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![Page 1: Atkins & de Paula: Elements of Physical Chemistry: 5e Chapter 9: Chemical Equilibrium: Electrochemistry.](https://reader034.fdocuments.in/reader034/viewer/2022050700/56649e375503460f94b2815e/html5/thumbnails/1.jpg)
Atkins & de Paula: Atkins & de Paula: Elements of Physical Chemistry: Elements of Physical Chemistry:
5e5e
Chapter 9: Chemical Equilibrium: Electrochemistry
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End of chapter 9 assignments
Discussion questions:• 1, 4
Exercises:• 1, 3, 9, 12, 13 (include last 3?)
Use Excel if data needs to be graphed
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Homework assignmentsHomework assignments
• Did you:– Read the chapter?– Work through the example problems?– Connect to the publisher’s website &
access the “Living Graphs”?– Examine the “Checklist of Key Ideas”?– Work assigned end-of-chapter
exercises?
• Review terms and concepts that you should recall from previous courses
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Build Yourself a Table…Build Yourself a Table…
TERM UNITS SYMBOLPotential volts V
Resistance ohms
Current amp I
Siemens -1 (ohm-1) S
Resistivity ohm meter
Conductivity ohm-1 meter-1
Molar conductivity
S m2 mol-1 m
Ionic conductivity
mol/dm3
mS m2 mol-1 + –T
erm
s, U
nits
, & S
ymbo
ls
is an uppercase
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Foundational conceptsFoundational concepts
• What is the most important difference between solutions of electrolytes and solutions of non-electrolytes?
• Long-range (Coulombic) interactions among ions in solutions of electrolytes
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The Debye-Hückel theoryThe Debye-Hückel theory
• Activity, a, is roughly “effective molar concentration”
• 9.1a aJ = JbJ/b b = 1 mol/kg
• 9.1b aJ = JbJ = activity
coefficient– treating b as the numerical value of molality
• If a is known, you can calculate chemical potential: μJ = μJ + RT ln aJ (9.2)
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The mean activity coefficientThe mean activity coefficient
Mean activity coefficient = (+ –)½
For MX, = (+ –)½
• For MpXq, = (+p
–q)1/s s = p+q
• So for Ca3(PO4)2, = (+3
–2)1/5
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Debye-Hückel theoryDebye-Hückel theory
• Fig 9.1 (203)• A depiction of the
“ionic atmosphere” surrounding an ion
• The energy of the central ion is lowered by this ionic atmosphere
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Debye-Hückel theoryDebye-Hückel theory
• Debye-Hückel limiting law:log = –A|z+z–| I ½
is the mean activity coefficient
– I = ionic strength of the solutionI = ½(z+
2 b+ + z–
2 b– ) [b =
molality]
– A is a constant; A = 0.509 for water– z is the charge numbers of the ions
p.203
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The extended Debye-Hückel The extended Debye-Hückel lawlaw
• log = – + C.I
is the mean activity coefficient
– I = ionic strength of the solutionI = ½(z+
2 b+ + z–
2 b– )
– A is a constant; A = 0.509 for water– B & C = empirically determined
constants– z = the charge numbers of the ions
A |z+z–| I ½
1 + B. I ½
p.203
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Debye-Hückel theoryDebye-Hückel theory
• Fig 9.2 (203)• (a) the limiting law
for a 1,1-electrolyte(B & C = 1)
• (b) the extended law for B = 0.5
• (c) the extended law extended further by the addition of the C I term
[in the graph, C=0.2]
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The migration of ionsThe migration of ions
• Ions move• Their rate of motion indicates:
– Size, effect of solvation, the type of motion
• Ion migration can be studied by measuring the electrical resistance in a conductivity cell
• V = IR
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The migration of ionsThe migration of ions
• V = IR• Resistivity () and conductivity ()• And = 1/ and = 1/• Drift velocity, s = uE • Where u (mobility) depends
on a, the radius of the ion and , the viscosity of the solution
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Conductivity cellConductivity cell
• Fig 9.3 (204)• The resistance is
typically compared to that of a solution of known conductivity
• AC is used to avoid decomposition products at the electrodes
Conductivity bridge
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Do you see any trends?Do you see any trends?T
9.1
Ioni
c co
nduc
tivi
ties
, /(
mS
m2/
mol
)*
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T9.
2 Io
nic
mob
ilit
ies
in w
ater
at 2
98 K
, u/
(10-
8 m
2 s-
1 V
-1)
Do you see any trends?Do you see any trends?
