Cell Chemistry
In 1771, Luigi Galvani discovered that the muscles of dead frog’s legs twitched when struck by a spark. Galvani's assistant touched an exposed sciatic nerve of the frog with a metal scalpel and the dead frog's leg kicked as if alive. Galvani coined the term animal electricity to describe the force that activated the muscles of his specimens. The
phenomenon was dubbed galvanism on the suggestion of his intellectual adversary Alessandro Volta. However Volta reasoned that animal electricity was a physical phenomenon, metallic electricity if you will.
Investigations by Volta led to the invention of an early battery, the Voltaic Pile (see below). Volta built the first battery to specifically disprove Galvani's theory of animal electricity.
An Original Voltaic Pile
Schematic of the Voltaic Pile The electrolyte was brine, NaCl(aq),
A simplified schematic of an electrochemical cell containing a spontaneous reaction is provided below. Once called a Galvanic Cell they are today called Voltaic Cells. The anode is the site of
oxidation – ALWAYS.
Anions always flow through the electrolyte to the anode
A voltaic cell has a
positive cell potential. It does not always occur, but electrode corrosion
can only occur at the
anode. corrosion is the oxidation of a metal element to its ion at the anode.
The cathode is the site of reduction – ALWAYS.
Cations always flow through the electrolyte to the cathode
An electrolytic cell has a
negative cell potential. It does not always occur, but electroplating can only occur at the
cathode. Electroplating is the reduction of a metal ion to its elements at the cathode
SOAC/GERC
Electrons always flow from the anode to the cathode of a cell.
For voltaic cells the anode is the negative electrode and the cathode is the positive electrode. For cells that contain a non-spontaneous reaction, electrolytic cells, the charges are opposite.
A standard voltaic cell operates at 25 °C with electrolytes at 1.0 mol/L concentration and all gaseous components at standard pressures.
A standard voltaic cell comprises two standard half cells/electrodes. A standard half cell contains a redox couple.
A redox couple comprises the reduced and oxidized forms of a chemical entity.
The 3 ways of schematically representing electrochemical cells are:
In practice, a schematic a voltaic cell in the laboratory setting looks as follows:
A more detailed diagram, such as would appear on a diploma examination will look like:
The two halves of the cell are obvious. The upturned tube connecting the two
Cell Nomenclature explained/demystified
C(s) | H
+(aq), NO
3–(aq) || Cu
2+(aq), NO
3–(aq) | Cu(s)
1st
electrode
line 1st electrolyte Porous
membrane
2nd electrolyte line 2nd
electrode The first, and the last, single vertical lines/solidus in the nomenclature represent the separation of the solid electrode from the electrolyte. The middle double vertical lines in the nomenclature represent the means of separating the liquid electrolytes while at the same time allowing ions to flow between the two half cells. This porous membrane can take many forms. Illustrated above is the salt bridge. This is a U shaped glass tube filled with a relatively inert electrolyte solution which will allow ions to flow between the electrolytes. Porous plugs of cotton, essentially cellulose, prevent the electrolyte
escaping from the upturned U-tube. A paper towel soaked in KNO3(aq) would also make
a reasonable salt bridge. Convention stipulates that the anode is written first.
The C(s) electrode is referred to as an inert electrode. An inert electrode is there solely to allow transport of electrons into, or out of, a half-cell and not to take part in any electrochemical reaction. Note: an inert electrode is not a catalyst!
A similar cell could be constructed with a porous membrane that was in fact a small, unglazed, ceramic beaker, called a porous cup or porous pot – see below
The nomenclature of this cell would be:
C(s)
|
H+(aq), K+(aq), MnO4–(aq)||
Cu2+(aq), SO42–(aq)|
Cu(s)Not all the chemical components in each electrolyte need to be specified. It is perfectly appropriate to designate only the relevant reactants and electrode components
The nomenclature of this cell would be: Zn(s)
|
Zn2+(aq)||
Cu2+(aq)|
Cu(s)Calculating Net Cell Potential, Eºcell
Once the anode and cathode identities of a cell or electrochemical reaction are identified the net cell potential is the difference between the cathode and the anode reduction potentials as follows
Eº
cell= Eº
cathode– Eº
anodeThe standard net potential of a voltaic cell, a cell which contains a spontaneous reaction, is always positive.
