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CH. 17 ELECTROCHEMISTRY
17.1 GALVANIC CELLS
Electrochemistry is the study of reactions that involve a loss of electrons by one species (_____________) and a
gain of electrons by another species (____________). The species being oxidized is called the __________agent
and is considered a donor of electrons. The species being reduced is called the _____________agent and is the
receiver of electrons.
BALANCING REDOX REDOX REACTIONS IN ACIDIC AND BASIC SOLUTIONS
Since electrochemistry is based on oxidation/reduction (redox) reactions. It is critical to be able to balance a redox
reaction. Refer to Ch 4 notes if you do not remember the steps for the following:
REDOX REACTIONS OCCURING IN ACIDIC SOLUTIONS
An acidified solution of potassium dichromate is added to a solution of hydrogen peroxide.
(a) Write the unbalanced net ionic equation.
(b) Determine the species being oxidized and the species being reduced.
(c) Determine the oxidizing agent and reducing agent.
(d) Balance the reaction:
IN BASIC SOLUTIONS Balance the following:
A solution of potassium permanganate is added to a solution of ferrous chloride in a basic solution.
Galvanic or Voltaic Cells
The redox reaction in a galvanic cell is a ____________reaction – which would mean Go ___0 and K __ 1 for a
galvanic cell. For this reason, galvanic cells are commonly used as batteries. Galvanic cell reactions supply energy
which is used to perform work. In a galvanic cell ____________energy is converted to _________ energy. The
energy is harnessed by situating the oxidation and reduction reactions in separate containers, creating a flow of
electrons through an external circuit. The two solid metals that are connected by the external circuit are called
___________ - the ____________ and ___________. Electrons (current) will always flow from the _________
where ___________ occurs to the ___________ where ____________ occurs.
In the galvanic cell pictured above the reaction taking place at the anode is
Zn(s)  Zn2+(aq) + 2eZinc is being __________– meaning it is the ________ agent – the provider/donor of electrons. As a result of the
Zn(s) becoming zinc ions, the zinc metal electrode is being consumed and its mass will decrease over time.
The reaction taking place at the cathode is
Cu2+(aq) + 2e-  Cu(s)
The copper (II) ion is being _________ – meaning it is the ________ agent – the acceptor/receiver or electrons.
As a result of Cu2+ being reduced to copper metal atoms, the mass of the copper electrode will increase over time.
The salt bridge is needed to maintain a balance of charge. The salt bridge will contain a strong ______________.
In the case of electrochemical cells, a common electrolyte used in a salt bridge is sodium nitrate – NaNO3.
If not for the salt bridge, the electrons that are leaving the anode would give the anode an overall positive charge
and as the electrons arrive at the cathode it would give the cathode an overall negative charge and the cell would
stop operating – no electrons would flow through the external circuit. The purpose of the salt bridge is to maintain
a _________ of electric charge – to prevent a build up of positive charge in the anode and negative charge in the
cathode. The cation (for example Na+) will migrate to the ________ to counteract the build up of negative charge
there and the anion (ex. NO3-) will migrate to the ____________ to counteract the build of positive charge there.
* Note – if a solid metal is not part of a an oxidation or reduction half reaction, then an inert electrode such as
platinum (s) or carbon (s) is used to aid in the transfer of electrons. The mass of an inert electrode does not change.
Another depiction of a galvanic cell shows a porous disk instead of a salt bridge to maintain a balance of electric
charge.
Zn(s)  Zn2+(aq) + 2eAnode (Oxidation)
Cu2+(aq) + 2e-  Cu(s)
Cathode (Reduction)
For the Galvanic cells depicted the overall reaction for the entire cell is the sum of the anode ½ reaction and the
cathode ½ reaction. NOTE: The electrons lost during oxidation must _______ the electrons gained by the reduction
reaction.
ANODE (OXIDATION)
Zn(s)  Zn2+(aq) + 2eCATHODE (REDUCTION) Cu2+(aq) + 2e-  Cu(s)
------------------------------------------------------------------OVERALL REACTION
LINE NOTATION (also called CELL NOTATION)
In this notation, the anode components are listed on the left and the cathode components are listed on the
right, separated by double vertical lines (indicating the salt bridge or porous disk). For example, the line notation
for the cell described above is
Note: a phase difference is indicated by a single vertical line and the substance constituting the anode is listed at
the far left and the substance constituting the cathode is listed at the far right.
