Transcript File - AP CHEMISTRY

Oxidation-Reduction Reactions
(REDOX) & Electrochemistry
What are some examples of
Redox Reactions?
What are some examples
of redox reactions?
What everday uses depend
on redox reactions?
•fires
•starting a car
•rusting steel
•calculator
•combustion in auto
engine
•digital watch
•metabolism of food in
the body
•portable radio
•portable CD player
Redox
Oxidation is the loss of electrons--oxidation
number becomes more positive.
Reduction is the gain of electrons--oxidation
number becomes more negative.
OIL RIG
Oxidation Is Loss.
Reduction Is Gain.
Redox Reactions
Loss and gain of electrons must be
simultaneous.
Loss and gain of electrons must be equal.
Why must the loss and gain of electrons be
equal? Law of Conservation of Matter
Basic Terminology
OXIDATION — loss of electron(s) by a species;
increase in oxidation number
REDUCTION — gain of electron(s); decrease in
oxidation number
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.
Oxidation Numbers
In order to keep
track of what loses
electrons and what
gains them, we
assign oxidation
numbers or
oxidation states.
Oxidation and Reduction
A species is oxidized when it loses electrons.
 Here,
zinc loses two electrons to go from neutral zinc
metal to the Zn2+ ion.
A species is reduced when it gains electrons.
 Here,
gas.
each of the H+ gains an electron to go form H2
Review of Oxidation States
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
peroxides such as H2O2 or Na2O2 it is –1.
4. The oxidation number of hydrogen is +1 except when
it is bonded to metals in binary compounds. In these
cases such as KH, 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. Basically, compound add to zero,
polyatomic ions add to its charge.
HCO3-
Oxidation numbers of all
the atoms in HCO3- ?
O = -2
H = +1
3x(-2) + 1 + ? = -1
C = +4
4.4
Charges & Oxidation States
Oxidation states are written as +2.
2
Ca
Charges are written 2+.
Ca
2
Determining Oxidation States
SF6
+6
-6 = 0
(-1 for each F)
[NO3]-
+5 -6 = -1
(-2 for each O)
The most electronegative element is assigned
a negative oxidation number!
Identifying Oxidation & Reduction
in a Reaction
Identify the element which is oxidized and the one
which is reduced.
3 CuCl2 + 2 Al  2 AlCl3 + 3 Cu
2Mg (s) + O2 (g)  2 MgO(s)
Redox
Oxidizing agent is the electron acceptor-usually a nonmetal.
Reducing agent is the electron donor--usually
a metal.
-4 +1
0
+4 -2
+1-2+1
CH4(g) + 2O2(g) ----> CO2(g) + 2HOH(g)
Carbon is oxidized.
CH4 is the reducing agent.
Oxygen is reduced.
O2 is the oxidizing agent.
Redox Reactions
Identify the substance oxidized and the substance reduced
as well as the oxidizing and reducing agents.
+2 -2
+2 -2
0
+4 -2
PbO(s) + CO(g) ---> Pb(s) + CO2(g)
oxidized
Carbon
reduced
Lead
oxidizing agent
PbO
reducing agent
CO
Redox Reactions
Identify the substance oxidized and the substance reduced
as well as the oxidizing and reducing agents.
+2 -2
0
+2 -2
+4 -2
2PbS(s) + 3O2(g) ---> 2PbO(s) + 2SO2(g)
oxidized
sulfur
reduced
oxygen
oxidizing agent
O2
reducing agent
PbS
https://www.youtube.com/watch?v=poV6lc2b070
Redox Reactions
Redox reactions are reactions in which
electrons are transferred.
Decomposition and synthesis reactions may be
redox.
Single replacement reactions are always redox.
Double replacement reactions are never redox.
Combustion reactions are always redox.
Balancing Reaction in Acid
Oxidation of Fe2+ by Cr2O72- in acid solution to form Fe3+ and Cr3+?
1. Write the unbalanced equation for the reaction in ionic form.
Fe2+ + Cr2O72-
Fe3+ + Cr3+
2. Separate the equation into two half-reactions.
+2
Fe2+
Oxidation:
Reduction:
+6
Cr2O72-
+3
Fe3+
+3
Cr3+
3. Balance the atoms other than O and H in each half-reaction.
Cr2O72-
2 Cr3+
Balancing Redox Equations
4. For reactions in acid, add H2O to balance O atoms and H+ to
balance H atoms.
Cr2O7214 H+ + Cr2O72-
2 Cr3+ + 7 H2O
2 Cr3+ + 7 H2O
5. Add electrons to one side of each half-reaction to balance the
charges on the half-reaction.
Fe2+
Fe3+ + 1e6e- + 14 H+ + Cr2O72-
2 Cr3+ + 7 H2O
6. If necessary, equalize the number of electrons in the two halfreactions by multiplying the half-reactions by appropriate
coefficients.