Examples : Calculate the net cell potential of the cells described by the following cell nomenclatures.
1. C(s)
|
H+(aq), NO3–(aq)||
Cu2+(aq), NO3–(aq)|
Cu(s)Looking up the table of selected electrochemical potentials in the Alberta Govt. Chemistry Data booklet gives the reduction potentials of each half cell as follows C(s)
|
H+(aq), NO3–(aq)
||
Cu2+(aq), NO3–(aq)|
Cu(s)Eº = +0.80 V Eº = +0.34 V
Since the C(s)
|
H+(aq), NO3–(aq) half contains the strongest oxidizing agent of themixture of chemical cell components this half cell is the cathode, Eº cathode = +0.80 V Since the Cu2+(aq), NO3–(aq)
|
Cu(s) half contains the strongest reducing agent of themixture of chemical cell components this half cell is the anode, Eº anode = +0.34 V Thus: Eºcell = Eºcathode – Eºanode = +0.80 V – (+0.34 V) = +0.46 V
In other words we can expect the cell with the nomenclature in question will produce a cell potential of +0.34 V when all the electrolyte components are at 1.0 mol/L and 25 ºC.
2. C(s)
|
H+(aq), K+(aq), MnO4–(aq)||
Cu2+(aq), SO42–(aq)|
Cu(s)Eº = +1.51 V Eº = +0.34 V
Since the C(s)
|
H+(aq), K+(aq), MnO4–(aq)half contains the strongest oxidizing agent ofthe mixture of chemical cell components this half cell is the cathode, Eº cathode = +1.51 V Since the Cu2+(aq), SO42–(aq)
|
Cu(s) half contains the strongest reducing agent of themixture of chemical cell components this half cell is the anode, Eº anode = +0.34 V Thus: Eºcell = Eºcathode – Eºanode = +1.51 V – (+0.34 V) = +1.17 V
In other words we can expect the cell with the nomenclature in question will produce a cell potential of +1.17 V when all the electrolyte components are at 1.0 mol/L and 25 ºC.
3. Zn(s)
|
Zn2+(aq)||
Cu2+(aq)|
Cu(s)From the chart Eº(Zn(s) | Zn2+(aq))= –0.76 V= Eº
anode From the chart Eº(Cu2+(aq)
|
Cu(s))= +0.34 V= Eºcathode
Thus: Eºcell = Eºcathode – Eºanode = +0.34 V – (–0.76 V) = +1.10 V
We can expect the cell with this nomenclature to produce a cell potential of +1.17 V when all the electrolyte components are at standard conditions.
Note: In cells 1-3 the Cu2+(aq)
|
Cu(s) was an anode or an cathode depending on electrical1. Identify the electrodes clearly and show a net cell potential (Eºcell) calculation for the
cells described by the following cell nomenclatures.
a. Cr(s) | Cr2+(aq) || Sn2+ (aq) | Sn(s) _____
b. C(s) | SO42–(aq), H+(aq), H2SO3(aq) || Co2+ (aq) | Co(s) _____
c. Ag(s) | Ag+(aq) || MnO4–(aq), H+(aq), Mn2+ (aq) | Pt(s) _____
d. Pt(s) | H2(g), OH–(aq) || O2(g),H+(aq) || Pt(s) _____
e. Pt(s) | Cr2O7–(aq), H+(aq), Cr3+ (aq) || Pb2+ (aq) | Pb(s) _____
f. Fe(s) | Fe2+ (aq) || Cr2+ (aq) | Cr(s) _____
Drawing Cell Diagrams
Draw complete diagrams of the following voltaic cells (include: electrolyte components; ion flows; electron flow; all half & net reaction equations; Eºcell; anode/cathode evidence of reaction; anode/cathode labels & charges. For clarity's sake use a salt bridge not a porous cup/pot.
a. Cu(s) | Cu2+(aq) || Ag+(aq) | Ag(s)
c. Cd(s) | Cd2+(aq) || H+(aq), NO3–(aq) | Pt(s)
Standard Electrode Potentials
1. A standard indium-gold voltaic cell is constructed and its Eºcell = + 1.84 V. The gold
electrode was observed to be the cathode. What is the reduction potential for the indium half cell?