Example: The following reaction takes place in a galvanic cell : 6H+ (aq) + 2Al (s)  3H2 (g) + 2Al3+ (aq)
A platinum electrode is used in the cell. Write the line notation for the cell.
17.2 STANDARD REDUCTION POTENTIAL
Electrons flow from the anode to the cathode because of a difference in potential energy. The potential
energy of electrons is higher in the anode than in the cathode and they spontaneously flow through an
external circuit from the ____________ to __________________.
The difference in potential energy per electrical charge (potential difference) between two electrodes is
measured in ________. One _______________is the potential difference required to impart 1 joule (J) of energy to
a charge of 1 coulomb (C). (note: one electron has a charge of 1.60 x 10-19 C.)
The potential difference between electrodes in an electrochemical cell provides the driving force that
“pushes” the electrons. It is called the electromotive force, or emf. The emf of a cell (Ecell) is called the cell
potential (sometimes called cell voltage) and is measured in volts.
note: under standard conditions (25°C, 1M, 1 atm) the Ecell is the standard emf, E°cell.
It is convenient to assign a potential to each half reaction so when we construct a cell from a given pair of half
reactions, we can obtain the cell potential (E°cell) by summing the half-cell potentials.
We need one absolute standard against which all other half-reactions can be compared.
The standard is the standard _______________________ electrode.
Table 17.1 in book (or refer to AP handout given) – has a list of standard reduction potentials for many half
reactions. Standard refers to a temperature of ______________, pressure of ______________, and concentration
of _____________.
Here are some things to be aware of in looking at this table:
- All reactions are shown in terms of the reduction reaction relative to the standard hydrogen electrode (SHE)
- The more positive the value of the voltage associated with the half-reaction (Eo), the more
readily the reaction occurs.
- The strength of the oxidizing agent increases as the value becomes more positive, and the
strength of the reducing agent increases as the value becomes more negative.
This table of standard reduction potentials can be used to write the overall cell reaction and to calculate
the standard cell potential (Eocell) , the potential associated with the cell at standard conditions.
To determine the Eo cell , we will use the following:
Since the standard cell potential is for a galvanic cell (involves spontaneous reaction), it must have a positive value.
Eocell ______ 0.
Calculating Standard Cell Potentials (Eocell)
1) Decide which element is oxidized and which is reduced using the table of reduction potentials. THE MORE POSITIVE
REDUCTION POTENTIAL GETS TO BE REDUCED.
2) Write the reduction equation AS IS from the table with its voltage. (E ored)
3) Reverse the equation that will be oxidized and change the sign of the voltage. (E oox)
4) Balance the 2 half-reactions
5) However, since a standard reduction potential is an ______________ property, do not multiple the potential (voltage)
by the integer required to balance the cell reaction.
6) Add the two half-reactions together and add the 2 voltages together.
Example #1: Given the following half-reactions: Cu2+ + e-  Cu+
Au3+ + 3e-  Au
Write the overall reaction for the galvanic cell and calculate Eocell.
Example #2 : Given the following half-reactions: Cu2+ +2 e-  Cu
Ag  Ag+ + eWrite the balance overall reaction and calculate the Eocell. Decide whether or not it’s a galvanic cell.
Oxidizing/Reducing Agents
The strength of the oxidizing agent increases as the Eored value
becomes more positive, and the strength of the reducing agent
increases as the Eored value becomes more negative. Remember,
however, the reducing agent are on the right side of the reduction halfreactions. Reducing agents themselves are being oxidized.
Example Rank the following in order of increasing strengths as
oxidizing agents: Al3+, Co2+, Zn2+, Pb2+
17.3 CELL POTENTIAL AND FREE ENERGY
In Ch 16, we showed the relationship between free energy and the spontaneity of a reaction.
A __________ ∆G indicated a spontaneous reaction. In the last section, we saw that a _____________
cell potential, Ecell, indicated a spontaneous process. The free energy and cell potential can be related in the
following equation.
∆G =
n=
E=
F=
Because both n and F are ______________, a _____________ cell potential (________________ reaction) will
produce a _____________ ∆G ( ___________________ reaction.)
At standard conditions:
Example #1: Calculate the ∆Go for the reaction: Zn2+ (aq) + Cu (s)  Cu2+ (aq) + Zn
Is this rxn spontaneous?
