Electrochemistry
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Transcript Electrochemistry
Electrochemistry
Chapter 4
4.4 and 4.8
Chapter 19
19.1-19.5 and 19.8
Oxidation-Reduction Reactions
Electron Transfer Reactions
Many redox reactions take place in water.
Half-reactions
OILRIG
Oxidation-Reduction
Reactions
2Mg (s) + O2 (g)
2MgO (s)
2Mg2+ + 4e- Oxidation half-reaction (lose e-)
2Mg
O2 + 4e-
2O2- Reduction half-reaction (gain e-)
2Mg + O2 + 4e2Mg + O2
2Mg2+ + 2O2- + 4e2MgO
Oxidation-Reduction Reactions
Oxidizing Agent- a substance that
accepts electrons from another substance,
causing the other substance to be
oxidized.
Reducing Agent- a substance that
donates electrons to another substance,
causing it to become reduced.
Oxidation-Reduction
Reactions
Zn (s) + CuSO4 (aq)
ZnSO4 (aq) + Cu (s)
Oxidation-Reduction
Reactions
Zn (s) + CuSO4 (aq)
Zn
Zn2+ + 2e-
ZnSO4 (aq) + Cu (s)
Zn is oxidized
Zn is the reducing agent
Cu2+ + 2e-
Cu
Cu2+ is reduced
Cu2+ is the oxidizing agent
Oxidation-Reduction
Reactions
Copper wire reacts with silver nitrate to form silver metal.
What is the oxidizing agent in the reaction?
Cu (s) + 2AgNO3 (aq)
Cu
Cu(NO3)2 (aq) + 2Ag (s)
Cu2+ + 2e-
Ag+ + 1e-
Ag
Ag+ is reduced
Ag+ is the oxidizing agent
Oxidation Number
Oxidation Number- the number of
charges the atom would have in a
molecule (or an ionic compound) if
electrons were transferred completely.
H2(g) + Cl2(g) → 2HCl(g)
0
0
+1 -1
Assigning Oxidation Numbers
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.
Assigning Oxidation Numbers
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.
5. Group IA metals are +1, IIA metals are +2 and fluorine is
always –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.
Assigning Oxidation Numbers
Oxidation numbers of all
the elements in HCO3- ?
HCO3O = -2
H = +1
3x(-2) + 1 + ? = -1
C = +4
Assigning Oxidation Numbers
Redox Titrations
Oxidizing agent/Reducing Agent
Equivalence point reached when reducing
agent is completely oxidized by the
oxidizing agent.
Indicator is needed
Redox Titrations
Electrochemistry
Branch of Chemistry that deals with
interconversion of electrical and chemical
energy.
Importance
Use
of electricity in everyday life
Interconversion of energy
Study of electrochemical processes
Electrochemistry
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
Balancing Redox Equations
Simple reactions easy to balance
Complex reactions
Usually
encountered in laboratory
Can include CrO42-, Cr2O72-, MnO4-, NO3- and
SO42 Follow balancing guidelines
Balancing Redox Reactions
Guidelines
Step 1: Write the unbalanced equation for the reaction in ionic
form.
Step 2: Separate the equation into two half-reactions
Step 3: Balance each half-reaction for number and type of atoms
and charges. (Look at medium)
Step 4: Add the two half-reactions together and balance the final
equation by inspection. The electrons on both sides must cancel.
(Be sure they are equal)
Step 5: Verify that the equation contains the same type and
numbers of atoms and the same charges on both sides of the
equation.
Balancing Redox Equations
The oxidation of Fe2+ to Fe3+ by Cr2O72- in acid solution?
1. Write the unbalanced equation for the reaction ion ionic form.
Fe2+ + Cr2O72-
Fe3+ + Cr3+
2. Separate the equation into two half-reactions.
+2
Fe2+
Oxidation:
+6
+3
Fe3+
+3
Cr3+
Reduction: Cr2O723. Balance the atoms other than O and H in each half-reaction.
Cr2O722Cr3+
Balancing Redox Reactions
4. For reactions in acid, add H2O to balance O atoms and H+ to
balance H atoms.
Cr2O722Cr3+ + 7H2O
14H+ + Cr2O722Cr3+ + 7H2O
5. Add electrons to one side of each half-reaction to balance the
charges on the half-reaction.
Fe2+
Fe3+ + 1e6e- + 14H+ + Cr2O722Cr3+ + 7H2O
6. If necessary, equalize the number of electrons in the two halfreactions by multiplying the half-reactions by appropriate
coefficients.
6Fe2+
6Fe3+ + 6e6e- + 14H+ + Cr2O722Cr3+ + 7H2O
Balancing Redox Reactions
7. Add the two half-reactions together and balance the final
equation by inspection. The number of electrons on both
sides must cancel.
Oxidation:
6Fe2+
Reduction: 6e- + 14H+ + Cr2O7214H+ + Cr2O72- + 6Fe2+
6Fe3+ + 6e2Cr3+ + 7H2O
6Fe3+ + 2Cr3+ + 7H2O
8. Verify that the number of atoms and the charges are balanced.
14x1 – 2 + 6x2 = 24 = 6x3 + 2x3
9. For reactions in basic solutions, add OH- to both sides of the
equation for every H+ that appears in the final equation.
