chapter 19...salt bridge -- a device used to maintain electrical neutrality in a galvanic cell. this...
TRANSCRIPT
Electrochemistry
Chapter 19
Electrochemistry—the study of the interchange of chemical and
electrical energy
OIL RIG – oxidation is loss, reduction is gain (of electrons)
Oxidation – the loss of electrons, increase in charge
Reduction – the gain of electrons, reduction of charge
Oxidation number – the assigned charge on an atom
Oxidizing agent (OA) – the species that is reduced and thus
causes oxidation
Reducing agent (RA) – the species that is oxidized and thus
causes reduction
Vocabulary
ELECTROCHEMISTRY INVOLVES
TWO MAIN TYPES OF PROCESSES:
A. Galvanic (voltaic) cells – which are spontaneous
chemical reactions (battery)
B. Electrolytic cells – which are non-spontaneous
and require external e− source (DC power source)
C. BOTH of these fit into the category entitled
Electrochemical cells
2Mg (s) + O2 (g) 2MgO (s)
2Mg 2Mg2+ + 4e-
O2 + 4e- 2O2-
Oxidation half-reaction (lose e-)
Reduction half-reaction (gain e-)
19.1
Electrochemical processes are oxidation-reduction reactions
in which:
• the energy released by a spontaneous reaction is
converted to electricity or
• electrical energy is used to cause a nonspontaneous
reaction to occur
0 0 2+ 2-
Oxidation number
The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred.
1. Free elements (uncombined state) have an oxidation
number of zero.
Na, Be, K, Pb, H2, O2, P4 = 0
2. In monatomic ions, the oxidation number is equal to
the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
3. The oxidation number of oxygen is usually –2. In H2O2
and O22- it is –1.
4.4
4. The oxidation number of hydrogen is +1 except when
it is bonded to metals in binary compounds. In these
cases, its oxidation number is –1.
6. The sum of the oxidation numbers of all the atoms in a
molecule or ion is equal to the charge on the
molecule or ion.
5. Group IA metals are +1, IIA metals are +2 and fluorine is
always –1.
HCO3-
O = -2 H = +1
3x(-2) + 1 + ? = -1
C = +4
Oxidation numbers of all
the atoms in HCO3- ?
4.4
Balancing Redox Equations
19.1
1. Write the unbalanced equation for the reaction ion ionic form.
The oxidation of Fe2+ to Fe3+ by Cr2O72- in acid solution?
Fe2+ + Cr2O72- Fe3+ + Cr3+
2. Separate the equation into two half-reactions.
Oxidation:
Cr2O72- Cr3+
+6 +3
Reduction:
Fe2+ Fe3++2 +3
3. Balance the atoms other than O and H in each half-reaction.
Cr2O72- 2Cr3+
Balancing Redox Equations
4. For reactions in acid, add H2O to balance O atoms and H+ to
balance H atoms.
Cr2O72- 2Cr3+ + 7H2O
14H+ + Cr2O72- 2Cr3+ + 7H2O
5. Add electrons to one side of each half-reaction to balance the
charges on the half-reaction.
Fe2+ Fe3+ + 1e-
6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O
6. If necessary, equalize the number of electrons in the two half-
reactions by multiplying the half-reactions by appropriate
coefficients.
6Fe2+ 6Fe3+ + 6e-
6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O
19.1
Balancing Redox Equations
7. Add the two half-reactions together and balance the final
equation by inspection. The number of electrons on both
sides must cancel.
6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O
6Fe2+ 6Fe3+ + 6e-Oxidation:
Reduction:
14H+ + Cr2O72- + 6Fe2+ 6Fe3+ + 2Cr3+ + 7H2O
8. Verify that the number of atoms and the charges are balanced.
14x1 – 2 + 6x2 = 24 = 6x3 + 2x3
19.1
9. For reactions in basic solutions, add OH- to both sides of the
equation for every H+ that appears in the final equation.
Anode--the electrode where oxidation occurs. After a period of time, the
anode may appear to become smaller as it falls into solution.
Cathode-- the anode where reduction occurs. After a period of time it may
appear larger, due to ions from solution plating onto it.
Inert electrodes -- used when a gas is involved OR ion to ion involved such
as Fe3+ being reduced to Fe2+ rather than Fe0. Made of Pt or graphite.
Salt bridge -- a device used to maintain electrical neutrality in a galvanic cell.
