ELECTRON TRANSFER
Reduction-Oxidation RX (redox)
A reaction in which electrons are transferred from one species to
another.
Combustion reactions are redox reactions
- oxidation means the loss of electrons
- reduction means the gain of electrons
- electrolyte is a substance dissolved in water which
produces an electrically conducting solution
- nonelectrolyte is a substance dissolved in water
which does not conduct electricity.
Rusting is a redox reaction:
4Fe(s) + 302(g)  2Fe2O3(s)
Electrochemistry involves redox reactions:
Cu(s) + 2AgNO3(aq)  2Ag(s) + Cu(NO3)2(aq)
REDOX Rx
(reduction-oxidation reactions)
“transfer of electrons”
Oxidation = Loss of electrons
Reduction = Gain of electrons
2Mg + O2  2MgO
OXIDATION STATES
1. All elements have an oxidation number of zero.
2. Group I and II have +1, +2 respectively
3. Oxygen has -2, Flouride has -1
0
3+
-2
0
3+ 2-
2Al + Fe2O3  2Fe + Al2O3
IDENTIFING REDOX RX
Element + compound  New element + New compound
A
+
BC

B
+
AC
Element + Element  Compound
A
+ B  AB
Check oxidation state (charges) of species
A change in oxidation # means redox reaction
Identify the Redox Rx:
Cu + AgNO3  Cu(NO3)2 + Ag
NO + O2  NO2
K2SO4 + CaCl2 KCl + CaSO4
C2H4O2 + O2  CO2 + H2O
LABELING COMPONENTS OF REDOX REACTIONS
The REDUCING AGENT is the species which
undergoes OXIDATION.
The OXIDIZING AGENT is the species which
undergoes REDUCTION.
CuO
+
H2

Cu +
H2O
A summary of redox terminology.
OXIDATION
One reactant loses electrons.
Zn loses electrons.
Reducing agent is oxidized.
Zn is the reducing
agent and becomes
oxidized.
Oxidation number increases.
The oxidation number
of Zn increases from x
to +2.
REDUCTION
Other reactant gains
electrons.
Oxidizing agent is reduced.
Hydrogen ion gains
electrons.
Hydrogen ion is the oxidizing agent
and becomes reduced.
Oxidation number decreases. The oxidation number of H
decreases from +1 to 0.
Key Points About Redox Reactions
•Oxidation (electron loss) always
accompanies reduction (electron gain).
•The oxidizing agent is reduced, and the
reducing agent is oxidized.
•The number of electrons gained by the
oxidizing agent always equals the number
lost by the reducing agent.
REDOX REACTIONS
For the following reactions, identify the
oxidizing and reducing agents.
MnO4- + C2O42-  MnO2 + CO2
acid: Cr2O72- + Fe2+  Cr3+ + Fe3+
base: CO2+ + H2O2  CO(OH)3 + H2O
As + ClO3-  H3AsO3 + HClO
Strongest
oxidizing
agent
Oxidizing/Reducing Agents
Most positive values of E° red
F2(g) +
2e-

•
•

H2(g)
•
+

Li(s)
Most negative values of E° red
Li+(aq)
Increasing
strength of
reducing
agent
•
2H+(aq) + 2e-
Increasing
strength of
oxidizing
agent
2F-(aq)
e-
Strongest
reducing
agent
More Positive
E
º
Cathode(reduction)
Red
Eº Red (cathode)
(V)
Eº cell
Eºred (anode)
Anode(oxidation)
More Negative
Standard Reduction Potentials in Water at 25°C
Standard Potential (V)
2.87
1.51
1.36
1.33
1.23
1.06
0.96
0.80
0.77
0.68
0.59
0.54
0.40
0.34
0
-0.28
-0.44
-0.76
-0.83
-1.66
-2.71
-3.05
Reduction Half Reaction
F2(g) + 2e-  2F-(aq)
MnO4-(aq) + 8H+(aq) + 5e-  Mn2+(aq) + 4H2O(l)
Cl2(g) + 2e-  2Cl-(aq)
Cr2O72-(aq) + 14H+(aq) + 6e-  2Cr3+(aq) + 7H2O(l)
O2(g) + 4H+(aq) + 4e-  2H2O(l)
Br2(l) + 2e-  2Br-(aq)
NO3-(aq) + 4H+(aq) + 3e-  NO(g) + H2O(l)
Ag+(aq) + e-  Ag(s)
Fe3+(aq) + e-  Fe2+(aq)
O2(g) + 2H+(aq) + 2e-  H2O2(aq)
MnO4-(aq) + 2H2O(l) + 3e-  MnO2(s) + 4OH-(aq)
I2(s) + 2e-  2I-(aq)
O2(g) + 2H2O(l) + 4e-  4OH-(aq)
Cu2+(aq) + 2e-  Cu(s)
2H+(aq) + 2e-  H2(g)
Ni2+(aq) + 2e-  Ni(s)
Fe2+(aq) + 2e-  Fe(s)
An2+(aq) + 2e-  Zn(s)
2H2O(l) + 2e-  H2(g) + 2OH-(aq)
Al3+(aq) + 3e-  Al(s)
Na+(aq) + e-  Na(s)
Li+(aq) + e-  Li(s)
ELECTROCHEMISTRY
A system consisting of electrodes that dip into an
electrolyte and in which a chemical reaction uses or
generates an electric current.
