Utilizes relationship between
chemical potential energy &
electrical energy
Redox Reactions
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Need battery to start car
Prevent corrosion
Bleach is an oxidizing agent
Na, Al, Cl prepared or purified by redox
reactions
• Breathing
– O2  H2O and CO2
Redox Reactions
• Synthesis
• Decomposition
• Single Replacement
Redox
• Double Replacement only is not redox
Predicting Redox Reactions
• Use Table J to predict if a given redox
reaction will occur.
• Any metal will donate its electrons to the
ion of any metal below it.
• Any nonmetal will steal electrons from the
ion of any nonmetal below it.
Predicting Single Replacement
Redox Reactions
• Element + Compound  New Element +
New
Compound
• If the element is above the swapable ion,
the reaction is spontaneous.
• If the element is below the swapable ion,
the reaction is not spontaneous.
Predicting Redox Reactions
A + BX  B + AX
A & B are metals. If metal A is above
metal B in Table J, the reaction is
spontaneous.
X + AY  Y + AX
• X & Y are nonmetals. If nonmetal X is
above nonmetal Y in Table J, the reaction
is spontaneous.
Which are spontaneous?
Yes
Cs + CuCl2 
Yes
I2 + NaCl 
No
Cl2 + KBr 
Yes
Fe + CaBr2 
No
No
Mg + Sr(NO3)2 
F2 + MgCl2 
Yes
• Li + AlCl3 
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Started with
Zn(NO3)2 &
Cu and
AgNO3 & Cu.
Which
beaker had
the Zn ions &
which had
the Ag ions?
Overview of Electrochemistry
• TWO kinds of cells (kind of opposites):
1. Galvanic or Voltaic (NYS – Electrochemical)
• Use a spontaneous reaction to produce a
flow of electrons (electricity). Exothermic.
2. Electrolytic
• Use a flow of electrons (electricity) to force
a nonspontaneous reaction to occur.
Endothermic.
Vocabulary
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Redox
Half-reaction
Oxidation
Reduction
Cell
Half-Cell
Electrode
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Anode
Cathode
Galvanic
Voltaic
Electrochemical
Electrolytic
Salt bridge
Electrochemical Cells
• Use a spontaneous single replacement
redox reaction to produce a flow of
electrons.
• Electrons flow from oxidized substance to
reduced substance.
• Called: Galvanic cells, voltaic cells, or
electrochemical cells (NYS)
Electrochemical Cells
• Redox reaction is arranged so the
electrons are forced to flow through a wire.
• When the electrons travel through a wire,
we can make them do work, like light a
bulb or ring a buzzer. OJ clock
• So the oxidation & reduction reactions
have to be separated physically.
Al / CuCl2 Lab
• Was a redox reaction.
• Did NOT force electrons to travel through
a wire. Got NO useful work out of system.
• Have to be clever in how we arrange
things.
2Al + 3Cu+2  2Al+3 + 3Cu
Got no useful
work
because halfreactions
weren’t
separated.
Half-Cell
• Where each of the half-reactions takes
place.
• Need 2 half-cells to have a complete redox
reaction.
• Need to be connected by a wire for the
electrons to flow through.
• Need to be connected by a salt bridge to
maintain electrical neutrality.
Schematic of Galvanic Cell
Parts of a Voltaic Cell
• 2 half-cells: oxidation & reduction
• Each half-cell consists of a container of an
aqueous solution & an electrode or
surface at which the electron transfer
takes place.
• Wire connecting electrodes.
• Salt bridge connects solutions.
How much work can you get out of
this reaction?
• You can measure the voltage by making
the electrons travel through a voltmeter.
• The galvanic cell is a battery. Of course,
it’s not a very easy battery to transport or
use in real-life applications.
Electrode
Surface at
which oxidation
or reduction
half-reaction
occurs.
Anode &
Cathode
An Ox Ate a Red Cat
• Anode – Oxidation
• The anode = location for the oxidation
half-reaction.
• Reduction – Cathode
• The cathode = location for the
reduction half-reaction.
Anode / Cathode
• How do you know which electrode is
which?
• Use Table J to predict which
electrode is the anode and which
electrode is the cathode.
Anode
• Anode = Oxidation = Electron Donor
• The anode is the metal that’s higher
in Table J.
Cathode
• Cathode = Reduction = Electron
Acceptor
• The cathode is the metal that’s lower in
Table J.
Zn is above Cu, Zn is anode
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf
Notation for Cells
ZnZn+2Cu+2Cu
Direction of Electron Flow
(wire)
Anode to Cathode
Direction of Positive Ion Flow
(salt bridge)
Anode to Cathode
Positive & Negative Electrode
• Negative electrode is where electrons
originate – here it’s the Zn electrode.
• Positive electrode is electrode that attracts
electrons – here it’s the Cu electrode.
Aqueous Solution
• Solution containing ions of the same
element as the electrode.
• Cu electrode: solution may be Cu(NO3)3
or CuSO4.
• Zn electrode: solution may be Zn(NO3)2
or ZnSO4.
Salt Bridge
• Allows for migration of ions between halfcells.
• Necessary to maintain electrical neutrality.
• Reaction will not proceed without salt
bridge.
A(s) + BX(aq)  B(s) + AX(aq)
• Single replacement rxn occurs during
operation of galvanic cell.
• One electrode will gain mass (B) and one
electrode will dissolve (A).
• The concentration of metal ions will
increase in one solution (making AX) &
decrease in one solution (using up BX).
Half-Reactions
Zn  Zn+2 + 2eCu+2 + 2e-  Cu
_________________________
Zn + Cu+2  Zn+2 + Cu
Which electrode is dissolving?
Which species is getting more
concentrated?
Zn+2
Zn
Zn + Cu+2  Zn+2 + Cu
• Which electrode is gaining mass? Cu
• Which species is getting more dilute? Cu+2
When the reaction reaches
equilibrium
• The voltage goes to 0.
Construct Galvanic Cell with Al & Pb
• Use Table J to identify anode & cathode.
• Draw Cell, put in electrodes & solutions
• Label anode, cathode, direction of electron
flow in wire, direction of positive ion flow in
salt bridge, positive electrode, negative
electrode.
• Negative electrode is where electrons
originate. Positive electrode attracts
electrons.
Electron flow 
wire
Al =
anode
-
Al+3
Positive ion flow 
Salt bridge
& NO3
-1
Pb =
cathode

Pb+2 & NO3-1
What are half-reactions?
Al  Al+3 + 3eAl metal is the electrode – it’s dissolving.
Al+3 ions go into the solution.
Pb+2 + 2e-  Pb
Pb+2 ions are in the solution. They pick
up 2 electrons at the surface of the Pb
electrode & plate out.
Overall Rxn
2(Al  Al+3 + 3e-)
3(Pb+2 + 2e-  Pb)
_____________________________
2Al + 3Pb+2  2Al+3 + 3Pb
2Al + 3Pb+2  2Al+3 + 3Pb
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Al
Which electrode is losing mass?
Which electrode is gaining mass? Pb
What’s happening to the [Al+3]? Increasing
What’s happening to the [Pb+2]? Decreasing
Application: Batteries
Dry Cell
Mercury battery
Fuel Cell
• Converts chemical to
electrical energy
– Use oxygen or other
oxidizing agent
• Constant supply of
fuel (hydrogen or
other hydrocarbon)
Application: Corrosion
Corrosion Prevention
What’s wrong with this picture?
Daniell Cell
• Invented in 1836 by
John Frederic Daniell
• Improved battery
technology (voltaic
pile –problem with
hydrogen bubbling)
• Definition of the volt