2. Identify the unknown half cells in the following standard voltaic cells (in each case the first half cell is the anode).
a. Zn(s) | Zn2+(aq) || X2–(aq), Y+(aq), Z(aq) | C(s) Eº
cell = +0.93 V
b. X(s) | X2+(aq) || MnO4– (aq), H+(aq), Mn2+(aq) | C(s) Eºcell = +1.17 V
c. Ni(s) | Ni2+(aq) || X–(aq), Y2(aq), | Pt(s) Eºcell = + 1.33 V
3. Calculate the reduction potential of the iodine redox couple if the standard reference electrode is no longer the S.H.E. but is the
a. lithium half cell b. fluorine half cell
c. mercury half cell d. nickel half cell
4. Create a table of reduction potentials given the following information.
Problem
What is the relative strength, in decreasing order, of four oxidizing agents?
Experimental Design
Several cells are investigated; each cell has at least one half-cell in common with one of the other cells. The cell potentials are measured and the electrodes identified.
Evidence
Cathode Anode Eºcell (V)
C(s) | Cr2O72–(aq), H+(aq) Pd(s) | Pd2+(aq)
+0.28 Tl(s) |Tl+(aq) Ti(s) |Ti+(aq)
+1.29 Pd(s) | Pd2+(aq) Tl(s) |Tl+(aq)
+1.29
5. The mercury cell is a special cell for products such as watches and hearing aids. The following reactions are involved in the operation of the cell.
ZnO(s) + H2O(1) + 2 e– Zn(s) + 2 OH–(aq) Eº = –1.25 V
HgO(s) + H2O(1) + 2 e– Hg(l) + 2 OH–(aq) Eº = +0.10 V
Determine the net reaction equation and the net potential for this cell.
6. In a methane fuel cell, the chemical energy of this compound is converted into electrical energy instead of the heat that would flow during the combustion of
methane. Using only the following half-reactions and reduction potentials, write a net reaction equation and determine the approximate potential for the methane fuel cell.
CO32– (aq) + 7 H2O(l) + 8 e– CH4(g) + 10 OH– (aq) Eº = +0.17 V
Question
1. Concern about increased air pollution and the increasing use of non-renewable
resources has accelerated research into alternatives to the internal combustion engine. One alternative is a battery-powered electric motor. Several efficient cells are being tested. The diagram below represents one of these cells.
The Aluminum- Air Cell
a. What role does the oxygen play in this cell?
b. Write the appropriate half and net reactions for this cell.
c. Estimate a standard potential difference of this cell.
Electrolytic Cells
Electrolytic cells are electrochemical cell that house a non-spontaneous chemical reaction – cells that will work only with power supplied from a D.C. power source (DC = direct current). While voltaic cells house chemical reactions that produce energy in the form of an electric current the opposite process occurs in electrolytic cells. In fact, electrolytic cells are used in industry to produce chemicals upon the application of an electric current. The prediction method for electrolytic cells is in all respects identical to that for voltaic cells. The actual design of the cell is another matter. Generally speaking there's no need for the electrodes of an electrolytic cell to be different. Often the electrodes are inert (in the laboratory this means using carbon or platinum foil electrodes; in an industrial setting other factors dictate the choice of electrode material). Electrolytic cells usually operate with a single electrolyte and a porous barrier separating the two electrodes is employed to keep reactive products apart but is not always needed.
Provide a suitable cell diagram and predict the cell reactions, electron flow, ion flow, anode/cathode, Eºcell, and all evidence of operation for the electrolysis of magnesium
sulfate using carbon electrodes.