Two Basic Types of Electrochemical cells:
Galvanic (Voltaic) Cell:
A spontaneous reaction generates an electric
current. Chemical energy is converted into electrical
energy
Electrolytic Cell:
An electric current drives a nonspontaneous
reaction. Electrical energy is converted into chemical
energy.
ELECTROCHEMICAL CELLS
CHEMICALS AND EQUIPMENT NEEDED TO BUILD A
SIMPLE CELL:
The Cell:
Voltmeter
Two beakers or glass jars
Two alligator clips
The Electrodes:
Metal electrode
Metal salt solution
The Salt Bridge:
Glass or Plastic u-tube
Na or K salt solution
ELECTROCHEMICAL CELLS
A CHEMICAL CHANGE PRODUCES ELECTRICITY
Theory:
If a metal strip is placed in a solution of it’s metal ions, one
of the following reactions may occur
Mn+ + ne-  M
M
 Mn+ + neThese reactions are called half-reactions or half cell reactions
If different metal electrodes in their respective solutions were connected by a
wire, and if the solutions were electrically connected by a porous membrane
or a bridge that minimizes mixing of the solutions, a flow of electrons will
move from one electrode, where the reaction is
M1  M1n+ + neTo the other electrode, where the reaction is
M2n+ + ne-  M2
The overall reaction would be
M1 + M2n+  M2 + M1n+
General characteristics of voltaic and electrolytic cells.
VOLTAIC CELL
System
Energydoes
is released
work on
from
its
spontaneous
surroundings
redox reaction
Oxidation half-reaction
X
X+ + e-
ELECTROLYTIC CELL
Surroundings(power
Energy is absorbed tosupply)
drive a
nonspontaneous
redox reaction
do work on system(cell)
Oxidation half-reaction
AA + e-
Reduction half-reaction
Y++ e- Y
Reduction half-reaction
B++ eB
Overall (cell) reaction
X + Y+
X+ + Y; DG < 0
Overall (cell) reaction
A- + B+
A + B; DG > 0
The electrolysis of water: nonspontaneous
Overall (cell) reaction
2H2O(l)
H2(g) + O2(g)
Oxidation half-reaction
2H2O(l) 4H+(aq) + O2(g) + 4e-
Reduction half-reaction
2H2O(l) + 4e2H2(g) + 2OH-(aq)
The spontaneous reaction between zinc and copper(II) ion.
ACTIVITY SERIES OF SOME SELECTED METALS
A brief activity series of selected metals, hydrogen and halogens are shown
below. The series are listed in descending order of chemical reactivity, with the most
active metals and halogens at the top (the elements most likely to undergo oxidation).
Any metal on the list will replace the ions of those metals (to undergo reduction) that
appear anywhere underneath it on the list.