Galvanic Cells
Oxidation-Reduction Reaction
Oxidizing agent and Reducing agent are not in
contact
External conducting medium
Galvanic Cell/Volatic Cell- the experimental
apparatus for generating electricity through the
use of a spontaneous reaction.
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Voltmeter
e–
Zinc
anode
e–
Copper
cathode
Cl– K +
Salt
bridge
SO 2–
4
Zn 2+
Zn2+
ZnSO4
solution
Cotton
plugs
Cu2+
SO 2–
4
CuSO4
solution
Cu2+
2e–
Cu
Zn is oxidized to
Zn2+ at anode.
Zn(s) Zn2+ (aq) + 2 e–
Cu2+ is reduced
to Cu at cathode.
2e– + Cu2+ (aq)
Net reaction
Zn2+ (aq) + Cu(s)
Zn(s) + Cu2+ (aq)
Cu(s)
Galvanic Cells
Cell Voltage- the difference in electrical
potential between the anode and the cathode.
Measured by a voltmeter.
Electromotive force (emf) and Cell Potential
Voltage of a cell depends on:
Composition
of electrode and ions
Concentration of ions
Temperature
Galvanic Cells
Cell Diagram
Zn (s) + Cu2+ (aq)
Cu (s) + Zn2+ (aq)
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
anode
cathode
Standard Reduction Potentials
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.
Used to calculate the emf of a galvanic
cell.
Use the values from anode and cathode
0 )
Standard emf (Ecell
0
0 = E0
Ecell
cathode - Eanode
Standard Reduction Potentials
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
Anode (oxidation):
Zn (s)
Cathode (reduction): 2e- + 2H+ (1 M)
Zn (s) + 2H+ (1 M)
Zn2+ (1 M) + 2eH2 (1 atm)
Zn2+ + H2 (1 atm)
Standard Reduction Potentials
0 = 0.76 V
Ecell
0 )
Standard emf (Ecell
0
0 = E0
Ecell
cathode - Eanode
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
0 = E 0 + - E 0 2+
Ecell
H /H2
Zn /Zn
0 2+
0.76 V = 0 - EZn
/Zn
0 2+
EZn
/Zn = -0.76 V
Zn2+ (1 M) + 2e-
Zn
E0 = -0.76 V
Standard Reduction Potentials
0 = 0.34 V
Ecell
0
0 = E0
Ecell
cathode - Eanode
0 = E 0 2+
0
Ecell
Cu /Cu – EH +/H 2
0 2+
0.34 = ECu
/Cu - 0
0 2+
ECu
/Cu = 0.34 V
Pt (s) | H2 (1 atm) | H+ (1 M) || Cu2+ (1 M) | Cu (s)
Anode (oxidation):
H2 (1 atm)
Cathode (reduction): 2e- + Cu2+ (1 M)
H2 (1 atm) + Cu2+ (1 M)
2H+ (1 M) + 2eCu (s)
Cu (s) + 2H+ (1 M)
19.3
Tips
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
Spontaneity of Redox Reactions
Relationship between Eºcell, ΔGº and K.
0
DG0 = -nFEcell
n = number of moles of electrons in reaction
J
F = 96,500
= 96,500 C/mol
V • mol
0
Ecell
0.0257 V
ln K
=
n
0
Ecell
0.0592 V
log K
=
n
Spontaneity of Redox Reactions
Relationships between Eºcell, ΔGº
and K
The Effect of Concentration on Cell
Emf
Not all reactions in galvanic cells can
occur under standard state conditions.
Nernst Equation- equation used to
calculate cell potential under nonstandard
state conditions.
E = E0 -
0.0257 V
ln Q
n
E = E0 -
0.0592 V
log Q
n
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)
Oxidation:
Cd
Reduction: 2e- + Fe2+
0
0
E0 = EFe
2+/Fe – ECd2+/Cd
E0 = -0.44 – (-0.40)
E0 = -0.04 V
Cd2+ + 2en=2
2Fe
0.0257 V
ln Q
n
0.010
0.0257 V
ln
E = -0.04 V 2
0.60
E = 0.013
E = E0 -
E>0
Spontaneous
19.5
Concentration Cells
Galvanic cells consisting of the same type of
electrodes in the same solution. Each solution is
a different concentration.
Zn(s) | Zn2+(0.10M) || Zn2+(1.0M) | Zn(s)
E = Eº - 0.0257V/ 2 x ln [Zn2+] dil / [Zn2+]conc
E = 0- 0.0257V/2 ln (0.10/1.0)
E = 0.0296 V
Concentration Cells
Importance
Unequal
ion concentrations
Membrane potential
Propagation
Electrolysis
Electrolysis- is the process in which
electrical energy is used to cause a
nonspontaneous chemical reaction to
occur.
Electrolytic Cell- an apparatus for
carrying out electrolysis.