The cell nomenclature for the electrolytic cell is: C(s) | Mg2+(aq), S042–(aq), H2O(l) | C(s)
The net and half reactions are as follows:
Cathode: 4 H2O(l) + 4 e–2 H2(g) + 4 OH–(aq) Eºcathode= –1.25 V
Anode: 2 H2O(l) O2(g) + 4 H+(aq) + 4 e– Eºanode = –1.25
Net Cell Equation: 2 H2O(l) 2 H2(g) + O2(g) Eºcell = –2.06 V
(Eºcell = Eºcathode– Eºanode = –0.83 V – (–1.23 V) = –2.06 V)
This cell will operate with the predicted chemistry if a minimum voltage of 2.06 V is applied from a DC power source. Evidence:
A gas that explodes with a squeaky pop will be produced at the cathode
The catholyte pH will increase {made OH–(aq)}
A gas that relights a glowing splint will be produced at the anode (oxygen)
• the anolyte pH will decrease
Electrolytic Cells
1. Predict the cathode and anode half- reactions, the net reaction, the Eºcell and the cathode and anode evidence of reaction/operation for the electrolyses (inert electrodes) of aqueous solutions containing the following compounds: a. cobalt(II) iodide
b. potassium sulfate
c. zinc bromide
e. nickel(II) nitrate
f. sodium hydroxide
3. Predict the cathode and anode half- reactions, the net reaction, the Eºcell and the cathode and anode evidence of reaction/operation for the electrolyses (inert
electrodes) of aqueous copper(II) sulfate solution using a copper anode and a copper cathode.
4. If you wish to electroplate an object with a metal, within an electrolytic cell, at which electrode in the cell should it be located? Explain.
Industrial Electrolysis (and other processes)
Barely a few years after Alessandro Volta developed battery technology in ca 1800 scientists in Britain passed an electronic through an acid solution and produced hydrogen and oxygen gases. Humphrey Davy, a renowned 19th century chemist, used electrolysis to discover a large number of metal elements and prepare pure samples of known metals. This chemical technology, electrolysis, is one of the most important means in use today of purifying metals.
1. Aluminium: The Hall-Héroult process. Alumina, Al2O3(s), is virtually insoluble in
water and melts at too high a temperature (>2000 ºC) to be economically recoverable by electrolysis of its aqueous solution or its melt. Instead alumina dissolves and dissociates in the molten mineral cryolite, Na3AlF6(1), at ca 950 ºC.
The aluminium produced at the cathode sinks to the bottom of the electrolyte and is removed continuously while the oxide ions reacts with carbon at the anodes to give CO2(g) as follows (“cry” signifies dissolved in cryolite):
Cathode: 4 Al3+(cry) + 12e– 4 Al(l) Anode : 6 O2– ( cry) + 6 C(s) 6 CO
2(g) + 12e–
Net : 4 Al3+(cry) + 6 O2– ( cry) + 6 C(s) 4 Al(l) + 6 CO2(g)
2. Sodium-The Downs’ Cell : Sodium is obtained by the electrolysis of molten sodium chloride in what is called, after its inventor J. Cloyd Downs, a Downs’ Cell.
The cell chemistry is as follows:
Cathode: 2 Na+(l) + 2 e– 2 Na(l) Anode : 2 Cl–(l) Cl2(g) + 2 e–
Net : 2 Na+(I) + 2 Cl–(l) 2 Na(l) + Cl2(g)
3. Chlorine: Chlor Alkali Process: Chlorine is produced on the large scale from the direct electrolysis of brine, a saturated aqueous solution of NaCl(aq). Although for this process Eºcell = –2.19 V, the operating voltage of is closer to 3 V. An applied
potential greater then the minimum necessary is called overpotential. It has been known for nearly 200 years that chlorine is the principal anode product. The anode product ought to be oxygen from the oxidation of water (Eº= +1.23 V) because it is a stronger reducing agent than chloride (Eº= +1.36 V). However the oxidation of chlorine has far lower activation energy than the oxidation of water. This means that chloride ion is oxidized at a far lower overpotential than water.
asbestos membrane
The asbestos membrane between the two half cells is a very polar tightly woven fabric. The H2(g) and Cl2(g) are collected in a way that does prevents them mixing
and reacting explosively. In solution they will not migrate through the membrane. The sodium ions pass through the membrane easily. The chemistry is as follows:
Cathode: 2 H2O(1) + 2 e– H2(g) + 2 OH–(aq)
Anode : 2 Cl–(aq) Cl2(g) + 2 e–
Net Cell : 2 H2O(1) + 2 Cl–(aq) H2(g) + Cl2(g) + 2 OH–(aq)
4. Purifying Copper: Copper Electrowinning. Impure copper, or blister copper as it is known, is obtained from the reaction of molten copper(II) sulfide in the hot oxygen atmosphere of a smelter. CuS(1) + O2(g) Cu(l) + SO2(g). This impure copper is
fashioned into anodes that are incorporated in a very large electrolytic cell containing an aqueous copper(II) sulfate electrolyte and thin cathodes of pure copper. The picture below shows the cathodes being removed from several such cells.
