Chemistry 100 – Chapter 20 Electrochemistry. Voltaic Cells.
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Transcript of Chemistry 100 – Chapter 20 Electrochemistry. Voltaic Cells.
![Page 1: Chemistry 100 – Chapter 20 Electrochemistry. Voltaic Cells.](https://reader031.fdocuments.in/reader031/viewer/2022031922/56649e7b5503460f94b7cba8/html5/thumbnails/1.jpg)
Chemistry 100 – Chapter 20
Electrochemistry
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Voltaic Cells
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A Schematic Galvanic Cell
e-
Reducing Agent
e-
e-
Oxidizing Agent
Anode Cathode
Porous Disk
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The Galvanic Cell Defined Galvanic cells – an electrochemical cell
that drives electrons through an external circuit as a result of the spontaneous redox reaction occurring inside.
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The Zn/Cu Galvanic Cell
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Voltaic Cells We expect the Zn electrode to lose mass
and the Cu electrode to gain mass. “Rules” of voltaic cells:
At the anode electrons are products. (Oxidation)
At the cathode electrons are reactants (Reduction)
Electrons flow from the anode to the cathode.
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The Anode and Cathode Galvanic cells - the anode is
negative and the cathode is positive.
Electrons are made to flow through an external circuit. (Rule 3.)
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Cell Potentials (Electromotive Force or EMF Values) Electromotive force (emf) - aka
the cell potential the force required to push electrons
through the external circuit. Ecell is the emf of a cell (old
notation). Now talk about the cell potential!
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Cell Reactions The difference in the RHS and the LHS
reactionCu2+ (aq) + Zn (s) Cu (s) + Zn2+ (aq)
For each half reaction, we can write the reaction quotient (see Chapter 15) as followsCu2+ (aq) + 2 e- Cu (s) Q = 1/ [Cu2+] Zn2+ (aq) + 2 e- Zn (s) Q = 1/ [Zn2+]
Overall Qcell = [Zn2+] / [Cu2+]
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The Cell Potential and G From the reaction Gibbs energy
cellcello
rxnrxn E F QlnRTGG
cellcell
orxnrxn E
F
QlnRT
F
G
F
G
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The Nernst Equation
E - standard cell potential Cell potential under
standard conditions. [Solutes] = 1.000
mole/L T = 298.15 K P = 1.00 atm pressure
cellcell QlnF
RTEE
F
GE
orxn
cell
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Cell Potentials
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Standard Reduction Potentials We cannot measure the potential
of an individual half-cell! We assign a particular cell as being
our reference cell and then assign values to other electrodes on that basis.
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Cell Potentials are Intensive Properties In the previous example, the cell
potential was simply the difference between the standard potential for the Sn4+/Sn2+ reduction and the Fe3+/Fe2+ reduction.
Reason: standard cell potentials are intensive quantities.
F
GE rxn
cell
F
GE
orxn
cell
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[H+] = 1.00
H2 (g)
e-
Pt gauze
The Standard Hydrogen electrode Eo (H+/H2) half-cell = 0.000 V
p{H2(g)} = 1.00 atm
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A Galvanic Cell With Zinc and the Standard Hydrogen Electrode.
Note - [Zn2+]= [H+] = 1.000 M
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The Cell Equation for the Zinc-Standard Hydrogen Electrode.
The cell reaction 2 H+ (aq) + Zn (s) H2 (g) + Zn2+ (aq)
When we measure the potential of this cell
Ecell = ERHS - ELHS but ERHS = E(H+/H2) = 0.000 V Ecell = E(Zn2+/Zn) = -0.763 V
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The Spontaneous Direction of a Cell reaction Examine the magnitude the of the
standard cell potential!
F
GE
orxn
cell
If Eo is positive, the rG is negative! Under standard conditions, the cell will proceed spontaneously in the direction written for the cell reaction.
