is less easily reduced then H+
d. Table 19.1 on page 870
1. half-reactions at the top have greatest tendency to occur as reduction, while those at the bottom have the least tendency to occur as reduction
VI. Using Standard Reduction Potentials
A. Predicting Spontaneous Redox reactions
1. the half-reaction with the more positive reduction potential always takes place as written (reduction) while the other half-reaction is forced to run in reverse (oxidation)
B. Determining the Cell Reaction and Cell Potential of a Galvanic Cell
1. Use Hess's Law to get the reaction, use the above formula and table 19.1 to get potential
2. An important point to remember: although the half reactions can be multiplied by factors, do not multiply the reduction potentials by thes factors
a. reduction potentials are Joules/Columb, they are dependant upon the number of electrons
C. Determining whether a Reaction is Spontaneous from the Calculated Cell Potential
1. In a galvanic cell, the calculated cell potential for the spontaneous reaction is always positive
2. If the calculated cell potential is negative, the reactionis spontaneous in the reverse direction
D. Using Reduction Potentials to Predict Electrolysis Reactions
1. The half-reaction with the larger reduction potential is more easily reduced at the cathode
2. The half-reaction with the smaller reduction potential is more easily reversed as oxidation at the anode
3. There are occassions when standard reduction potentials do not sucessfully predict electolysis products
VII. Cell Potentials and Thermodynamics
A. Determining a Free Energy Change from a Cell Potential
1. Ch 18 --> -DeltaG = maximum work
2. maximum work = nFE(cell)
a. n = moles, F = Faraday constant, E(cell) = potential in volts
3. Delta G = -nFE(cell)
a. Delta Go = -nFEo(cell)
B. Determining Equilibrium Constants
1. Delta Go = -RTlnKc
2. Eo(cell) = (RT)/(nF)lnKc
a. R = 8.314 J/mol K, T = Kelvin, F = 9.65 x 10(4), n = moles
VIII. Effect of Concentration on Cell Potentials
A. The Nerst Equation: The Relationship of Cell Potential to Ion Concentrations
1. Delta G = Delta Go + RTlnQ
2. -nFE(cell) = -nFEo(cell) + RTlnQ
3. E(cell) = Eo(cell) - (RT)(nF)lnQ
a. Nerst Equation
B. Determining Concentrations from Experimental Cell Potentials
1. See example 19.15
2. since concentrations are usually very small, they can be obtained by simply measuring the potential generated by the electrochemical cell
Outline based upon:
Brady, J. E., Holum, J. R., Russell, J. W. (2000). Chemistry: The Study of Matter and Its Changes. (3rd ed.). New York: John Wiley & Sons, Inc. pp. 845-889.