The cell chemistry is as follows:
Cathode : Cu2+(aq) + 2e– Cu(s, pure) Anode : Cu(s,blister) Cu2+(aq) + 2e–
Net : Cu(s,blister) + Cu2+(aq) Cu2+(aq) + Cu(s, pure)
4. Chromium Plating: Chromium metal is much more resistant to oxidation by the environment than iron or other metals. The thin oxide coat that readily forms on chromium is resistant to further oxidation. We say that the Cr(s) is passivated. Oxide coats on aluminum and zinc passivate zinc. Chromium is plated onto iron objects that have already been nickel plated (the bumper in the diagram) acting as cathodes in the electrolysis of a solution of chromium(IV) oxide in sulfuric acid (diagram) as it appeared in a past Chemistry 30 diploma examination.
1 = anode, 2 = error, 3= current flow, 4 = cathode, 5, = electrolyte cation, 6 = electrolyte anion, 7 = anion flow, 8 = cation flow
The chemistry of this cell is complex. The overall reaction at the cathode can be summarized as:
Cathode : CrO3(aq) + 6H+(aq) + 6e– Cr(s) + 2H2(1)
By manipulating the length of time the cell operates, and varying the current magnitude, it is possible to obtain the desired thickness of chromium plating.
Silver electroplating is little different from
Faraday's Laws of Electrolysis
During the operation of an electrochemical cell the mass of reactant consumed, product produced is dependent on several cell and reactant/product parameters:
1. I - the electric current in coulombs of charge per second – 1 A= 1C/s 2. t - the cell operation time in seconds.
3 The molar mass of the product/reactant as appropriate.
4. The number of mole of e– lost/gained to produce/consume a mole of element X. The Faraday, symbol F, is the quantity of electric charge on one mole of electrons, namely 9.65 × 104 C/mol e– (Coulombs). Once the number of moles of electrons passing through a cell is known a simple stoichiometry calculation using the appropriate half reaction will give an answer to the problem. Here are a few worked examples.
Example 1 What mass of copper is reduced when a standard copper-zinc voltaic cell operates with an electric current of 400 mA for exactly one hour?
Cu2+(aq) + 2 e– Cu(s)
400 mC/s 63.55 g/mol
3600 s mCu = ?
-4
-400 1 1 63.55
3600 474
9.65 10 2 1
mC mol e molCu g
mCu s mg of Cu
s C mol e mol Cu
The mass of copper produced at the cathode will be 474 mg.
A sound understanding of cell stoichiometry is all that is needed to tackle the next problem.
Example 2 Estimate how much time it takes to produce 28 kg of copper electroplate in an electrowinning cell supplied with an electric current of 100 kA?
- 4
2
-2 9.65 10
28 8.5 10 ( 14 min)
63.55 1 100
mol Cu mol e C s
t kg s ca
g mol Cu mol e kC
Example 3 Estimate the electric current required to produce 1.000 tonne of pure aluminum in a Hall-Héroult cell every hour by the industrial electrolysis of an aqueous copper(U) sulfate. (1 tonne = 106 g)
- 4
-3 9.65 10 1
1000 724 /
26.98 1 3600
mol Al mol e C
I kg kC s
g mol Al mol e s
The predicted current to produce 1000 kg/h of aluminum is 724 kC/s, 724 kA.
Example 4 What mass of element appears at one of the electrodes of a standard silver – zinc voltaic cell if 10.0 g of element is consumed at the other electrode? Cathode/Reduction : 2 Ag+(aq) + 2 e– 2 Ag(s)
Anode/Oxidation : Zn(s) Zn2+(aq) + 2 e–
Net Equation: 2 Ag+(aq) + Zn(s) 2 Ag(s) + Zn2+(aq)
It is now obvious that the zinc is consumed at the anode (corroded) and the silver is produced at the cathode (electroplated).
