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Unit 9:
(Chapter 4 &18)
Electrochemistry
I. Oxidation-Reduction (Redox)
A. The transfer of electrons: Oxidation and Reduction always occur together
1. Oxidation
a. Originally defined as the reaction with oxygen; now there is a more general definition.
b. The loss of electrons (LEO)
c. Oxidation number INCREASES.
C + O2  CO2
2. Reduction
a. Originally defined as the reduction of oxygen atoms.
b. The gaining of electrons (GRR)
c. Oxidation number DECREASES.
CO2  C + O2
B. Oxidation Numbers
1. The actual charge on a monatomic ion; the hypothetical charge on an atom in a molecule.
2. Oxidation numbers can be considered to be the relative charges on the atoms.
C. Rules to Determine Oxidation Numbers:
1. In free elements, each atom has an oxidation state of _______.
(i.e. H2, Br2, Li, etc.)
2. For monatomic ions, the oxidation state is equal to the __________ on the ion.
(i.e. Li+ has an oxidation state of +1)
3. The oxidation state of oxygen in most compounds is ________, but in peroxides its oxidation state is
_________.
4. The oxidation state of hydrogen when it is bonded to a non-metal is ______. When it is bonded to a
metal its oxidation state is ______.
(i.e. HCl = +1; LiH = -1)
5. Fluorine has an oxidation state of ______ in ALL its compounds. Other halogens can have
____________ oxidation states when bonded to oxygen.
6. In a neutral compound, the sum of the oxidation states is ______. In a polyatomic ion, the sum of the
oxidation states is the ___________ of the ion.
D. Assign oxidation numbers to the following:
KClO4
Cr2O72-
CaH2
Na2O2
Fe3O4
OF2
NH3
thiosulfate
CH4
H2O2
arsenite
P4
E. Oxidizing and Reducing Agents
1. Oxidizing agent: the reactant that allows the other substance to be oxidized; it is reduced.
2. Reducing agent: the reactant that allows the other substance to be reduced; it is oxidized.
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F. Oxidation Number of Nonmetals
1. The maximum oxidation number for a nonmetal is the group number. (Max. for Cl = ______).
2. The minimum oxidation number is equal to the group number minus 8. (Min. for Cl = ______).
3. Substances with the nonmetal in its maximum ox. number = oxidizing agent. (NO3- = ______).
4. Substances with the nonmetal in its minimum ox. number = reducing agent. (H2S = ______).
5. Substances with an intermediate oxidation number = oxidizing or reducing agent.
G. Metals as Reducing Agents
1. Metal atoms can oxidize to their ionic (cation) form, which makes them reducing agents.
Na  Na+
Ca  Ca2+
2. Relative strength of metals as reducing agents (activity series of metals):
K
Ca
Na
Ag
Hg
Au
3. Single displacement reaction:
Metal will displace an ion that is below it on the activity series:
(a) K(s) + AgCl(aq) 
(b) Au(s) + NaCl(aq) 
G. Redox Reaction: One substance is oxidized (loses electrons, oxidation number increases), and another
substance is reduced (gains electrons, oxidation number is reduced).
H. Example Reactions:
1. Example: Assign oxidation numbers to the atoms in the following reactions in order to determine the
oxidized and reduced species and the oxidizing and reducing agents:
(a) SnCl2 + PbCl4  SnCl4 + PbCl2
(b) Aluminum powder reacts with ferric oxide.
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(c) Magnesium metal reacts with aqueous cupric chloride (write the molecular and net ionic rxn).
(d) The combustion of propanol.
(e) Decomposition of hydrogen peroxide (Disproportionation Reaction)
II. Balancing Redox Reactions: In order to balance redox reactions, both ____________and ______________must be
conserved! Each half reaction is balanced separately, then they are added to give the overall reaction.
A. Steps to Balance Redox Reactions in Acidic Solutions
1. Assign oxidation numbers to each atom.
2. Split the skeleton reaction into half reactions.
3. Complete and balance each half reaction.
i. Balance all atoms except O and H.
ii. Balance O atoms by adding H2O molecules to one side of the rxn.
iii. Balance H atoms by adding H+ ions to one side of the rxn.
iv. Balance electrons.
