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Electrochemistry Investigation
Chemistry 226H Notes
1.
Dr. Richard Nafshun
Oregon State University
Consider C6H12O6 (aq) + [Re2Cl8]2- (aq) → Re2+ (aq) + CO2 (g) in acid. Balance and
identify details (species reduced, oxidized...)
Step 1: Identify the half reactions
C6H12O6 (aq) → CO2 (g)
Loose the chloride ion (let’s call it spectator)
2 Re3+ (aq) → Re2+ (aq)
Step 2: Balance all atoms except H and O
C6H12O6 (aq) → 6 CO2 (g)
2 Re3+ (aq) → 2 Re2+ (aq)
Step 3: Balance O atoms by adding water molecules
6 H2O (l) + C6H12O6 (aq) → 6 CO2 (g)
2 Re3+ (aq) → 2 Re2+ (aq)
Step 4: Balance H atoms by adding protons (hydrogen ion)
6 H2O (l) + C6H12O6 (aq) → 6 CO2 (g) + 24 H+ (aq)
2 Re3+ (aq) → 2 Re2+ (aq)
Step 5: Balance change by adding electrons
6 H2O (l) + C6H12O6 (aq) → 6 CO2 (g) + 24 H+ (aq) + 24 e2 e- + 2 Re3+ (aq) → 2 Re2+ (aq)
Step 6: Use multipliers so the number of electrons lost equals the number electrons gained
6 H2O (l) + C6H12O6 (aq) → 6 CO2 (g) + 24 H+ (aq) + 24 e12 x [2 e- + 2 Re3+ (aq) → 2 Re2+ (aq)]
24 e- + 24 Re3+ (aq) + 6 H2O (l) + C6H12O6 (aq) → 24 Re2+ (aq) + 6 CO2 (g) + 24 H+ (aq) + 24 e-
2.
Consider the cell: Cu0 (s) | Cu2+ (aq) || Ag+ (aq) | Ag0 (s) at 25ºC
anode
oxidation
salt bridge
cathode
reduction
What is the cell potential (volts) of this cell constructed of 1.0 M solutions?
Cu0 (s) → Cu2+ (aq) + 2 e-
e- + Ag+ (aq) → Ag0 (s)
The half reactions above show that copper metal is giving up two electrons and silver ion is
gaining one electron. When these two half reactions are combined, two electrons are transferred:
Cu0 (s) → Cu2+ (aq) + 2 e2 x [e- + Ag+ (aq) → Ag0 (s)]
_______________________
2 e- + 2 Ag+ (aq) + Cu0 (s) → Cu2+ (aq) + 2 e- + 2 Ag0 (s)
Cancel the electrons on both sides. We are transferring electrons: electrons leave the copper
compartment and travel through the external circuit to the silver compartment.
2 Ag+ (aq) + Cu0 (s) → Cu2+ (aq) + 2 Ag0 (s)
Each Ag+ (aq) is gaining an electron
Ag+ (aq) is reduced
Ag+ (aq) is the oxidizing agent; it is taking an electron from Cu0 (s)—note that two silver ions
are required to remove two electrons from Cu0 (s) resulting in the formation of Cu2+ (aq)
_____
Each Cu0 (s) is losing an electron
Cu0 (s) is oxidized
Cu0 (s) is the reducing agent; it is giving two electrons to two Ag+ (aq)
_____
The table (standard reduction potential table) depicts (shows, pictures, illustrates, displays,
represents, portrays, describes):
Ag+ (aq) + e- → Ag(s)
Fe3+(aq) + e- → Fe2+(aq)
O2(g) + 2H+ (aq) + 2 e- → H2O2(aq)
I2(s) + 2 e- → 2 I- (aq)
Cu2+(aq) + 2 e- → Cu(s)
+0.800
+0.771
+0.695
+0.535
+0.340
The difference between the silver and copper half reactions is 0.460 V—this is the cell potential.
After some use, one solution increases to 1.10 M. Which solution is this?
According to the balanced reaction: 2 Ag+ (aq) + Cu0 (s) → Cu2+ (aq) + 2 Ag0 (s)
Ag+ (aq) is being consumed and Cu2+ (aq) is being generated. Therefore, Cu2+ (aq) is 1.10 M.
What is the concentration of the other solution?
The concentration of Ag+ (aq) is 0.80 M because the concentration of Cu2+ (aq) increased by
0.10 M (from 1.00 M to 1.10 M) and the stoichiometry of silver ion to copper ion is 2:1.
What is the cell potential?
For equal molar solutions at standard conditions, as discussed above, the cell potential is the
difference between +0.800 V and +0.340 V or +0.460 V. For non-equal molar solutions, we use
the Nernst Equation:
E cell  E  cell 
RT
ln Q
nF
where Q is the reaction quotient (
E cell  E  cell
or
E cell  E  cell 
0.0592
log Q
n
products
).
reacatants
RT [Cu 2 (aq)]

ln
 0.460 V 
nF [Ag  (aq)] 2
J
)( 298 K )
1.10 M
mol  K
ln
Coulombs
[0.80 M]2
(2)(96,485
)
mole  e 
(8.314
E cell  0.460 V  0.00V  0.456 V
The cell potential drops as the cell reaction proceeds because of the increase in Cu2+ (aq)
concentration.
3.
Discuss rust and cathodic protection.