2 107.87
10.0 33.0 ( )
65.41 1
mol Zn mol Ag g
mAg g g of Ag s
g mol Zn mol Ag
Cell Stoichiometry &. Faraday's Laws
1. What mass of zinc is consumed at the anode of a Daniel cell, operating with a current of 500 mA, over a period of 24 hours?
2. Estimate how much time it takes to produce 1000 kg of aluminium in a Hall-Héroult cell supplied with an electric current of 125 kA?
3. A Downs cell, for producing sodium metal from molten sodium chloride, operates with an electric current of 30 000 Amperes at a temperature of 600°C. How many seconds does it take such a cell to produce exactly one mole of sodium?
4. In a Chemistry 30 demonstration a current of 6.00 amps is passed through an aqueous potassium iodide solution using carbon electrodes. What mass of product is expected at the anode after 30 minutes?
6. At SATP, what volume of hydrogen gas is collected above the cathode during the electrolysis of aqueous potassium sulfate with a current of 5.00 A for 30 minutes?
7. If 10.0 g of element are produced at one electrode of a copper zinc standard cell what mass of element is consumed at the other electrode?
8. At SATP, what volume of chlorine gas is collected above the anode during the electrolysis of aqueous sodium chloride with a current of 5.00 A for 15 minutes?
9. If 1.08 g of element are produced in a silver copper standard cell operating with a current of 200 mA for how long was the cell allowed to operate?
10. The appropriate potential and current is supplied to several electrolytic cells
Cells and Batteries
There are three general types of cells used to provide power to portable devices
Primary cells: single use only. Attempting to recharge these can lead to fires/explosions. The products/by-products of primary cell function are either too delicate to survive recharging or are unsuitable for recharging.
Secondary cells: Can be recharged multiple times once run down to 0.0 V, (nickel cadmium cell) or continuously (car battery). The
recharging process reverse the redox reaction that occurred when the cell produced electricity/ discharged.
Fuel cells: provide power if the relevant fuels are supplied to the anode and cathode( hydrogen oxygen fuel cell)
The chemistry occurring at either electrode within each of the above type of cells are essentially peculiar to each cell. You are expected to be able to recognize the anode and the cathode half reactions given the appropriate equations and electrical potentials, and ultimately construct the net cell reaction equation.
A battery comprises two or more cells arranged in series. The total potential of a battery is the sum of the potentials of all the cells that it contains.
For example a common 9 V battery, shown, contains six lozenge shaped 1.5 V cells stacked one atop the other.
Primary cells
Zinc-carbon cell/standard carbon cell(Enet = 1.5 V )
Zinc-carbon chemistry is used in all inexpensive AA, C and D dry-cell batteries. The electrodes are zinc and carbon, with an acidic paste between them that serves as the electrolyte.
Cathode : 2 MnO2(s) + 2 NH4+(aq) + 2 e– 2 Mn2O3(s) + 2 NH3(aq) + H2O(l)
Anode : Zn(s) Zn2+ (aq) + 2e–
The cathode mix is a paste of NH4Cl(aq),
MnO2(s) and C(s) for a
dry carbon zinc cell.
Alkaline battery (Enet = 1.5 V) - Used in common Duracell and Energizer batteries, the electrodes are zinc and carbon/manganese-oxide, with an alkaline electrolyte (potassium hydroxide)
Cathode : 2 MnO2(s) + H2O(l) + 2 e– 2 Mn2O3(s) + 2 OH–(aq)
Anode : Zn(s) + 2 OH–(aq) ZnO(s) + 2e–
Secondary cells:
Lead-acid battery/accumulator (Enet= 2.05 V) Used in automobiles, the electrodes are made of mossy lead and lead(IV) oxide with a sulfuric acidic electrolyte.