METALS
HALOGENS
K (most oxidized
Ca
Na
Mg
Al
Zn
Fe
Ni
Sn
Pb
H
Cu
Ag
Hg
Au(least oxidized)
F2
Cl2
Br2
l2
Oxidation refers to the loss of
electrons and reduction refers to the
gain of electrons
The reaction of calcium in water: spontaneous
Oxidation half-reaction
Ca(s)
Ca2+(aq) + 2e-
Reduction half-reaction
2H2O(l) + 2eH2(g) + 2OH-(aq)
Overall (cell) reaction
Ca(s) + 2H2O(l)
Ca2+(aq) + H2(g) + 2OH-(aq)
Activity Series
Relative Reactivities (Activities) of Metals
1. Metals that can displace H from acid
2. Metals that cannot displace H from acid
3. Metals that can displace H from water
4. Metals that can displace other metals from
solution
K
Ba
Ca
Na
Mg
Al
Zn
Cr
Fe
Ni
Sn
Pb
H2
Cu
Ag
Hg
Au
A voltaic cell based on the zinc-copper reaction.
Oxidation half-reaction
Zn(s)
Zn2+(aq) + 2e-
Reduction half-reaction
Cu2+(aq) + 2eCu(s)
Overall (cell) reaction
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
Why Does a Voltaic Cell Work?
The spontaneous reaction occurs as a result of the different
abilities of materials (such as metals) to give up their electrons
and the ability of the electrons to flow through the circuit.
Ecell > 0 for a spontaneous reaction
1 Volt (V) = 1 Joule (J)/ Coulomb (C)
Notation for a Voltaic Cell
components of
anode compartment
components of
cathode compartment
(oxidation half-cell)
(reduction half-cell)
phase of lower phase of higher
oxidation state oxidation state
phase of higher
oxidation state
phase of lower
oxidation state
phase boundary between half-cells
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu (s)
Examples:
Zn(s)
Zn2+(aq) + 2e-
Cu2+(aq) + 2e-
Cu(s)
graphite | I-(aq) | I2(s) || H+(aq), MnO4-(aq) | Mn2+(aq) | graphite
inert electrode
NOTATION FOR VOLTAIC CELLS
Zn + Cu2+  Zn2+ + Cu
Zn(s)/Zn2+(aq) //
Cu2+(aq)/Cu(s)
Anode
Cathode
oxidation
reduction
salt bridge
write the net ionic equation for:
Al(s)/Al3+(aq)//Cu2+(aq)/Cu(s)
Sample Problem:
PROBLEM:
PLAN:
Diagramming Voltaic Cells
Diagram, show balanced equations, and write the notation for a
voltaic cell that consists of one half-cell with a Cr bar in a Cr(NO3)3
solution, another half-cell with an Ag bar in an AgNO3 solution, and
a KNO3 salt bridge. Measurement indicates that the Cr electrode is
negative relative to the Ag electrode.
Identify the oxidation and reduction reactions and write each halfreaction. Associate the (-)(Cr) pole with the anode (oxidation) and the
(+) pole with the cathode (reduction).
Voltmeter
e-
SOLUTION:
Oxidation half-reaction
Cr(s)
Cr3+(aq) + 3e-
salt bridge
Cr
Ag
K+
NO3-
Reduction half-reaction
Ag+(aq) + eAg(s)
Cr3+
Ag
+
Overall (cell) reaction
Cr(s) + Ag+(aq)
Cr3+(aq) + Ag(s)
Cr(s) | Cr3+(aq) || Ag+(aq) | Ag(s)
Voltages of Some Voltaic Cells
Voltaic Cell
Voltage (V)
Common alkaline battery
1.5
Lead-acid car battery (6 cells = 12V)
2.0
Calculator battery (mercury)
1.3
Electric eel (~5000 cells in 6-ft eel = 750V)
0.15
Nerve of giant squid (across cell membrane)
0.070
Alkaline Battery
Mercury and Silver (Button) Batteries
Lithium battery.
Nickel-metal hydride (Ni-MH) battery
Lithium-ion battery.
A voltaic cell using inactive electrodes.
Oxidation half-reaction
2I-(aq)
I2(s) + 2e-
Reduction half-reaction
MnO4-(aq) + 8H+(aq) + 5eMn2+(aq) + 4H2O(l)
Overall (cell) reaction
2MnO4-(aq) + 16H+(aq) + 10I-(aq)
2Mn2+(aq) + 5I2(s) + 8H2O(l)
The Hydrogen Electrode:
At the hydrogen electrode, the half reaction involves
a gas.