Rust: http://en.wikipedia.org/wiki/Rust
Rust is the substance formed when iron compounds corrode in the presence of oxygen and water. It is a
mixture of iron oxides and hydroxides. Rusting is a common term for corrosion, and usually corrosion of
steel.
Iron is found naturally in the ore haematite as iron oxide, and metallic iron tends to return to a similar
state when exposed to air, (hydrogen, oxygen, nitrogen, etc.) and water. This corrosion is due to the
oxidation reaction when iron metal returns to an energetically favourable state. Energy is given off when
rust forms. The process of rusting can be summarized as three basic stages: The formation of iron(II)
ions from the metal; the formation of hydroxide ions; and their reaction together, with the addition of
oxygen, to create rust.
Iron is the main component of steel and the corrosion of steel is observed more frequently, since iron is
rarely used without alloying in the present day.
When steel contacts water, an electrochemical process starts. On the surface of the metal, iron is
oxidized to iron(II):
Fe → Fe2+ + 2e−
The electrons released travel to the edges of the water droplet, where there is plenty of dissolved
oxygen. They reduce the oxygen and water to hydroxide ions:
4e− + O2 + 2H2O → 4OH−
The hydroxide ions react with the iron(II) ions and more dissolved oxygen to form iron oxide. The
hydration is variable, however in its most general form:
Fe2+ + 2OH− → Fe(OH)2
4Fe(OH)2 + O2 → 2(Fe2O3.xH2O) + 2H2O
Hence, rust is hydrated iron(III) oxide. Corrosion tends to progress faster in seawater than fresh water
due to higher concentration of sodium chloride ions, making the solution more conductive. Rusting is
also accelerated in the presence of acids, but inhibited by alkalis. Rust can often be removed through
electrolysis, however the base metal object can not be restored through this method.
CP:
http://en.wikipedia.org/wiki/Cathodic_protection
Cathodic protection (CP) is a technique to control the corrosion of a metal surface by making that
surface the cathode of an electrochemical cell.
It is a method used to protect metal structures from corrosion. Cathodic protection systems are most
commonly used to protect steel, water, and fuel pipelines and tanks; steel pier piles, ships, and offshore
oil platforms.
A side effect of improperly performed cathodic protection may be production of molecular hydrogen,
leading to its absorption in the protected metal and subsequent hydrogen embrittlement.
Cathodic protection is an effective method of preventing stress corrosion cracking.
Origins
The first use of CP was in 1824, when Sir Humphry Davy, of the British Navy, attached chunks of iron
to the external, below water line, hull of a copper clad ship. Iron has a stronger tendency to corrode
(rust) than copper and when connected to the hull, the corrosion rate of the copper was dramatically
reduced.
Galvanic CP
Today, galvanic or sacrificial anodes are made in various shapes using alloys of zinc, magnesium and
aluminium. The electrochemical potential, current capacity, and consumption rate of these alloys are
superior for CP than iron.
Galvanic anodes are designed and selected to have a more "active" voltage (technically a more negative
electrochemical potential) than the metal of the structure (typically steel). For effective CP, the potential
of the steel surface is polarized (pushed) more negative until the surface has a uniform potential. At that
stage, the driving force for the corrosion reaction is halted. The galvanic anode continues to corrode,
consuming the anode material until eventually it must be replaced. The polarization is caused by the
current flow from the anode to the cathode. The driving force for the CP current flow is the difference in
electrochemical potential between the anode and the cathode.
Impressed Current CP
For larger structures, galvanic anodes cannot economically deliver enough current to provide complete
protection. Impressed Current Cathodic Protection (ICCP) systems use anodes connected to a DC power
source (a cathodic protection rectifier). Anodes for ICCP systems are tubular and solid rod shapes or
continuous ribbons of various specialized materials. These include high silicon cast iron, graphite, mixed
metal oxide, platinum and niobium coated wire and others.
A cathodic protection rectifier connected to a pipelineA typical ICCP system for a pipeline would
include an AC powered rectifier with a maximum rated DC output of between 10 and 50 amperes and 50
volts. The positive DC output terminal is connected via cables to the array of anodes buried in the
ground (the anode groundbed). For many applications the anodes are installed in a 60 m (200 foot) deep,
25 cm (10-inch) diameter vertical hole and backfilled with conductive coke (a material that improves the
performance and life of the anodes). A cable rated for the expected current output connects the negative
terminal of the rectifier to the pipeline. The operating output of the rectifier is adjusted to the optimum
level by a CP expert after conducting various voltage measurements.
Testing
Electrochemical potential is measured with reference electrodes. Copper-copper(II) sulfate electrodes
are used for structures in contact with soil or fresh water. Silver chloride electrodes are used for seawater
applications.
Galvanized Steel
Most modern cars have galvanized (zinc-coated) steel frames and panels. Unprotected steel forms a
layer of iron oxide, which is permeable to air and water and allows corrosion to continue underneath.
However, zinc oxide (produced on the surface of zinc-protected objects) is impermeable (see
passivation). As long as the zinc and zinc oxide layers are undisturbed (i.e. not scraped or sanded off),
the steel underneath will not rust.
Galvanised steel has some self repairing properties; small scratches where the steel becomes exposed
will be recovered by the zinc. This happens because the zinc from the surrounding area will dissolve and
be deposited on the steel, replacing what was lost to the scratch
4.
Consider the colors of transition metal complexes.
The colors of transition metal complexes are due to electronic transitions within the d-orbitals.