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The Composition Dependence of the Cell Potential
Nernst equation the nonstandard cell potential (Ecell) will
be a function of the concentrations of the species in the cell reaction.
cellcell QlnF
RTEE
To calculate Ecell, we must know the cell reaction and the value of Qcell.
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Electrochemical Series Look at the following series of reactions Cu2+ (aq) + 2 e- Cu (s) E(Cu2+/Cu) = 0.337 VZn2+ (aq) + 2 e- Zn (s) E(Zn2+/Zn) = -0.763 V Zn has a thermodynamic tendency to
reduce Cu2+ (aq) Pb2+ (aq) + 2 e- Pb (s) E(Pb2+/Pb) = -0.13 VFe2+ (aq) + 2 e- Fe (s) E(Fe2+/Fe) = -0.44 V
Fe has a thermodynamic tendency to reduce Pb2+ (aq)
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Differences in Reduction Potentials
• The larger the difference between Ered values, the larger Ecell.
• In a voltaic (galvanic) cell (spontaneous) Ered(cathode) is more positive than Ered(anode).
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Oxidizing and Reducing Agents The more positive Ered the stronger
the oxidizing agent on the left. The more negative Ered the
stronger the reducing agent on the right.
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Spontaneous Oxidation Processes A species on the higher to the left of
the table of standard reduction potentials will spontaneously oxidize a species that is lower to the right in the table.
Any species on the right will spontaneously reduce anything that is higher to the left in the series.
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Oxidizing and Reducing Agents
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Concentration Cells Two identical half-cells.
RHS AgCl (s) + e- Ag (s) + Cl- (aq, 0.10 M)
LHS AgCl (s) + e- Ag (s) + Cl- (aq, 0.50 M)
Electrolyte concentration cell – the electrodes are identical; they simply differ in the concentration of electrolyte in the half-cells.
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The Nernst equation for the cell
LHS
RHS
cellcell
]Cl[
]Cl[ln
F
RT
QlnF
RTE
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Cells at Equilibrium When the electrochemical cell has
reached equilibrium
cellcellcell KQV 0E
Kcell = the equilibrium constant for the cell reaction.
RT
FE KlnKln
F
RTE cellcell
Knowing the E° value for the cell, we can estimate the equilibrium constant for the cell reaction.
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Equilibrium Constants from Cell Potentials Examine the following cell.
Half-cell reactions. Sn4+ (aq) + 2 e- Sn2+ (aq) E(Sn4+/Sn2+) =
0.15 V Fe3+ (aq) + e- Fe2+ (aq) E (Fe3+/Fe2+) =
0.771 V Cell Reaction
Sn4+ (aq) + 2 Fe3+ (aq) Sn2+ (aq) + 2 Fe2+ (aq)
Ecell = (0.771 - 0.15 V) = 0.62 V
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Lead-Acid Battery A 12 V car battery - 6 cathode/anode
pairs each producing 2 V.Cathode: PbO2 on a metal grid in sulfuric
acid:PbO2(s) + SO4
2-(aq) + 4H+(aq) + 2e- PbSO4(s) + 2H2O(l).
Anode: Pb:Pb(s) + SO4
2-(aq) PbSO4(s) + 2e-
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Lead-Acid Battery The overall electrochemical reaction is PbO2(s) + Pb(s) + 2SO4
2-(aq) + 4H+(aq) 2PbSO4(s) + 2H2O(l)
for whichEcell = ERHS - ELHS
= (+1.685 V) - (-0.356 V)= +2.041 V.
Wood or glass-fiber spacers are used to prevent the electrodes form touching.