6 Fe2+
6 Fe3+ + 6e6e- + 14 H+ + Cr2O72-
2 Cr3+ + 7 H2O
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:
6 Fe2+
Reduction: 6e- + 14 H+ + Cr2O72-
14H+ + Cr2O72- + 6Fe2+
6 Fe3+ + 6e2 Cr3+ + 7 H2O
6Fe3+ + 2Cr3+ + 7H2O
8. Verify that the number of atoms and the charges are balanced.
14(+1) – 2 + 6(+2) = 24 = 6(+3) + 2(+3)
9. For reactions in basic solutions, add OH- to both sides of the
equation for every H+ that appears in the final equation. You
should combine H+ and OH- to make H2O.
Electrochemistry
The study of the interchange of
chemical and electrical energy.
Voltaic Cells
anode
oxidation
cathode
reduction
-
+
spontaneous
redox reaction
19.2
Electrochemical <---> Electrolytic
Electrochemical <---> Electrolytic
Spontaneous <---> Nonspontaneous
Energy released <---> Energy absorbed
Cu2+(aq) + Mg(s) <---> Cu(s) + Mg2+(aq)
Electrochemical cell -- chemical energy to electrical energy.
Electrolytic cell -- electrical energy to chemical energy.
Half-Reactions
The overall reaction is split into two half-reactions,
one involving oxidation and one reduction.
8H+ + MnO4 + 5Fe2+  Mn2+ + 5Fe3+ + 4H2O
Reduction: 8H+ + MnO4 + 5e  Mn2+ + 4H2O
Oxidation: 5Fe2+  5Fe3+ + 5e
Schematic of a method for separating the oxidizing
and reducing agents in a redox reaction
Electron flow
Ion flow
Electric circuit
Schematic of a battery (voltaic cell)
Electrons are transferred at the interface between the
electrodes and the solution.
voltaic Cell
A device in which chemical
energy is changed to electrical
energy. The basic parts are:
anode
cathode
electrochemical solution
porous disk or salt bridge
Anode and Cathode
OXIDATION occurs at the ANODE.
REDUCTION occurs at the CATHODE.
RED CAT
AN OX
17_363
e–
e–
e–
e–
Zn( s)
+
Zn 2
–
SO 4 2
1.0 M Zn 2
solution
Anode
+
Cu 2
–
SO 4 2
+
1.0 M Cu
solution
Cu( s)
2+
Cathode
A voltaic cell (Daniell Cell) involving Zn and Cu
electrodes. This cell was the energy source for telegraphy
during the War Between the States.
Cell Potential
Cell Potential or Electromotive
Force (emf): The “pull” or driving
force on the electrons.
Cell Potential
Cell potential is measured in volts (V).
J
1V=1
C
Official AP
Reduction Table
How would you
make this an
oxidation table?
Oxidizing and Reducing Agents
The strongest oxidizers
have the most positive
reduction potentials.
The strongest reducers
have the most negative
reduction potentials.
Standard Reduction Potentials
The E values corresponding to reduction halfreactions with all solutions at 1.0 M and all gases
at 1 atm.
Cu2+ + 2e  Cu
E = 0.34 V vs. SHE
SHE = Standard Hydrogen Electrode
(arbitrarily set to zero as a reference point)
Zinc electrode compared to a Standard Hydrogen Electrode
(SHE). The Zn has a potential of 0.76 V.
Cell Potential Calculations
To Calculate cell potential using Standard
Reduction Potentials:
1. One reaction and its cell potential’s sign
must be reversed--it must be chosen such
that the overall cell potential is positive.
2. The half-reactions must often be multiplied
by an integer to balance electrons--this is
not done for the cell potentials.
Cell Potential Calculations
Continued
Fe3+(aq) + Cu(s) ----> Cu2+(aq) + Fe2+(aq)
Fe3+(aq) + e- ----> Fe2+(aq)
Eo = 0.77 V
Cu2+(aq) + 2 e- ----> Cu(s)
Eo = 0.34 V
Reaction # 2 must be reversed.
Cell Potential Calculations
Continued
2 (Fe3+(aq) + e- ----> Fe2+(aq)) Eo = 0.77 V
Cu(s) ----> Cu2+(aq) + 2 e-
Eo = - 0.34 V
2Fe3+(aq) + Cu(s) ----> Cu2+(aq) + 2Fe2+(aq)
Eo = 0.43 V
Cell Potentials
 = E  (cathode) − E  (anode)
Ecell
red
red
17_369
Reference solution of
dilute hydrochloric acid
Silver wire coated with
silver chloride
Thin-walled membrane
Ion selective electrodes are glass electrodes that measures
a change in potential when [H+] varies. Used to measure pH.