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The hydrodynamic radiusThe hydrodynamic radius
• The equation for drift velocity (s) and the equation for mobility (u) together indicate that the smaller the ion, the faster it should move…
• But the Group 1A cations increase in radius and increase in mobility! The hydrodynamic radius can explain this phenomenon.
• Small ions are more extensively hydrated.
s = uE
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Proton conduction through Proton conduction through waterwater
• Fig 9.4 (207) The Grotthus mechanism• The proton leaving on the right side is not
the same as the proton entering on the left side
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Determining the Isoelectric Determining the Isoelectric PointPoint
• Fig 9.5 (207)• Speed of a
macro-molecule vs pH
• Commonly measured on peptides and proteins (why?)
• Cf “isoelectric focusing”
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Types of electrochemical Types of electrochemical rxnsrxns
• Galvanic cell—a spontaneous chemical rxn produces an electric current
• Electrolytic cell—a nonspontaneous chemical rxn is “driven” by an electric current (DC)
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Anatomy of electrochemical cellsAnatomy of electrochemical cells
Fig 9.6 (209) Fig 9.7 (209) The salt bridge overcomes difficulties that the liquid junction introduces into interpreting measurements
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Half-reactionsHalf-reactions
• For the purpose of understanding and study, we separate redox rxns into two half rxns: the oxidation rxn (anode) and the reduction rxn (cathode)
• Oxidation, lose e–, increase in oxid #• Reduction, gain e–, decrease in oxid #• Half rxns are conceptual; the e– is
never really free
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Dir
ect
ion
of
e– flow
in
ele
ctro
chem
ical
cells
Fig
9.8
(21
3)
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Reactions at electrodesReactions at electrodes
• Fig 9.9 (213)• An electrolytic cell• Terms:
– Electrode – Anode– Cathode
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• Fig 9.10 (213) Standard Hydrogen Electrode
• Is this a good illustration of the SHE?
• Want to see a better one?
A gas electrodeA gas electrode
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19.3
E0 = 0 V
2e- + 2H+ (1 M) 2H2 (1 atm)
Reduction Reaction
Standard hydrogen electrode (SHE)
SStandard tandard HHydrogen ydrogen EElectrodelectrode
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19.3
E0 = 0 V
Standard hydrogen electrode (SHE)
Standard Hydrogen ElectrodeStandard Hydrogen Electrode
H2 (1 atm) 2H+ (1 M) + 2e-
Oxidation Reaction
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H2 gas, 1 atm
Pt electrode
SHE acts as cathode SHE acts as anode
Standard Hydrogen ElectrodeStandard Hydrogen Electrode
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Metal-insoluble-salt electrodeMetal-insoluble-salt electrode
• Fig 9.11 (214)• Silver-silver
chloride electrode• Metallic Ag coated
with AgCl in a solution of Cl–
• Q depends on aCl
ion
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Variety of cellsVariety of cells
• Electrolyte concentration cell• Electrode concentration cell• Liquid junction potential
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Redox electrode
• Fig 9.12 (215)• The same element in two non-
zero oxidation states
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The Daniell cellThe Daniell cell
• Fig 9.13 (215)• Zn is the anode• Cu is the
cathode
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The cell reactionThe cell reaction
• Anode on the left; cathode on the right
Cell Diagram
Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)
[Cu2+] = 1 M & [Zn2+] = 1 M
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
anode cathode
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Measuring cell emfMeasuring cell emf
Fig 9.13 (217) Cell emf is measured by balancing the cell against an opposing external potential. When there is no current flow, the opposing external potential equals the cell emf.
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The electromotive forceThe electromotive force
• The maximum non-expansion work (w’max) equals G [T,p=K] (9.12)
• Measure the potential difference (V) and convert it to work to calculate G
rG = –FE (F = 96.485 kC/mol)
• E = – rGF
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The electromotive forceThe electromotive force
rG = –FE
rG = rG + RT ln Q
• E = E – ln Q
• E =
• At 25°C, = 25.693 mV
• E is independent of how the rxn is balanced
RTRTFF
rGrGFF
RTRTFF
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Cells at equilibriumCells at equilibrium
ln K = FERT
• At equilibrium, Q = K and a rxn at equilibrium can do no work, so E = 0
• So when Q = K and E = 0, the Nernst equation
E = E – ln Q , becomes….