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A Picture of a Car Battery
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An Alkaline Battery Anode: Zn cap:
Zn(s) Zn2+(aq) + 2e- Cathode: MnO2, NH4Cl and carbon paste:
2 NH4+(aq) + 2 MnO2(s) + 2e- Mn2O3(s) +
2NH3(aq) + 2H2O(l) Graphite rod in the center - inert cathode. Alkaline battery, NH4Cl is replaced with KOH. Anode: Zn powder mixed in a gel:
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The Alkaline Battery
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Fuel Cells Direct production of electricity from
fuels occurs in a fuel cell. H2-O2 fuel cell was the primary source
of electricity on Apollo moon flights. Cathode: reduction of oxygen:
2 H2O(l) + O2(g) + 4e- 4OH-(aq) Anode:
2H2(g) + 4OH-(aq) 4H2O(l) + 4e-
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Fuel Cells
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Corrosion of Iron Since E(Fe2+/Fe) < E(O2/H2O) iron can
be oxidized by oxygen. Cathode
O2(g) + 4H+(aq) + 4e- 2H2O(l).
Anode Fe(s) Fe2+(aq) + 2e-.
Fe2+ initially formed – further oxidized to Fe3+ which forms rust, Fe2O3• xH2O(s).
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Rusting (Corrosion) of Iron
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Preventing the Corrosion of Iron Corrosion can be prevented by coating
the iron with paint or another metal. Galvanized iron - Fe is coated with Zn. Zn protects the iron (Zn - anode and Fe -
the cathode)
Zn2+(aq) +2e- Zn(s), E(Zn2+/Zn) = -0.76 VFe2+(aq) + 2e- Fe(s), E(Fe2+/Fe) = -0.44 V
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Preventing the Corrosion of Iron
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Preventing the Corrosion of Iron To protect underground pipelines,
a sacrificial anode is added. The water pipe - turned into the
cathode and an active metal is used as the sacrificial anode.
Mg is used as the sacrificial anode:
Mg2+(aq) +2e- Mg(s), E(Mg2+/Mg) = -2.37 VFe2+(aq) + 2e- Fe(s), E(Fe2+/Fe) = -0.44 V
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Corrosion Prevention
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Electrolysis of Aqueous Solutions Nonspontaneous reactions require
an external current in order to force the reaction to proceed.
Electrolysis reactions are non-spontaneous.
In voltaic and electrolytic cells: reduction occurs at the cathode, and oxidation occurs at the anode.
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Voltaic vs.Electrolytic Cells Electrolytic cells – electrons are
forced to flow from the anode to cathode.
In electrolytic cells the anode is positive and the cathode is negative. (In galvanic cells the anode is negative and the cathode is positive.)
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Electrolysis of Aqueous Solutions
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Electrolysis of Molten Salts Decomposition of molten NaCl.
Cathode: 2Na+(l) + 2e- 2Na(l) Anode: 2Cl-(l) Cl2(g) + 2e-.
Industrially, electrolysis is used to produce metals like Al.
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Electrolysis With Active Electrodes
Active electrodes: electrodes that take part in electrolysis.
Example: electrolytic plating.
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Electrolysis With Active Electrodes (cont’d) Consider an active Ni electrode and
another metallic electrode placed in an aqueous solution of NiSO4: Anode: Ni(s) Ni2+(aq) + 2e-
Cathode: Ni2+(aq) + 2e- Ni(s). Ni plates on the inert electrode. Electroplating is important in
protecting objects from corrosion.
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Quantitative Aspects of Electrolysis Consider the reduction of Cu2+ to Cu.
Cu2+(aq) + 2e- Cu(s). 2 mol of electrons 1 mol of Cu. How
much material is obtained?Q = I t
current (I) time (t) of the plating process.
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Gibbs energy – the maximum amount of useful work that can be obtained from a system. nFEw
nFEG
wG
max
max
Gibbs Energy and Work
Note – if wmax is negative, then work is performed by the system and E is positive.
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Electrical Work Eelectrolytic cell – external source
of energy is required to force the reaction to proceed.
External emf must be greater than Ecell.
From physics: work has units watts.1 W = 1 J/s.
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Units of Electrical Work Electric utilities use units of
kilowatt-hours:
J. 106.3
W1
J/s 1
h 1
s 3600h 1 W1000kWh 1
6