This may be filled with agar which contains a neutral salt or it may be
replaced with a porous cup.
Electron flow -- always from anode to cathode. (through the wire)
Standard cell notation (line notation) - anode/solution// cathode solution/
cathode Ex. Zn/Zn2+ (1.0 M) // Cu2+ (1.0 M) / Cu
Voltmeter - measures the cell potential (emf) . Usually is measured in volts.
GALVANIC CELLS
Parts of the voltaic or galvanic cell:
MnO4- + Fe2+ → Mn2+ + Fe3+ [acidic]
Why separate beakers?
Why connected by wire?
What do we need?
-Salt Bridge
-Cell potential -Volts (V) J/Coulomb)
-Volt meter
Galvanic Cells
19.2
spontaneous
redox reaction
anode
oxidation
cathode
reduction
Galvanic Cells
19.2
The difference in electrical
potential between the anode
and cathode is called:
• cell voltage
• electromotive force (emf)
• cell potential
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
Standard Reduction Potentials
19.3
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
2e- + 2H+ (1 M) H2 (1 atm)
Zn (s) Zn2+ (1 M) + 2e-Anode (oxidation):
Cathode (reduction):
Zn (s) + 2H+ (1 M) Zn2+ + H2 (1 atm)
Standard Reduction Potentials
19.3
Standard reduction potential (E0) is the voltage associated
with a reduction reaction at an electrode when all solutes
are 1 M and all gases are at 1 atm.
E0 = 0 V
Standard hydrogen electrode (SHE)
2e- + 2H+ (1 M) H2 (1 atm)
Reduction Reaction
19.3
E0 = 0.76 Vcell
Standard emf (E0 )cell
0.76 V = 0 - EZn /Zn0
2+
EZn /Zn = -0.76 V02+
Zn2+ (1 M) + 2e- Zn E0 = -0.76 V
E0 = EH /H - EZn /Zncell0 0
+ 2+2
Standard Reduction Potentials
E0 = Ecathode - Eanodecell0 0
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
STANDARD REDUCTION POTENTIALS• Each half-reaction has a cell potential
• Each potential is measured against a standard which is the
standard hydrogen electrode [consists of a piece of inert
Platinum that is bathed by hydrogen gas at 1 atm]. The
hydrogen electrode is assigned a value of ZERO volts.
• Standard conditions—1 atm for gases, 1.0M for solutions
and 25°C for all (298 K)
• naught, ° -- we use the naught to symbolize standard
Conditions. That means Ecell , Emf , or εcell become Ecello , Emf
o ,
or εcello when measurements are taken at standard conditions.
The diagram to the right
illustrates what really
happens when a Galvanic
cell is constructed from zinc
sulfate and copper(II) sulfate
using the respective metals
as electrodes.
Notice that 1.0 M solutions of each salt
are used
Notice an overall voltage of 1.10 V for
the process
Standard Reduction Potentials
19.3
Pt (s) | H2 (1 atm) | H+ (1 M) || Cu2+ (1 M) | Cu (s)
2e- + Cu2+ (1 M) Cu (s)
H2 (1 atm) 2H+ (1 M) + 2e-Anode (oxidation):
Cathode (reduction):
H2 (1 atm) + Cu2+ (1 M) Cu (s) + 2H+ (1 M)
E0 = Ecathode - Eanodecell0 0
E0 = 0.34 Vcell
Ecell = ECu /Cu – EH /H 2+ +2
0 0 0
0.34 = ECu /Cu - 002+
ECu /Cu = 0.34 V2+0
19.3
• E0 is for the reaction as
written
• The more positive E0 the
greater the tendency for the
substance to be reduced
• The half-cell reactions are
reversible
• The sign of E0 changes
when the reaction is
reversed
• Changing the stoichiometric
coefficients of a half-cell
reaction does not change
the value of E0
Elements that have the most positive reduction
potentials are easily reduced (in general, non-metals)
Elements that have the least positive reduction potentials
are easily oxidized (in general, metals)
The table can also be used to tell the strength of various
oxidizing and reducing agents.
Can also be used as an activity series. Metals having
less positive reduction potentials are more active and will
replace metals with more positive potentials.
HOW CAN WE DETERMINE WHICH SUBSTANCE IS
BEING REDUCED AND WHICH IS BEING OXIDIZED??
The MORE POSITIVE reduction potential gets to be
reduced - IF you are trying to set up a cell that can act as
a battery.