4. Multiply each half reaction by a factor to balance the electrons.
5. Combine half reactions to write the full reaction.
6. Simplify reaction by canceling anything that occurs on both side of the rxn.
B. Steps to Balance Redox Reactions in Basic Solutions
1. Balance the rxn as if it were in acidic solution.
2. Add as many OH- ions to both sides as there are H+ ions.
3. Combine H+ and OH- to form water; simplify the reaction.
Example: Balance the following redox reaction in an acidic solution:
MnO4- + I-  I2 + Mn2+
Oxidation half rxn:
Reduction half rxn:
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Example: Balance the following redox reaction in an acidic solution:
Zn + NO3-  Zn2+ + NH4+
Oxidation half rxn:
Reduction half rxn:
Example: Balance the following redox reaction in Basic Solution
MnO4- + SO32-  MnO2 + SO42Oxidation half rxn:
Reduction half rxn:
Example: Balance the following redox reaction in Basic Solution
Br2  Br- + BrO3-
Oxidation half rxn:
Reduction half rxn:
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III. Electrochemical Cells: A contained system in which a redox reaction occurs in conjunction with the passage of
an electric current.
A. Voltaic Cells (Galvanic Cells): A redox reaction that occurs _________________________.
ΔG = ____; EMF (_____________________________, cell potential) = ____
1. These redox reactions can supply ___________ and are used to do _______.
2. The oxidation and reduction reactions are placed in separate containers called half-cells. The halfcells are connected by an apparatus that allows for the flow of ______.
3. Oxidation occurs at the _______________; Reduction occurs at the _________________.
a. An Ox and a Red Cat……….or…………vowels vs. consonants.
b. Because electrons are produced at the anode, it is considered the _____________ electrode.
c. Since the cathode receives electrodes, it is the _________________ electrode.
d. The electrons at the anode have higher potential energy than the electrons in the cathode,
therefore electrons flow from the anode to the cathode (alphabetical order!).
4. Electric potential (like the height of a waterfall) depends on the relative differences of potential energy
of the electrons at the anode versus the cathode. The larger the potential difference, the greater
the ______________. Volts are defined as energy per unit charge.
1 V = 1 J/C
5. A voltmeter measures the potential difference in an electric circuit. The cell potential (voltage) is
known as the Ecell. This is also known as the cell voltage.
6. Electrode: Strip of ________________ used in the electrochemical cell.
7. Example of a voltaic cell:
8. The salt bridge contains an inert salt (ie Na2SO4) which allows for the movement of charged particles
(ions). Without the salt bridge, there would be a buildup of electrons in the ________________, and
excess positive charge would buildup on the _______________. Electrons carry the charge through the
external part of the circuit, but ions carry the charge within the solutions.
9. During the reaction, electrons flow from the zinc anode and through the wire and voltmeter. This
electric current can be used for energy.
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10. Example of a voltaic cell: __________________ car battery. Each cell in the car battery produces
2 V of electricity, so there are six connected cells to produce a 12 V battery.
B. Cell Diagrams: More convenient way to right the reactions of an electrochemical cell than to draw a picture.
1. The anode is written on the left.
2. The cathode is on the right (alphabetical).
3. Single vertical line represents the boundary within the half-cell.
4. A double vertical line represents the salt bridge between the two half-cells.
5. The overall reaction in the above voltaic cell: _____________________________________________
Cell diagram:
6. In many reactions, the electrode only acts as a site for the reaction, but does not take place in the
redox reaction.
Example:
Cl2(g) + 2 e- -------> Cl-(aq)
Cell diagram:
Pt(s)Cl2(g)  Cl-(aq)
Example:
Cl-(aq) -------> Cl2(g) + 2 e-
Cell diagram:
Cl-(aq)Cl2(g) Pt(s)
7. Write the half and overall rxns for the following:
(__________________)
(__________________)
Pt(s)Fe2+(aq)  Fe3+(aq)  Cl2(aq)Cl-(g) Pt(s)
C. Standard Electrode Potentials
1. Individual half-cell potentials cannot be determined, because there must be a potential difference
(waterfall analogy). Therefore in order to determine the cell potentials of half reactions, there must be a
standard reaction that all others are compared to.
2. Standard Hydrogen Electrode (SHE): The zero point of electrode potentials. All other electrode
potentials are relative to this reaction that is arbitrarily assigned a voltage of zero.