The cell chemistry is:
Cathode : PbO2(s) + SO42–(aq) + 4 H+(aq) + 2 e– PbSO4(s) + 2 H2O(1)
Anode. : Pb(s) + SO42–(aq) 2 PbSO4(s) + 2 e–
The nickel cadmium rechargeable cell is popular for applications that require mobile power sources *saws, drills, etc). A 12 V car battery contains six cells each providing 2.00 V or thereabouts.
Cathode :
2NiO(OH)(s)+2H2O(l) +2e–2Ni(OH)2(s) +2OH–(aq)
Anode :
A Fuel Cell
The hydrogen oxygen fuel cell(Enet= 0.6-0.7 V) (PEM = proton exchange membrane or polymer electrolyte membrane) pioneered by Ballard industries of BC is described by the following diagram.
The chemistry is as follows
Cathode : O2(g) + 4 H+(pem) + 4e– 2 H2O(l)
Anode : 2 H2(g) 4 H+(pem) + 4e–
Notes : H+(pem) refers to hydrogen ions incorporated/dissolved in the PEM electrolyte. The PEM electrolyte is a highly modified version of the polymer component of Gore-tex® fabric. The oxygen is 20% of the air supplied to the cathode. Platinum metal particles catalyze the half
reactions that occur at each electrodes – hence the expense of these cells. The hydrogen oxygen alkaline fuel cell, AFC,(Enet= ca 0.7 V) designed for NASA’s space shuttle uses a hot 30–45% KOH(aq) as the electrolyte.
The chemistry is as follows
Cathode : O2(g) + 2 H2O(l) + 4 e– 4 OH–(aq)
Anode : 2 H2(g) + 4 OH–(aq) 4 H2O(l) + 4 e–
Rust and Corrosion.
Rust is the product of the corrosion of iron. If one part of a continuous piece of iron is exposed to oxygen and moisture (the oxidizing agent) it becomes cathodic, while the piece of iron (the reducing agent) distant from the oxygen and moisture becomes anodic. When iron is like an anode it can corrode/oxidize as shown below:
(From: http://www.corrosionist.com/corrosion-drop-water.jpg, accessed 10/3/10) The hydroxide anions produced at the cathode area migrate towards the anode area and react with the cations that migrate from the anode area. Rust is actually a complex variety of hydrated iron(III) oxides and hydroxides produced by further oxidation of the Fe(OH)2(s) formed above. The reactions that initiate the rusting process are:
Cathodic area : O2(g) + 2 H2O(l) + 4 e– 4 OH–(aq)
Anodic area : 2 Fe(s) 2 Fe2+(aq) + 4 e–
Net Reaction : O2(g) + 2 H2O(l) + 2 Fe(s) 2 Fe2+(aq) + 4 OH–(aq) (or 2 Fe(OH)2(s)
Familiar Rust/Corrosion Images
Bridge bolts (opposite) The rust on the bolt surface
indicates that the underlying iron has corroded/oxidized. more importantly the bolt underneath is worn away somewhat.
A burst rust blister (opposite) on a steel plate. Rust blisters form because rust, being less dense, takes up more space than iron. The paint blisters out and then cracks open. In this way more rust can form.
(From: http://en.wikipedia.org/wiki/Rust, accessed 10/3/10)
1. Paint, grease, oil, and plastic coatings can protect iron by preventing access of oxygen and water. This is why bicycle chains are greased and car bodies are covered in primer and several layers of paint.
2. Metal electroplates accomplish the same effect as the coatings in 1. However, when Sn(s) or Ni(s) coatings are scratched the iron remains the SRA and it is oxidized faster than if unprotected. If Zn(s), Cr(s) or Al(s) coatings on iron are scratched the metal coating acts as the SRA and the iron is still protected – see below.
Iron sheeting coated in a thin applied coating or electroplate of zinc is said to be
galvanized. Iron coated with an electroplate of nickel covered in turn by an
3. Cathodic protection with Sacrificial Anodes prevents corrosion by making all the iron in a structure cathodic by providing a source of electrons. This involves
attaching iron structures to pieces of Zn(s), Al(s), or Mg(s) anodes. These metals, are (i) stronger reducing agents than Fe(s), (ii) oxidize instead of the iron, and (iii) as they oxidize they release electrons to the iron making it a cathode. Metals other than Zn(s), Al(s), and Mg(s) can actually cause iron to corrode and rust! (Copper is very notorious in this regard – it nearly caused the collapse of the Statue of Liberty.)