How to Calculate “Standard Cell Potential”, Symbolized
by E°cell OR Emf° OR εcell° (terms are interchangable)
1. Decide which element is oxidized or reduced using the table of
reduction potentials.
Remember: THE MORE POSITIVE REDUCTION POTENITAL
GETS TO BE REDUCED.
2. Write both equations AS IS from the chart with their voltages.
3. Reverse the equation that will be oxidized and change the sign of
the voltage - This is now: E°oxidation
4. Balance the two half reactions **do not multiply voltage values**
5. Add the two half reactions and the voltages together.
6. E°cell = E°oxidation + E°reduction (where ° means standard conditions: 1atm, 1M, 25°C)
▪AN OX – oxidation occurs at the anode (may show mass decrease)
▪RED CAT – reduction occurs at the cathode (may show mass increase)
▪FAT CAT – The electrons in a voltaic or galvanic cell ALWAYS flow
“From the Anode To the CAThode”
▪Ca+hode – the cathode is + in galvanic (voltaic) cells
▪EPA--in an electrolytic cell, there is a positive (+) anode.
▪Salt Bridge – bridge between cells whose purpose is to provide ions
to balance the charge. Usually made of a salt filled agar (KNO3) or a
porous cup.
Neumonic Devices - Handy when constructing a
“spontaneous cell” --one that can act as a battery:
What is the standard emf of an electrochemical cell made
of a Cd electrode in a 1.0 M Cd(NO3)2 solution and a Cr
electrode in a 1.0 M Cr(NO3)3 solution?
Cd2+ (aq) + 2e- Cd (s) E0 = -0.40 V
Cr3+ (aq) + 3e- Cr (s) E0 = -0.74 V
Cd is the stronger oxidizer
Cd will oxidize Cr
2e- + Cd2+ (1 M) Cd (s)
Cr (s) Cr3+ (1 M) + 3e-Anode (oxidation):
Cathode (reduction):
2Cr (s) + 3Cd2+ (1 M) 3Cd (s) + 2Cr3+ (1 M)
x 2
x 3
E0 = Ecathode - Eanodecell0 0
E0 = -0.40 – (-0.74) cell
E0 = 0.34 V cell
19.3
Sample Problem:
a) Calculate the cell voltage for the following reaction.
b) Draw a diagram of the galvanic cell for the reaction and
label completely.
Fe3+ (aq) + Cu(s) → Cu2+ (aq) + Fe2+ (aq)
Exercise 2
a) Calculate the cell voltage for the galvanic cell that would
utilize silver metal and involve iron(II) ion and iron(III) ion.
b) Draw a diagram of the galvanic cell for the reaction and
label completely. Given: Fe3+ + e-
Fe2+ Eo = 0.77 V
Ag+ + e- Ag(s) Eo = 0.80 V
E°cell= 0.03 V
Reaction Quotient• Q: equal to equilibrium expression but does not have to
be at equilibrium
• used to tell if a reaction is at equilibrium or not
• relationship between Q and K tells which way the
reaction will shift
– Q=K: at equilibrium, no shift
– Q > K: too large, forms reactants, shift to left
– Q < K: too small, forms products, shift to right
A B
13.1
rate = -D[A]
Dt
rate = D[B]
Dt
time
Law of Mass Action
• created in 1864 by Guldberg and Waage
(Norweigen)
• For a reaction: jA + kB ⇄ lC + mD
– equilibrium constant: K
Law of Mass Action
• because
– Ratef = kf[A]j[B]k
– Rater = kr[C]l[D]m
– so if Ratef = Rater
– kf[A]j[B]k = kr[C]l[D]m
19.4
Spontaneity of Redox Reactions
DG = -nFEcell
DG0 = -nFEcell0
n = number of moles of electrons in reaction
F = 96,500J
V • mol = 96,500 C/mol
DG0 = -RT ln K = -nFEcell0
Ecell0 =
RT
nFln K
(8.314 J/K•mol)(298 K)
n (96,500 J/V•mol)ln K=
=0.0257 V
nln KEcell
0
=0.0592 V
nlog KEcell
0
Spontaneity of Redox Reactions
19.4
DG0 = -RT ln K = -nFEcell0
2e- + Fe2+ Fe
2Ag 2Ag+ + 2e-Oxidation:
Reduction:
What is the equilibrium constant for the following reaction
at 250C? Fe2+ (aq) + 2Ag (s) Fe (s) + 2Ag+ (aq)
=0.0257 V
nln KEcell
0
19.4
E0 = -0.44 – (0.80)
E0 = -1.24 V
0.0257 V
x nE0cell
expK =
n = 2
0.0257 V
x 2-1.24 V= exp
K = 1.23 x 10-42
E0 = EFe /Fe – EAg /Ag0 0
2+ +
The Effect of Concentration on Cell Emf
DG = DG0 + RT ln Q DG = -nFE DG0 = -nFE 0
-nFE = -nFE0 + RT ln Q
E = E0 - ln QRT
nF
Nernst equation
At 298
19.5
-0.0257 V
nln QE0E = -
0.0592 Vn
log QE0E =
CELL POTENTIAL, ELECTRICAL WORK & FREE ENERGY
Combining thermodynamics and electrochemistry, and a bit of
physics:
• The work that can be accomplished when electrons are
transferred through a wire depends on the“push” or emf which
is defined in terms of a potential difference [in volts] between
two points in the circuit.