Red:
2 H+ (1 M) + 2 e-  H2(g) (1 atm)
E = 0 V (exactly)
Ox:
H2(g) (1 atm)  2 H+ (1 M) + 2 e-
E = 0 V (exactly)
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3. Standard Reduction Potential: Determined relative to the standard hydrogen electrode.
[STANDARD conditions (): All gases at ~1 atm, aqueous solutions 1 M].
4. Calculation of cell potential:
Ecell = Eox + Ered
5. Example: Determine the standard reduction potential of Cu2+.
Oxidation:
Reduction:
Overall:
6. Example: Determine the standard reduction potential of Zn2+.
Oxidation:
Reduction:
Overall:
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7. The more POSITIVE reduction potential, the more likely a substance will undergo reduction.
Example: Which substance is more easily reduced: Zn2+ or Cu2+?
Example: What are the oxidation potentials for Zn and Cu? Which is more likely to oxidize?
D. Standard Reduction Potentials:
1. The oxidation potential is the negative of the reduction potential. (reverse rxn = reverse sign).
2. Substances at the top of the table are EASILY reduced, therefore they make good ______________
agents.
3. Substances at the bottom are EASILY oxidized, so they make good _______________ agents.
4. Example: Organize the following in order of increasing strength of oxidizing agents:
Cl2, NO3-, MnO4-
E. Calculating standard cell potentials:
Ecell = Eox + Ered
1. Example: Calculate the std cell potential of
Pt(s)Fe2+(aq)  Fe3+(aq)  Cl2(aq)Cl-(g) Pt(s).
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2.
(a) Balance the following redox rxn in acidic solution:
O2(g) + I-(aq)  H2O(l) + I2(s)
(b) Calculate the standard cell potential.
3. Determine the standard cell potential for the following:
(a) 2 Al(s) + 3 Cu2+(aq)  2 Al3+(aq) + 3 Cu(s)
(b) Nitrate ion reacts with plumbous ion in acidic solution to form nitrogen monoxide and
plumbic oxide.
(c) Manganous ion reacts with S2O82- ion to form permanganate ion and sulfate ion in
acidic solution.
(d) Zinc metal reacts with hypochlorite ion in basic solution.
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4. Can ferric ion oxidize bromide ion to bromine at standard conditions?
5. Will hydrogen peroxide act as an oxidizing agent or a reducing agent with the following substances
at standard conditions:
(a) I2 / I-
(b) Fe3+ / Fe2+
E. Gibbs’ Free Energy and Spontaneity:
1. Faraday constant (F) is the electric charge per mole of electrons:
1 e- = 1.602 x 10-19 C
1 mol e- = (1.602 x 10-19 C)(6.022 x 1023) = 1 F
1 F = 96,485 C
2. How many Faradays are needed to reduce one mol of nitrate to nitrogen monoxide?
3. Rxn in a voltaic cell  work is done by the electric charges in motion  Free energy is released.
G  nFEcell
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4. n = number of moles of electrons, E = voltage of the cell.
5. If the cell voltage is positive (_________________ cell), then the reaction is spontaneous, therefore
the G= _____. If G is positive, then the reaction requires an input of energy to make it occur
(_______________________ cell). If G = 0, then the reaction is at equilibrium.
6. Determine the G for each of the above reactions 3(a) and 3(c).
(a) 2 Al(s) + 3 Cu2+(aq)  2 Al3+(aq) + 3 Cu(s)
(b) Manganous ion reacts with S2O82- ion to form permanganate ion and sulfate ion in
acidic solution.
7. Will copper metal react with each of the following acids? Explain
(a) Hydrochloric acid
(b) Nitric acid
8. Calculate the G for each of the above reactions.
(a)
(b)
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F. Redox and Equilibrium
1. Voltage is a very precise measurement in the lab, therefore it is a very useful way to determine
ΔG and Keq.
