Transcript Document
Electrochemistry
Chapter 19
Electron Transfer Reactions
•
Electron transfer reactions are oxidation reduction or redox reactions.
•
Results in the generation of an electric current (electricity) or be caused by imposing an electric current.
•
Therefore, this field of chemistry is often called 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 0 0
2Mg (
s
) + O
2 2+ 2-
(
g
) 2MgO (
s
)
2Mg 2Mg 2+ + 4e -
Oxidation
half-reaction (lose e ) O 2 + 4e 2O 2-
Reduction
half-reaction (gain e ) 19.1
Terminology for Redox Reactions
• • • •
OXIDATION —loss of electron(s) by a species; increase in oxidation number; increase in oxygen.
REDUCTION —gain of electron(s); decrease in oxidation number; decrease in oxygen; increase in hydrogen.
OXIDIZING AGENT —electron acceptor; species is reduced. (an agent facilitates something; ex. Travel agents don’t travel, they facilitate travel) REDUCING AGENT —electron donor; species is oxidized.
You can’t have one… without the other!
• Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons.
• You can’t have 2 oxidations or 2 reductions in the same equation. Reduction has to occur at the cost of oxidation
LEO
o s e l e c r t o n s i t o n x i d a
the lion says
GER
!
a i n l e c r t o n s e d u c i t o n GER!
Another way to remember
•OIL RIG
x i d a i t o n s o s e e d u c i t o n s a i n
Review of Oxidation numbers
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, H
2
, O
2
, P
4
=
0 2. In monatomic ions, the oxidation number is equal to the charge on the ion.
Li
+
, Li =
+1
; Fe
3+
, Fe =
+3
; O
2-
, O =
-2 3. The oxidation number of oxygen is usually and O 2 2 it is –1 . –2 . In H 2 O 2 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 .
5. Group IA metals are +1 , IIA metals are +2 always –1 .
and fluorine is 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.
HCO
3 Oxidation numbers of all the atoms in HCO 3 ?
O =
-2
H =
+1
3x(
-2
) +
1
+
?
= -1 C =
+4 4.4
Balancing Redox Equations
The oxidation of Fe 2+ to Fe 3+ by Cr 2 O 7 2 in acid solution?
1. Write the unbalanced equation for the reaction ion ionic form.
Fe 2+ + Cr 2 O 7 2 Fe 3+ + Cr 3+ 2. Separate the equation into two half-reactions.
Oxidation: Reduction: +2 Fe 2+ +6 Cr 2 O 7 2 +3 Fe 3+ +3 Cr 3+ 3. Balance the atoms other than O and H in each half-reaction.
Cr 2 O 7 2 2Cr 3+ 19.1
Balancing Redox Equations
4. For reactions in acid, add H 2 O to balance O atoms and H + balance H atoms.
to Cr 2 O 7 2 14H + + Cr 2 O 7 2 2Cr 2Cr 3+ 3+ + 7H + 7H 2 2 O O 5. Add electrons to one side of each half-reaction to balance the charges on the half-reaction.
Fe 2+ 6e + 14H + + Cr 2 O 7 2 Fe 3+ 2Cr 3+ + 1e + 7H 2 O 6. If necessary, equalize the number of electrons in the two half reactions by multiplying the half-reactions by appropriate coefficients.
6Fe 2+ 6Fe 3+ + 6e 6e + 14H + + Cr 2 O 7 2 2Cr 3+ + 7H 2 O 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. You should also cancel like species.
Oxidation: 6Fe 2+ 6Fe 3+ + 6e Reduction: 14H + + Cr 2 O 6e 7 2 + 14H + 6Fe + 2+ + Cr 2 O 7 2 2Cr 3+ + 7H 2 O 6Fe 3+ + 2Cr 3+ + 7H 2 O 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 equation for every H + to
both sides
of the that appears in the final equation. You should combine H + and OH to make H 2 O.
19.1
CHEMICAL CHANGE ---> ELECTRIC CURRENT
•To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire.
This is accomplished in a GALVANIC VOLTAIC cell.
or
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf
A group of such cells is called a
battery
.
anode oxidation -
Galvanic Cells
cat
hode
red
uction + spontaneous redox reaction 19.2
Galvanic Cells
The difference in electrical potential between the anode and cathode is called: •
cell voltage
•
electromotive force (emf)
•
cell potential
Zn (
s
) + Cu 2+ Cell Diagram (
aq
) Cu (
s
) + Zn 2+ (
aq
) [Cu 2+ ] = 1
M
& [Zn 2+ ] = 1
M
Zn (
s
) | Zn 2+ (1
M
) || Cu 2+ (1
M
) | Cu (
s
) anode cathode 19.2
Standard Electrode Potentials
Zn (
s
) | Zn 2+ (1
M
) || H + (1
M
) | H 2 (1 atm) | Pt (
s
) Anode (oxidation): Zn (
s
) Zn 2+ (1
M
) + 2e Cathode (reduction): 2e + 2H + (1
M
) H 2 (1
atm
) Zn (
s
) + 2H + (1
M
) Zn 2+ + H 2 (1
atm
) 19.3
Standard Electrode Potentials
Standard reduction potential (E 0 )
is the voltage associated with a
reduction reaction
at an electrode when all solutes are 1
M
and all gases are at 1 atm.