2 H+(aq) + 2e-  H2(g)
so an inert material must serve as the reaction site
(Pt)
H+(aq)/H2(g)/Pt cathode
Pt/H2(g)/H+(aq) anode
Therefore:
Al(g)/Al3+(aq)//H+(aq)/H2(g)/Pt
Determining an unknown E0half-cell with the standard
reference (hydrogen) electrode.
Oxidation half-reaction
Zn(s) Zn2+(aq) + 2e-
Overall (cell) reaction
Zn(s) + 2H3O+(aq) Zn2+(aq) + H2(g) + 2H2O(l)
Reduction half-reaction
2H3O+(aq) + 2eH2(g) + 2H2O(l)
The laboratory measurement of pH.
Pt
Glass
electrode
Reference
(calomel)
electrode
Hg
Paste of
Hg2Cl2 in
Hg
AgCl on
Ag on Pt
1M HCl
Thin glass
membrane
KCl
solution
Porous ceramic
plugs
Rechargeable Batteries
The tin-copper reaction as the basis of a voltaic and an electrolytic
cell.
voltaic cell
Oxidation half-reaction
Sn(s) Sn2+(aq) + 2eReduction half-reaction
Cu2+(aq) + 2eCu(s)
Overall (cell) reaction
Sn(s) + Cu2+(aq) Sn2+(aq) + Cu(s)
electrolytic cell
Oxidation half-reaction
Cu(s)
Cu2+(aq) + 2eReduction half-reaction
Sn2+(aq) + 2eSn(s)
Overall (cell) reaction
Sn(s) + Cu2+(aq) Sn2+(aq) + Cu(s)
Lead-acid battery.
The processes occurring during the discharge and recharge of a
lead-acid battery.
VOLTAIC(discharge)
ELECTROLYTIC(recharge)
The corrosion of iron.
OTHER APPLICATIONS
FOR REDOX REACTIONS
Corrosion
Photosynthesis
Cleaning agents and other household supplies
Chemicals
The effect of metal-metal contact on the corrosion of iron.
faster corrosion
In Photosynthesis, carbohydrates are produced in the
green leaves of plants
carbon cycle
6CO2 + 6H2O + hv
(photosynthesis)

C6H12O6 + 6O2
(respiration)
glucose
In body tissues, Glucose is oxidized in metabolic
reactions (respiration) and chemical energy is released to
do work in the cells.
Polysaccharides react with water to yield
monosaccharides, (Hydrolysis)
H+
C12H22O11 + H2O  C6H12O6 + C6H12O6
sucrose
glucose
fructose
COMMON OXIDIZING AGENTS
-
O2 (gaseous)
H2O2 (antiseptics) tincture of Iodine
K2Cr2O7 (breathalyzer)
benezoyl peroxide
Ca (Ocl)2 (swimming pools) disinfectant
NaOCl (bleaches)
COMMON REDUCING AGENTS
- coal (coke) for producing metals from ores
- more reactive metals (ore)
- H2 (ores & hydrocarbons)
- B/W photograph developer
C6H4(OH)2 + 2 Ag+  C6H4O2 + 2Ag + 2 H+
- antioxidants
ascorbic acid “C”
tocopherol “E”
beta-carotene “A”
HYDROGEN H2
- most abundent element in universe
- rare on earth
- flammable as an element
- usually found as compound
Acids = HCl
hydrocarbons = CH4
water = H2O
- many times required a catalyst to react.
Catalyst speeds up the reaction but is not consumed.
(Pt, Ni, Pd)
Pt
CuO + H2  Cu + H2O
- produced by reacting an acid with a reactive metal
Mg + 2HCl  MgCl2 + H2
OXYGEN O2
- occurs in atmosphere, hydrosphere , and lithosphere
O2
H2O
SiO2
- percent composition of H2O = 89% O
2/3 body weight is O
- needed for combustion
HC + O2  CO2 + H2O
- used in rusting or corrosion
2Cu + O2  2CuO
4Fe + 3O2  2Fe2O3
- nonmetal oxides
C + O2  CO2
S + O2  SO2
- needed for burning