G  nFEcell  RT ln K eq
Ecell 


Ecell
RT
(8.314)(298.15)
ln K eq 
ln K eq
nF
n 96,485
0.025693V

ln K eq
n
2. Example: Calculate the ΔG and Keq at 298 K for the rxn: Cu(s)Cu2+(aq)  Ag+(aq)Ag(s).

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G. Cell Voltage at Non-standard Conditions: The Nernst Equation
1. LeChatelier’s Principle determines the effect of concentration on cell voltage.
2. Example: Zn(s) + Cu2+ (1.0 M)  Zn2+ (1.0 M) + Cu(s)
Ecell = ________
(a) Cu2+ (1.5 M), Zn2+ (0.075 M)
Ecell = 1.142 V
(b) Cu2+ (0.075 M), Zn2+ (1.5 M)
Ecell = 1.064 V
3. Nernst Equation:
E cell  Ecell
0.0592V

log Q
n
4. Example: Calculate the expected voltage for the following voltaic cell:
Cu(s) + Fe3+ (0.20 M)  Cu2+ (0.50 M) + Fe2+ (0.10 M)

5. Example: Calculate the expected voltage of the following voltaic cell:
Zn(s) + Cu2+ (0.050 M)  Zn2+ (2.0 M) + Cu(s)
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H. Concentration Cells:
1. A voltaic cell in which the same substance is used in both the anode and the cathode.
2. The voltage is produced due to a difference in concentrations in the two sides.
3. Example:
anode:
Ni(s)  Ni2+ (aq, 0.00100 M) + 2 eEox = +0.28 V
cathode:
Ni2+ (aq, 1.00 zm) + 2 e-  Ni(s)
Ered = -0.28 V
_______________________________________________________________
overall:
Ecell = ______
Calculate the voltage of the above cell:
4. The reaction will continue until the two sides have ______________ concentrations.
5. Examples of concentration cells:
a. pH meter: Two hydrogen electrodes (one standard, one unknown concentration).
Overall rxn:
2 H+ (1 M)  2 H+ (x M)
b. The heart (pacemaker cells): a concentration gradient of K+ ions creates a potential
difference in these cells. The ratio of K+ ion between the inside and outside of these cells is
about 30:1. If you have too much K+ ion in the blood (excess dietary supplements and kidney
disease) or too little K+ ion (excessive vomiting, etc), the potential difference is altered and
cardiac arrest can result.
I. Fuel Cells: Fuel + oxygen  oxidation products
1. Fuel cells are more efficient at producing electrical energy (up
to 90% efficient).
2.
Anode:
2 [H2(g) + 2 OH-  H2O + 2 e-]
Cathode:
O2(g) + 2 H2O + 4 e-  4 OH__________________________________________
Ecell =________ + ________ = ________
3. Currently the hydrogen gas used in the fuel cell is made
from _____________ (a fossil fuel).
4. The benefit of using hydrogen fuel cells is that the only
product of the combustion is ________.
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IV. Electrolytic Cells
A. A redox reaction that is ____________________ ; ΔG = ____; EMF = ____
1. Electrical energy is required to induce reaction (ie battery).
2. Oxidation and reduction reactions are generally placed in one container.
3. Example: Electrolysis of Water
Anode:
Cathode:
____________________________________________
Overall rxn:
Ecell =________
B. Electroplating
1. Electrolysis used to coat one metal onto another.
2. Example: Electroplating silver onto a spoon.
Anode: Oxidation of oxygen in water
Cathode: Reduction of silver ion to silver metal
Anode:
Cathode:
___________________________________
Overall rxn:
Ecell =________ + ________ = ________
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C. Quantitative Electrolysis
1. Electric current is measured in amperes (A); it is the rate flow of electric charge.
1 A = 1 C/s
2. The gold-plating process involves the following reaction:
Au3+ + 3e-  Au(s).
If 0.600 g of Au is plated onto a metal, how many Coulombs are used?
3. A sample of gold is dissolved and all the gold is converted to Au3+, which is then
reduced back to Au(s) on an electrode of known mass. What mass of gold will be deposited in 1.00 h
by a current of 1.50 A?
4. For how many minutes must the electrolysis of a solution of CuSO4(aq) be carried out, at a current of
2.25 A, to deposit 1.00 g of Cu(s) at the cathode?
5. One type of electrolytic cell contains platinum electrodes in AgNO 3. If 2.175 g of solid silver is
deposited on the cathode in an electrolysis lasting 21 minutes and 12 seconds, (a) what is the amount of
charge (C) passing through the cell? (b) What is the current in amperes?
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