Example #2: Calculate the ∆Go for the reaction: Cu2+ (aq) + Fe (s)  Fe2+ (aq) + Cu
Is this rxn spontaneous?
17.4 DEPENDENCE OF CELL POTENTIAL ON CONCENTRATION
So far we have described cells under standard conditions which are
MOLARITY ________ PRESSURE _________ TEMPERATURE_______
Because cell potentials depend on concentration, we can construct galvanic cells where the molarity of the ion(s) in the ½ cell
are a value other than 1.0-molar.
A relationship was developed by ______________ which can be used to determine the cell potential when ion concentrations
are at a value other than 1.0-molar. The relationship is
Recall from Ch16:
R=
T=
F=
Q=
n=
This equation can also be written using the base 10 logarithm as opposed to the natural log:
NOTE: If the cell is operating at standard temperature, then T k, F, and R, are all essentially constants and a simpler version of
the Nernst equation can be used as follows.
PRACTICE PROBLEM: Determine the Ecell based on the following ½ reactions at 298 K.
VO2+ + 2H+ e-  VO2+ + H2O
Zn2+ + 2e-  Zn
Eored = 1.00 V
Eored = -0.76V
[VO2+] = 2.0 M [H+] = 0.50 M [VO2+] = 0.01 M [Zn2+] = 0.10 M
CONCENTRATION CELLS
A cell in which both compartments have the same components but different _____________ is called a
__________________ _________.
Reaction proceeds in the direction that will ______________ the ion concentration in each half-cell.
The difference in the concentration is the only factor that produces a cell potential, and the voltages are typically
small.
When concentrations are equal, ________________ occurs and Ecell = _____.
Example: A concentration cell is constructed with two Fe(s)/Fe2+(aq) half-cells. One half-cell has [Fe2+] = 0.010 M and the
other has [Fe2+] = 0.10 M. Determine the anode and cathode and the E cell.
CALCULATION OF EQUILIBRIUM CONSTANTS FOR REDOX REACTIONS
The quantitative relationship between E0 and ∆G0 allows calculations of equilibrium constants for redox reactions.
For a cell at equilibrium, Ecell = ______and Q ______K.
Applying these conditions to the Nernst equation valid at 250C:
EXAMPLE: For the oxidation-reduction reaction:
S4O62-(aq) + Cr2+(aq)  Cr3+(aq) + S2O32-(aq)
The appropriate ½ reactions are
S4O62-(aq) + 2e-  2 S2O32-(aq) Eored = 0.17V
Cr3+(aq) + e-  Cr2+(aq)
Eored = -0.50V
Determine Eocell
Write the overall balanced equation for the cell:
Determine the value of K for the cell.
PAST AP FR ESSAY
The following questions refer to the electrochemical cell shown in the diagram above.
(a) Write a balanced net ionic equation for the spontaneous reaction that takes place in the cell.
(b) Calculate the standard cell potential, E°, for the reaction in part (a).
(c) In the diagram above,
(i) label the anode and the cathode on the dotted lines provided, and
(ii) indicate, in the boxes below the half-cells, the concentration of AgNO3 and the concentration of
Zn(NO3)2 that are needed to generate E°.
(d) How will the cell potential be affected if KI is added to the silver half-cell? Justify your answer.
17.5 BATTERIES AND 17.6 CORROSION
The AP Chem test does not heavily stress these sections, I will go over these sections briefly when we
review for the AP Exam.
17.7 ELECTROLYSIS
Galvanic cells are based on ____________________ redox reactions which produce _____________
energy. Conversely it is possible to use ________________ energy to cause _______________________
redox reactions to occur. Such processes which are driven by an outside source of electrical energy are called
________________________ reactions and take place in ______________________ cells.
An electrolytic cell is often used for the electroplating of a metal on another object.
Example: Electroplating a paper clip with copper
Requirements to accomplish this:
1) The object to be plated (to receive the metal) must be at the ___________. Why?
2) The metal that is to be plated on the object must be at the _______________.Why?
3) In solution, you need to have a salt solution of the metal that is being plated on the object.
4) You need a ______________________ to accomplish this reaction.
Sketch the electrolytic cell needed to accomplish this.
Example Problem: Determine the mass of copper that is plated on the paper clip if a current of 10.0 amps
is passed for 30.0 minutes through a solution of copper (II) sulfate.