Reduction Reaction 2e + 2H + (1
M
) H 2 (1
atm
)
E 0
= 0 V Standard hydrogen electrode (SHE) 19.3
E θ
TABLE 4B: STANDARD REDUCTION POTENTIALS/
TABEL 4B: STANDAARD REDUKSIEPOTENSIALE
Li + + e K + Ba + e 2+ + 2e Ca 2+ + 2e Na + + e Mg 2+ + 2e Aℓ 3+ + 3e Mn 2+ + 2e 2H 2 O + 2e Zn 2+ + 2e Cr 2+ + 2e Cr 3+ + 3e Fe 2+ + 2e Cr 3+ + e Cd 2+ + 2e Co 2+ + 2e Ni 2+ + 2e Sn 2+ Pb 2+ + 2e + 2e Fe 3+ + 3e 2H + + 2e S + 2H + + 2e Sn 4+ + 2e Cu 2+ + e SO + 4H + + 4e Cu 2+ + 2e 2H 2 O + O 2 + 4e SO 2 + 4H + + 2e Cu + + e I 2 + 2e O 2 (g) + 2H + + 2e Fe 3+ + e NO + 2H + + e Hg 2+ + 2e Ag + + e NO + 4H + + 3e Br 2 (ℓ) + 2e O 2 (g) + 4H + + 3e MnO 2 + 4H + + 2e Cr 2 O + 14H + + 6e Cℓ 2 (g) + 2e MnO + 8H + + 5e Co 3+ + e F 2 (g) + 2e
Half-reactions/Halfreaksies
⇌ Li ⇌ K ⇌ Ba ⇌ Ca ⇌ Na ⇌ ⇌ ⇌ ⇌ ⇌ Mg Aℓ Mn H 2 (g) + 2OH Zn ⇌ ⇌ ⇌ ⇌ ⇌ ⇌ ⇌ ⇌ ⇌ ⇌ ⇌ ⇌ ⇌ ⇌ Cr Cr Fe Cr 2+ Cd Co Ni Sn Pb Fe H 2 (g) H 2 S(g) Sn 2+ Cu + ⇌ SO 2 (g) + 2H 2 O ⇌ Cu ⇌ 4OH ⇌ ⇌ ⇌ ⇌ ⇌ ⇌ ⇌ ⇌ ⇌ ⇌ ⇌ S + 2H 2 O Cu 2I H 2 O 2 Fe 2+ NO 2 (g) + H 2 O Hg(ℓ) Ag NO(g) + 2H 2 O 2Br 2H 2 O ⇌ ⇌ ⇌ ⇌ ⇌ ⇌ Mn 2+ + 2H 2 O 2Cr 3+ + 7H 2 O 2 Cℓ Mn 2+ + 4H 2 O Co 2+ 2F
(V)
3,04 2,92 2,90 2,87 2,71 2,37 1,66 1,18 0,83 0,76 0,74 0,74 0,44 0,41 0,40 0,28 0,25 0,14 0,13 0,04 0,00 + 0,14 + 0,15 + 0,16 + 0,17 + 0,34 + 0,40 + 0,45 + 0,52 + 0,54 + 0,68 + 0,77 + 0,78 + 0,78 + 0,80 + 0,96 + 1,06 + 1,23 + 1,28 + 1,33 + 1,36 + 1,52 + 1,82 + 2,87 • • • • •
E 0
is for the reaction as written The more positive
E 0
the greater the tendency for the substance to be reduced The half-cell reactions are reversible The sign of
E 0
changes when the reaction is reversed Changing the stoichiometric coefficients of a half-cell reaction
does not
the value of
E 0
change 19.3
Standard Electrode Potentials
Standard emf (E 0 cell )
E 0 cell
=
E 0 cathode E 0 anode
If the reaction is backwards, be sure to flip the sign!
Zn (
s
) | Zn 2+
E 0 cell
(1
M
) || H + =
E 0 H /H
2
+ E 0 Zn /Zn
+2 (1
M
) | H 2 (1 atm) | Pt (
s
) Zn 2+ (1
M
) + 2e Zn
E 0
= -0.76 V So
E o Zn/Zn
+2
= + 0.76 V E 0 cell
= 0 + 0.76 V = 0.76 V 19.3
Standard Electrode Potentials
E 0 cell
= 0.34 V
E 0 cell
=
E 0 cathode E 0 anode E 0 cell
=
E 0 Cu /Cu
-
0 E H /H+
2 0.34 =
E 0 Cu /Cu
- - 0
E 0 Cu /Cu
= 0.34 V Pt (
s
) | H 2 (1
atm
) | H + (1
M
) || Cu 2+ (1
M
) | Cu (
s
) Anode (oxidation): H 2 (1
atm
) 2H + (1
M
) + 2e Cathode (reduction): 2e + Cu 2+ (1
M
) Cu (s) H 2 (1
atm
) + Cu 2+ (1
M
) Cu (
s
) + 2H + (1
M
) 19.3
What is the standard emf of an electrochemical cell made of a Cd electrode in a 1.0 electrode in a 1.0
M M
Cd(NO 3 ) 2 Cr(NO 3 ) 3 solution?