(V = J/C)
Thus “one joule of work” is produced [or required] when one
coulomb of charge is transferred between two points in the
circuit that differ by a potential of one volt- IF work flows OUT it is assigned a MINUS sign
When a cell produces a current, the cell potential is positive
and the current can be used to do work THEREFORE “ε”and “work” have opposite signs!
A “faraday” (F) - the charge on one MOLE of electrons =
96,485 coulombs (AP Exam uses 96,500) And: q = # moles of
electrons × F
For a process carried out at constant temperature and
pressure, wmax [neglecting the very small amount of energy that is lost as friction or heat] is
equal to ΔG, therefore….
ΔGo = −nFEo
G = Gibb’s free energy
n = number of moles of electrons
F = Faraday constant 9.6485309 × 104 J/V • mol
So it follows that:
−Eo implies nonspontaneous.
+Eo implies spontaneous (would be a good battery!)
.
Will the following reaction occur spontaneously at 250C if
[Fe2+] = 0.60 M and [Cd2+] = 0.010 M?
Fe2+ (aq) + Cd (s) Fe (s) + Cd2+ (aq)
2e- + Fe2+ Fe
Cd Cd2+ + 2e-Oxidation:
Reduction:n = 2
E0 = -0.44 – (-0.40)
E0 = -0.04 V
E0 = EFe /Fe – ECd /Cd0 0
2+ 2+
-0.0257 V
nln QE0E =
-0.0257 V
2ln -0.04 VE =
0.010
0.60
E = 0.013
E > 0 Spontaneous
19.5
Bell Work - 9-27-2013 (Exercise 3)Using the table of standard reduction potentials, calculate ΔG° for the
reaction
Cu2+(aq) + Fe(s) → Cu(s) + Fe2+(aq)
Exercise 4Using the table of standard reduction potentials, predict whether 1 M
HNO3 will dissolve gold metal to form a 1 M Au3+ solution.
DEPENDENCE OF CELL POTENTIAL ON CONCENTRATION
Voltaic cells at NON-standard conditions: LeChatlier’s
principle can be applied. An increase in the concentration
of a reactant will favor the forward reaction and the cell
potential will increase. (The converse is also true!)
Exercise 5 - For the cell reaction:
2Al(s) + 3Mn2+(aq) → 2Al3+(aq) + 3Mn(s) E°cell = n/a
predict whether Ecell is larger or smaller than E°cell for the
following cases.
a. [Al3+ ] = 2.0 M, [Mn2+ ] = 1.0 M
b. [Al3+ ] = 1.0 M, [Mn2+] = 3.0 M
Bell Work Oct 19, 2013
For the cell reaction:Al(s) + Mn2+(aq) → Al3+ (aq) + Mn(s)
1) Balance the equation
2) Calc E°cell
3) Predict whether Ecell is larger or smaller than E°cell for the following
cases.
a. [Al3+ ] = 2.0 M, [Mn2+ ] = 1.0 M
b. [Al3+ ] = 1.0 M, [Mn2+] = 3.0 M
4) Calc. the Ecell for each of the above solutions.
5) Calc ΔG° for the “std conditions rxn” (#2) and ΔG for each of the
different concentrations from #3
For a more quantitative approach….. When cell is not at
standard conditions, use Nernst Equation:
R = Gas constant 8.315 J/K• mol
F = Faraday constant
Q = reaction quotient [products]coefficient/[reactants]coefficient
E = Energy produced by reaction
T = Temperature in Kelvins
n = # of electrons exchanged in BALANCED redox equation
As E declines with reactants converting to products, E
eventually reaches zero.