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
Cd is the stronger oxidizer
Cr3+ (aq) + 3e-
Cr (s) E0 = -0.74 V
Cd will oxidize Cr
Cr3+ (1 M) + 3e- x 2
Cr (s)
Anode (oxidation):
Cathode (reduction): 2e- + Cd2+ (1 M)
2Cr (s) + 3Cd2+ (1 M)
x3
3Cd (s) + 2Cr3+ (1 M)
0
0
E0cell = Ecathode
+ Eanode
E0cell = -0.40 + (+0.74)
E0cell = 0.34 V
Cd (s)
NATURAL OCCURENCE
Batteries
A battery is a voltaic cell or, more
commonly, a group of voltaic cells
connected in series. Sometimes, a
voltaic cell is referred to as a
galvanic cell.
+
17_370
–
Lead storage
battery
H2SO4
electrolyte
solution
Anode (lead
grid filled with
spongy lead)
Anode:
Cathode:
Pb (s) + SO42- (aq)
Cathode (lead
grid filled with
spongy PbO2)
PbSO4 (s) + 2e-
PbO2 (s) + 4H+ (aq) + SO42- (aq) + 2e-
Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO42- (aq)
PbSO4 (s) + 2H2O (l)
2PbSO4 (s) + 2H2O (l)
17_371
Anode
(zinc inner case)
Cathode
(graphite rod)
Dry cell
Paste of MnO2,
NH4CL, and
carbon
Zn (s)
Anode:
Cathode:
2NH4 +(aq) + 2MnO2 (s) + 2e-
Zn (s) + 2NH4 +(aq) + 2MnO2 (s)
Zn2+ (aq) + 2eMn2O3 (s) + 2NH3 (aq) + H2O (l)
Zn2+ (aq) + Mn2O3 (s) + 2NH3 (aq) + H2O (l)
Fuel cell
A fuel cell is an
electrochemical cell
that requires a
continuous supply
of reactants to keep
functioning
Anode:
Cathode:
2H2 (g) + 4OH- (aq)
O2 (g) + 2H2O (l) + 4e2H2 (g) + O2 (g)
4H2O (l) + 4e4OH- (aq)
2H2O (l)
Corrosion
Some metals, such as copper, gold,
silver and platinum, are relatively
difficult to oxidize. These are often
called “noble” metals.
With other more
valuable metals, why
is Superman, the
“man of steel”?
Corrosion of Iron
Net Ionic Equation:
Protecting Metals
Some metals such as aluminum, copper, and silver
form a protective coating that keeps them from
corroding further.
The protective coating for other metals i.e. iron
flakes away opening new layers of metal to
corrosion.
About 1/5 of all iron produced in industry is used to
replace metal that has oxidized.
Prevention of Corrosion
Coating--painting or applying oil to keep out
oxygen and moisture.
Galvanizing--dipping a metal in a less active
metal -- galvanized steel bucket.
Alloying -- mixing metals with iron to prevent
corrosion -- stainless steel.
Cathodic protection -- attaching a more active
metal. Serves as sacrificial metal--used to
protect ships, gas lines, and gas tanks.
Cathodic Protection
Electrolysis
Process in which
electrical energy is
used to cause a
nonspontaneous
chemical reaction
to occur.
. . . forcing a current through a cell to produce
a chemical change for a setup in which the
cell potential is negative.
Napoleon III of
France served his
most
honored guests
with aluminum
silverware.
Why was
aluminum so darn
expensive?
Al2O3
Bauxite
Electrolysis of Bauxite
6 O2- (l)  3 O2 (g) + 12 e4 Al3+ (l) + 12 e-  4 Al (s)
Molten aluminum sinks and is tapped off at the bottom.
Bauxite-cryolite mixture floats on top and is electrolyzed
by the carbon electrodes.
Electrolysis of Halite
2 Na+ (l) + 2 e-  2 Na (s)
2 Cl- (l)  2 Cl2 (g) + 2 eChlorine gas is a poisonous gas. How could you modify
the setup so that you do not kill everyone in the lab?
19.8
Electrolysis of Water
Why is the
volume of H2
gas greater
than the
volume of O2?
Mathematic of Electrolysis
What determines the amount of products that
can be produced in an electrolytic reaction?
Moles of Electrons
How do we determine how many electrons are
flowing through a wire?
Current (measure in “amps”)
The number of electrons transferred is
proportional to current and time.
Electrolysis and Mass Changes
1 amp = 1 Coulomb / sec
1 mole e- = 96,500 C = 1 Faraday
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:
2Cl- (l)
Ca2+ (l) + 2eCa2+ (l) + 2Cl- (l)
Cl2 (g) + 2eCa (s)
Ca (s) + Cl2 (g)
2 mole e- = 1 mole Ca
C
s
1 mol e- 1 mol Ca
mol Ca = 0.452
x 1.5 hr x 3600
x
x
s
hr 96,500 C
2 mol e= 0.0126 mol Ca
= 0.50 g Ca