solution and a Cr Cd Cr 2+ 3+ (
aq
) (
aq
) + 2e + 3e Anode (oxidation): Cd (
s
) Cr (
s
)
E 0
= -0.40 V
E 0
= -0.74 V Cd is the stronger oxidizer Cd will oxidize Cr Cr (s) Cr 3+ (1
M
) + 3e x 2 Cathode (reduction): 2e + Cd 2+ (1
M
) Cd (s) x 3 2Cr (
s
) + 3Cd 2+ (1
M
) 3Cd (
s
) + 2Cr 3+ (1
M
)
E 0 cell
=
E 0 cathode + E 0 anode E 0 cell
= -0.40 + (+0.74)
E 0 cell
= 0.34 V 19.3
Le Chatelier’s Principal and EMF of a cell
If the forward reaction is favoured then emf increases If the reverse reaction is favoured the emf decreases SO If the concentration of reactants increases relative to products, the cell reaction becomes more spontaneous and the emf increases. As the cell operates, the reactants are used up as more product is formed causing the emf to decrease 19.5
If a cell reaction reaches equilibrium , the cell is said to be flat If the cell reaction goes to completion, (one of the reactants used up) the cell is said to be flat At equilibrium, the reverse rxn pushes electrons out of the cathode at the same rate as the forward rxn pushes electrons out of the anode, and there is no net current flow - the battery is dead. The only thing preventing the voltaic cell from reaching equilibrium is a conductor between anode and cathode (which you supply when you want to use some of the energy in the cell).
The greater the temperature, the greater the emf of a cell A voltaic cell is NOT at equilibrium. To produce a net flow of energetic electrons out of the anode, the forward rxn must be able to occur at a greater rate than the reverse rxn. When the temperature of a voltaic cell increases, both forward and reverse rxn rates increase proportionally, but because the forward rxn rate was already greater than the reverse rxn rate, the forward rxn rate becomes proportionally greater when temperature increases. Ordinarily the reaction rate is limited by the electrical resistance of the path between anode and cathode. In this case, the difference in rxn rates shows up as potential energy. Electrical potential energy per electron is voltage, EMF. When the difference in rxn rates increases, potential energy and EMF of the cell increases 19.5
Charging a Battery
When you charge a battery, you are forcing the electrons backwards (from the + to the -). To do this, you will need a higher voltage backwards than forwards. This is why the ammeter in your car often goes slightly higher while your battery is charging, and then returns to normal.
In your car, the battery charger is called an alternator. If you have a dead battery, it could be the battery needs to be replaced OR the alternator is not charging the battery properly.
Dry cell
Leclanch é cell
Batteries
Anode: Zn (s) Zn 2+ (
aq
) + 2e Cathode: + 2NH 4 (
aq
) + 2MnO 2 (
s
) + 2e Mn 2 O 3 (
s
) + 2NH 3 (
aq
) + H 2 O (
l
) Zn (
s
) + 2NH 4 (
aq
) + 2MnO 2 (
s
) Zn 2+ (
aq
) + 2NH 3 (
aq
) + H 2 O (
l
) + Mn 2 O 3 (
s
) 19.6
Mercury Battery
Batteries
Anode: Cathode: Zn(Hg) + 2OH (
aq
) ZnO (
s
) + H 2 O (
l
) + 2e HgO (
s
) + H 2 O (
l
) + 2e Hg (
l
) + 2OH Zn(Hg) + HgO (
s
) ZnO (
s
) + Hg (
l
) (
aq
) 19.6
Batteries
Lead storage battery Anode: Pb (s) + SO 2 4 (
aq
) PbSO 4 (
s
) + 2e Cathode: PbO 2 (
s
) + 4H + (
aq
) + SO 2 4
(aq
)
+
2e PbSO 4 (
s
) + 2H Pb (
s
) + PbO 2 (
s
) + 4H + (
aq
) + 2SO 2 4 (
aq
) 2PbSO 4 (
s
) + 2H 2 O (
l
) 2 O (
l
) 19.6
Batteries
Solid State Lithium Battery 19.6
Batteries
A
fuel cell
is an electrochemical cell that requires a continuous supply of reactants to keep functioning Anode: Cathode: 2H 2 (
g
) + 4OH (
aq
) 4H 2 O (
l
) + 4e O 2 (
g
) + 2H 2 O (
l
) + 4e 4OH (
aq
) 2H 2 (
g
) + O 2 (
g
) 2H 2 O (
l
) 19.6
Corrosion
19.7
Cathodic Protection of an Iron Storage Tank
19.7
Electrolysis
is the process in which electrical energy is used to cause a
nonspontaneous
chemical reaction to occur.
19.8
Electrolysis of Water 19.8
Chemistry In Action:
Dental Filling Discomfort 2+ Hg 2 /Ag 2 Hg 3 0.85 V 2+ Sn /Ag 3 Sn -0.05 V 2+ Sn /Ag 3 Sn -0.05 V