The cutaway diagram opposite illustrates the basic principle of cathodic protection by sacrificial anode. The sacrificial anode, a block of Zn(s), Al(s), or Mg(s), is connected directly to the iron. The sacrificial anode oxidizes instead of the iron and loses electrons directly to the iron pipe – which now becomes a cathode. Without the sacrificial anode it would probably have been an anode and subject to corrosion.
As shown opposite the sacrificial anode, Al(s), can be connected to the iron structure by an electrical wire (anode connection). As the aluminum corrodes electrons are shunted into the iron structure. The size of the anode is calculated to deliver the current necessary to keep the iron structure/ cathode free of corrosion for its projected life-span. It may be replaced at some time in the future.
Zinc sacrificial anodes, the blobs attached to the reinforcing bars, provide localized corrosion protection in reinforced concrete structures. Their size and location are crafted to deliver sufficient current for the projected life of the rebar within the concrete.
From:
http://www.ssihawaii.com/xp.htm,
Magnesium sacrificial anode being directly attached/bolted to a boat hull. The wires attach to the steel hull elsewhere on the hull. Blobs of zinc can also be welded
directly onto the hull in widely spread locations and have the same protective effect.
4. Cathodic protection by Impressed Current involves directing an electric current, form a power source, into an iron structure and causing it to be a cathode. The size of the current is adjusted to be just sufficient to prevent iron oxidation.
Use the following information to answer the next question
A student prepared the following experimental design to investigate a common reduction-oxidation phenomenon.
1. Describe the most likely outcome of this experiment. Your answer should also include
• balanced half reaction equations, and electrical potentials for the reaction of the appropriate metal in each of the beakers
• an explanation for the reaction that occurs
Use the following information to answer the next question In the 18th century, it was common for wooden ships to have their hulls sheathed in copper to protect them against attack by barnacles which could slow down the ship. Iron bolts were used to fasten the copper sheeting to the hull (see below).
.
In countless instances, corrosion caused the heavy copper bottom to fall off, and the ships, suddenly top-heavy, promptly sank.
2. Use what you know of corrosion chemistry to explain the chemistry behind these sinkings.
Your answer should include
• the identity of the metal that corroded • balanced half and net reaction equations
Cell Chemistry Review
1. Provide a complete cell description for each of the following cells. Your response must include a labeled diagram, net cell equation, Eºnet, and all evidence of reaction.
a. standard nitric acid - zinc cell
b. standard copper – acidified permanganate cell
c. electrolysis of aqueous sodium chloride
2. Write and label the cathode, anode and net reaction equations for the following cell. Cu(s) | Cu2+(aq) || Cl2(g), Cl–(aq) | Pt(s)
3. The electrochemical cell represented by Nb(s) | Nb3+(aq) || Cu2+(aq) | Cu(s) has a standard cell potential of +2.54 V. Copper found to be the positive electrode, the cathode. Use this information to determine the electrical potential for niobium-niobium(III) redox couple.
4. Compare voltaic and the electrolytic cells. In what ways are they similar? How are they different? (A table would be good here.)
5. If the reference electrode is changed to the standard copper half cell, determine a reduction potential for the standard iodine half cell.
7. In a Downs’ cell, a current of 12.0 A is applied for 30 minutes. What mass of sodium will be produced at the cathode during this electrolysis?
8. For what length of time must a 1.00 kA current be applied to form 52.0 kg of aluminum in a Hall-Héroult cell?
9. Write out the half and net reaction equations for the corrosion of unprotected iron in the environment.
11. Complete the Prediction section of the following investigation report. Draw and label a cell diagram as part of your answer.
Problem: What is the complete cell description (half-reactions, net reaction, cell
potential, electrodes, electrolytes, cathode and anode with polarities, and the electron and ion flow) of the maximum-voltage cell assembled from the materials available?
Materials metal strips of iron, cobalt, tin, lead, and nickel. Several 250 mL beakers.1.0 mol/L solutions of the metal ion nitrates. Connecting wires, U-tube salt bridge,
voltmeter.