Zero potential means reaction is at equilibrium [dead
battery]. Also, Q = K AND ΔG = 0 as well.
CONCENTRATION CELLSWe can construct a cell where both compartments contain the
same components BUT at different concentrations.
Notice the difference in the concentrations pictured at
left. Because the right compartment contains 1.0 M
Ag+ and the left compartment contains 0.10 M Ag+,
there will be a driving force to transfer electrons from
left to right. Silver will be deposited on the right
electrode, thus lowering the concentration of Ag+ in
the right compartment. In the left compartment the
silver electrode dissolves [producing Ag+ ions] to raise
the concentration of Ag+ in solution.
Exercise 6
Determine the direction of electron flow and designate the
anode and cathode for the cell represented here.
Exercise 7
Determine Eocell and Ecell based on the following half-reactions:
VO2+ + 2H+ + e- → VO2+ + H2O E°= 1.00 V
Zn2+ + 2e- → Zn E° = -0.76V
Where:
T = 25°C
[VO2+] = 2.0 M
[H+] = 0.50 M
[VO2+] = 1.0 x 10-2 M
[Zn2+] = 1.0 x 10-1 M
Eocell = 1.76 V
Ecell = 1.89 V
SUMMARY OF GIBB’S FREE ENERGY AND CELLS−Eo implies NON-spontaneous.
+Eo implies spontaneous (would be a good battery!)
E = 0, equilibrium reached (dead battery)
The larger the voltage, the more spontaneous the reaction
ΔG will be negative in spontaneous reactions
K > 1 are favored (product favored)
Two important equations:
ΔG = −nFE° -- [“minus nunfe”]
ΔG = −RTlnK -- [“ratlink”]
G = Gibbs free energy [Reaction is spontaneous if ΔG is negative]
n = number of moles of electrons.
F = Faraday const. 9.6485309×104 J/V (1 mol of electrons carries 96,500C )
E = cell potential
R = 8.31 J/mol•K
T = Kelvin temperature
K = equilibrium constant -- [products]coeff /[reactants]coeff
**Favored conditions: Ecell > 0 ΔG < 0 K > 1**
Exercise 8
For the oxidation-reduction reaction:
S4O62- (aq) + Cr2+ (aq) → Cr3+ (aq) + S2O3
2-(aq)
The appropriate half-reactions are:
S4O62- + 2e- → 2 S2O3
2- E° = 0.17V
Cr3+ + e- → Cr2+ E° = -0.50V
Balance the redox reaction, and calculate E° and K (at 25°C).
E° = 0.67 V
K = 1022.6 = 4 × 1022
Batteries
19.6
Leclanché cell
Dry cell
Zn (s) Zn2+ (aq) + 2e-Anode:
Cathode: 2NH4 (aq) + 2MnO2 (s) + 2e- Mn2O3 (s) + 2NH3 (aq) + H2O (l)+
Zn (s) + 2NH4 (aq) + 2MnO2 (s) Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s)
Batteries
Zn(Hg) + 2OH- (aq) ZnO (s) + H2O (l) + 2e-Anode:
Cathode: HgO (s) + H2O (l) + 2e- Hg (l) + 2OH- (aq)
Zn(Hg) + HgO (s) ZnO (s) + Hg (l)
Mercury Battery
19.6
Batteries
19.6
Anode:
Cathode:
Lead storage
battery
PbO2 (s) + 4H+ (aq) + SO2- (aq) + 2e- PbSO4 (s) + 2H2O (l)4
Pb (s) + SO2- (aq) PbSO4 (s) + 2e-4
Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO2- (aq) 2PbSO4 (s) + 2H2O (l)4
Batteries
19.6Solid State Lithium Battery
Batteries
19.6
A fuel cell is an
electrochemical cell
that requires a
continuous supply of
reactants to keep
functioning
Anode:
Cathode: O2 (g) + 2H2O (l) + 4e- 4OH- (aq)
2H2 (g) + 4OH- (aq) 4H2O (l) + 4e-
2H2 (g) + O2 (g) 2H2O (l)
Chemistry In Action: Bacteria Power
CH3COO- + 2O2 + H+ 2CO2 + 2H2O
Corrosion
19.7
Cathodic Protection of an Iron Storage Tank
19.7
Electrolysis
19.8
Electrolysis is the process in which electrical energy is used
to cause a nonspontaneous chemical reaction to occur.