Note: The unit of electric current (I) is the ampere (amp). I = ____________. “q” is charge in
_____________________ and “t” is time in ___________________. Therefore 1 A =
_______________.
To solve this problem – we will use the following steps (done in dimensional analysis.)
Time(s)  total charge  moles of e-  (mole ratio) moles of metal to be plated  grams of metal
ELECTROLYSIS HALF REACTIONS
In an electrolytic cell, a ____________________________reaction occurs, and Eocell is __________________ and
∆G is __________________________. By using an externally supplied ______________ source, we can get these
non-spontaneous reactions to go forward
Thus, for an electrolysis reaction, (using the table of standard reduction potentials) - the half-reaction with the
greater (more positive) reduction potential will be oxidized and the half-reaction with lower reduction potentials
will be reduced.
Higher Reduction Potential

Lower Reduction Potential

Example #1 Electrolysis of H2O (l)
2 H2O (l) + 2 e-  H2 + 2OHO2 (g) + 4 H+ + 4 e-  2 H2O (l)
E0red = -0.83 V
E0red = 1.23 V
_________________gas is produced at the anode, ___________________ at the cathode.
Example #2 Electrolysis of NaF (l)
Based on the table of reduction potentials, we predict that fluorine gas will spontaneously react with sodium metal. The
fluorine will be reduced and the sodium will be oxidized in this spontaneous reaction forming ___________.
Using an electrolytic cell we can run the reverse reaction, reacting sodium fluoride to make _______________ and
_____________________. Solid sodium fluoride does not conduct electricity since the ions are fixed in their crystal
lattice. So we need liquid (molten) sodium fluoride to conduct an electric current.
Net Ionic:
Reduction ½ reaction:
Oxidation ½ reaction:
Balanced net ionic:
_________________ gas is produced at the _______________________.
Example #3 Electrolysis of NaF (aq)
Sodium fluoride dissolved in water will also conduct an electric current. However, with water present, we have it
as a possible reactant along with H+, OH- , H2, and O2.
Possible reduction ½ reactions: (the one with the higher (more positive) E0red will occur)
Possible oxidation ½ reactions: (the one with the higher (more positive) E0oxid will occur)
Overall net ionic:
Helpful Hints for Predicting Products of Electrolysis
If there is no water present and you have a pure molten ionic compound, then:
- the cation will be reduced
- the anion will be oxidized
If water is present and you have an aqueous solution of the ionic compound, then:
- you need to figure out if the ions are reacting or water is reacting
 no group 1 or 2 metal will be reduced in an aqueous solution, water will be reduced instead.
 no polyatomic ion will be oxidized in an aqueous solution, water will be oxidized instead.
HOMEWORK – (separate sheet of paper)
HOMEWORK – 3/7 (Galvanic Cells)
1) The nickel-cadmium (NiCad) battery, a rechargeable “dry cell” used in battery-operated devices, uses the following redox
reaction to generate electricity: Cd(s) + NiO2 (s) + 2 H2O(l) → Cd(OH)2(s) + Ni(OH)2(s)
Identify the species being oxidized and species being reduced. Identify the oxidizing and reducing agents.
2) Balance the following reactions in an acid solution:
(a) An acidified solution of potassium dichromate is added to a solution of sodium iodide.
(b) Aluminum metal is added to a potassium permanganate solution.
3) Balance the following reaction in basic solution:
In a basic solution sodium phosphite is mixed with a solution of potassium permanganate. NOTE: The phosphite
ion will become the phosphate ion.
4) Sketch a spontaneous galvanic cell (using salt bridge) for following ½ reactions
Ag+ + e-  Ag
Fe2+ + 2e-  Fe
HOMEWORK – 3/18 (Eocell , ∆)
pp. 880-882 26b, 34b, 35, 36, 38 (refer to 26b), 47,48, 55, and 57
Will need to use stand. reduction table in book (p 843) for a few of these!
HOMEWORK 3/19
AP FRQ #1
HOMEWORK – 3/21
1. The electrolysis of BiO+ produces pure bismuth. How long (in minutes) would it take to produce 10.0 g
of Bi by the electrolysis of a BiO+ solution using a current of 25.0 A?
2. Electrolysis of a molten metal chloride (MCl3) using a current of 6.50 A for 1397 seconds deposits
1.41 g of the metal at the cathode. What is the metal?
3. Show the ½ reactions and balanced net ionic equation for the electrolysis of Zn(NO3)2 (aq).
What gas is produced at the anode?