RTF
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Cells at equilibriumCells at equilibrium
ln K = FERT
Is simply an electrochemical expression of
rG = – RT ln K
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Cells at equilibriumCells at equilibrium
• If E > 0, then K > 1 and at equilibrium the cell rxn favors products
• If E < 0, then K < 1 and at equilibrium the cell rxn favors reactants
218218
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Standard potentialsStandard potentials
• SHE is arbitrarily assigned E = 0 at all temperatures, and the standard emf of a cell formed from any pair of electrodes is their difference:
• E = Ecathode – Eanode OR
• E = Eright – Eleft
• Ex 9.6: Measure E, then calculate K
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The variation of potential The variation of potential with pHwith pH
If a redox couple involves H3O+, then the potential varies with pH
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Table 9.3 Table 9.3 Standard reduction potentials at 25°C Standard reduction potentials at 25°C
(1)(1) Reduction half-reaction Eo/V
Oxidizing agent Reducing agent
Strongly oxidizing
F2 2 e 2 F 2.87
S2O 82– 2 e 2 SO
42– 2.05
Au e Au 1.69
Pb4 2 e Pb2 1.67
Ce4 e Ce3 1.61
MnO 4– 8 H 5 e Mn2 4 H2O 1.51
Cl2 2 e 2 Cl 1.36
Cr2O 72– 14 H 6 e 2 Cr3 7 H2O 1.33
O2 4 H 4 e 2 H2O 1.23, 0.81 at pH 7
Br2 2 e 2 Br 1.09
Ag e Ag 0.80
Hg22 2 e 2 Hg 0.79
Fe3 e Fe2 0.77
I2 e 2 I 0.54
O2 2 H2O 4 e 4 OH 0.40, 0.81 at pH 7
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Table 9.3 Table 9.3 Standard reduction potentials at 25°C Standard reduction potentials at 25°C
(2)(2) Reduction half-reaction Eo/V
Oxidizing agent Reducing agent
Cu2 2 e Cu 0.34
AgCl e Ag Cl 0.22
2H 2 e H2 0, by definition
Fe3 3 e Fe 0.04
O2 H2O 2 e HO 2– OH 0.08
Pb2 2 e Pb 0.13
Sn2 2 e Sn 0.14
Fe2 2 e Fe 0.44
Zn2 2 e Zn 0.76
2 H2O 2 e H2 2 OH 0.83, 0.42 at pH 7
Al3 3 e Al 1.66
Mg2 2 e Mg 2.36
Na e Na 2.71
Ca2 2 e Ca 2.87
K e K 2.93
Li e Li 3.05
Strongly reducing
For a more extensive table, see the Data section.
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The determination of pHThe determination of pH
• The potential of the SHE is proportional to the pH of the solution
• In practice, the SHE is replaced by a glass electrode (Why?)
• The potential of the glass electrode depends on the pH (linearly)
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A glass electrodeA glass electrode
• Fig 9.15 (222) • The potential of a
glass electrode varies with [H+]
• This gives us a way to measue pKa electrically, since pH = pKa when [acid] = [conjugate base]
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The electrochemical seriesThe electrochemical series
• A couple with a low standard potential has a thermodynamic tendency to reduce a couple with a higher standard potential
• A couple with a high standard potential has a thermodynamic tendency to oxidize a couple with a lower standard potential
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• E0 is for the reaction as written
• The more positive E0 the greater the tendency for the substance to be reduced
• The more negative E0 the greater the tendency for the substance to be oxidized
• Under standard-state condi-tions, any species on the left of a given half-reaction will react spontaneously with a species that appears on the right of any half-reaction located below it in the table (the diagonal rule)
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• The half-cell reactions are reversible
• The sign of E0 changes when the reaction is reversed
• Changing the stoichio-metric coefficients of a half-cell reaction does not change the value of E0
• The SHE acts as a cath-ode with metals below it, and as an anode with metals above it
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The determination of thermodynamic The determination of thermodynamic functionsfunctions
• By measuring std emf of a cell, we can calculate Gibbs energy
• We can use thermodynamic data to calculate other properties (e.g., rS)
rS =
F(E – E’)T – T ’
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Determining thermodynamic Determining thermodynamic functionsfunctions
• Fig 9.16 (223)
• Variation of emf with temperature depends on the standard entropy of the cell rxn
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Key Key IdeasIdeas
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Key Key IdeasIdeas
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Key Key IdeasIdeas
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The EndThe End…of this chapter…”
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Box 9.1 pp207ff
Ion channel
s and pumps