ELECTROLYSIS AND ELECTROLYTIC CELLS
Electrolysis- the use of electricity to bring about chemical
change. Literal translation “split with electricity”
Electrolytic cells [NON spontaneous cells]:
Used to separate ores or plate out metals.
Important differences between a voltaic/galvanic cell and an
electrolytic cell:1) Voltaic cells are spontaneous and electrolytic cells are forced to occur
by using an electron pump or battery or any DC source.
2) A voltaic cell is separated into two half cells to generate electricity; an
electrolytic cell occurs in a single container.
3) A voltaic [or galvanic] cell IS a battery, an electrolytic cell NEEDS a
battery
4) AN OX and RED CAT still apply BUT the polarity of the electrodes is
reversed. The cathode is Negative and the anode is Positive (remember
E.P.A – electrolytic positive anode). Electrons
still flow FATCAT.
5) Usually use inert electrodes
Predicting the Products of Electrolysis:
If NO water is present and you have a pure molten ionic
compound, then:
the cation will be reduced (gain electrons/go down in charge)
the anion will be oxidized (lose electrons/go up in charge)
If water IS present and you have an aqueous solution of
the ionic compound, then:
you’ll need to figure out if the ions are reacting or the water
is reacting. You can look at a reduction potential table to
figure it out but, OR as a rule of thumb:
•No group IA or IIA metal will be reduced in an aqueous
solution
• Water will be reduced instead.
•No polyatomic will be oxidized in an aqueous solution
• Water will be oxidized instead.
Electrolysis of Water
19.8
Calculating the Electrical Energy of ElectrolysisHow much metal could be plated out?
How long would it take to plate out?
Faraday’s Law: The amount of a substance being oxidized
or reduced at each electrode during electrolysis is directly
proportional to the amount of electricity that passes through
the cell.
Use dimensional analysis (unit factor, T-chart) for these
calculations, remembering: # coulombs = It where:
1 Volt = 1 Joule/Coulomb
1 Amp = 1 Coulomb/second (current is measured in amp, but
symbolized by I)
Faraday = 96,500 Coulombs/mole of electrons
Balanced redox equation gives #moles of e−/mole of
substance
Formula weight gives grams/mole
Electrolysis and Mass Changes
charge (C) = current (A) x time (s)
1 mole e- = 96,500 C
19.8
How much Ca will be produced in an electrolytic cell of
molten CaCl2 if a current of 0.452 A is passed through the
cell for 1.5 hours?
Anode:
Cathode: Ca2+ (l) + 2e- Ca (s)
2Cl- (l) Cl2 (g) + 2e-
Ca2+ (l) + 2Cl- (l) Ca (s) + Cl2 (g)
2 mole e- = 1 mole Ca
mol Ca = 0.452C
sx 1.5 hr x 3600
s
hr 96,500 C
1 mol e-
x2 mol e-
1 mol Cax
= 0.0126 mol Ca
= 0.50 g Ca
19.8
Exercise 9
How long must a current of 5.00 A be applied to a solution of
Ag+ to produce 10.5 g silver metal?
= 31.3 min
Exercise 10
An acidic solution contains the ions Ce4+, VO2+, and Fe3+ .
Using the E° values [listed in Table 17.1, Zumdahl] give the order of
oxidizing ability of these species and predict which one will be
reduced at the cathode of an electrolytic cell at the lowest
voltage.
Ce4+ > VO2+ > Fe3+
Applications of electrolytic cells:▪Production of pure forms of elements from mined ores
Aluminum from Hall-Heroult process
▪Separation of sodium and chlorine (Down's cell)
▪Purify copper for wiring
▪Electroplating—applying a thin layer of an expensive metal
to a less expensive one Jewelry --- 14 K gold plated
▪Bumpers on cars --- Chromium plated
▪Charging a battery --- i.e. your car battery when the
alternator functions
Corrosion – The process of returning metals to their
natural state, the ores.
▪Involves oxidation of the metal which causes it to lose its
structural integrity and attractiveness.
▪The main component of steel is iron.
▪20% of the iron and steel produced annually is used to
replace rusted metal!
▪most metals develop a thin oxide coating to protect them,
patina’s, tarnish, rust, etc.