Chapter 21: Electrochemistry

1
Chapter 21 Electrochemistry

• Batteries serve as power sources for all types of
  gadgets
• The energy in a battery comes from a spontaneous
  redox reaction where the electron transfer is
  forced to take place through a wire
• The apparatus that provides electricity in this
  way is called a galvanic or voltaic cell

2

A galvanic cell. The cell consists of two
half-cells where the oxidation and reduction
half-reactions take place. The salt bridge is
required for electrical neutrality. The overall
cell reaction is 2Ag(aq)Cu(s)?
2Ag(s)Cu2(aq)
3

• Cell reactions are obtained by adding the
  half-reactions
• Half-reactions are balanced using the
  ion-electron method (see Section 6.2)
• The electrodes are assigned the name anode or
  cathode
• Reduction (electron gain) occurs at the cathode
• Electrons appear as reactants in the
  half-reaction
• Oxidation (electron loss) occurs at the anode
• Electrons appear as products in the half-reaction

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Changes that take place at the anode and cathode
of a copper-silver galvanic cell. At the anode,
Cu2 ions enter the solution when copper atoms
are oxidized. At the cathode, Ag ions leave
solution and become silver atoms.

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• There are two types of electrical conduction in a
  galvanic cell
• Metallic conduction occurs when electrons move
  through the wires
• Electrolytic conduction occurs through the liquid
  by movement of ions, not electrons
• The movement of ions through the salt bridge and
  in solution is required for charge neutrality
• Cations move in the general direction of the
  cathode
• Anions move in the general direction of the anode

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• The anode has negative polarity because the
  electrons left behind by the Cu2 ions give it a
  slightly negative charge
• The cathode has positive polarity because of the
  Ag ions joining the electrode give it a
  slightly positive charge
• For convenience, a standard cell notation has
  been developed by chemists
• Anode half-cell is specified on the left
• Cathode half-cell is specified on the right
• Phase boundaries are indicated using
• The salt bridge separates the anode and cathode
  and is indicated using

7

• The cell diagram for the copper-silver galvanic
  cell is
• Cu(s)Cu2(aq)Ag(aq)Ag(s)
• (anode) (cathode)
• Galvanic cells can push electrons through a wire
• The magnitude of this ability is expressed as a
  potential
• The maximum potential a given cell can generate
  is called the cell potential, Ecell

8

• The cell potential depends on the temperature and
  composition
• The standard cell potential, Eocell, is the cell
  potential measured at 298 K (25oC) with all ion
  concentration 1.00 M
• Standard cell potentials are rarely more than a
  few volts
• Eocell for the copper-silver galvanic cell is
  0.46 V
• Eocell for a single cell in a car battery is
  about 2 V

9

• The tendency for a species to gain electrons and
  be reduced is its reduction potential
• When measured at standard condition, it is called
  the standard reduction potential, Eo
• When two half-cells are are connected
• The one with the larger reduction potential will
  acquire electrons and undergo reduction
• The half-cell with the lower reduction potential
  will give up electrons and undergo oxidation

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• The difference in the two standard reduction
  potentials gives the standard cell potential
• It is not possible to measure the reduction
  potential of an isolated half-cell
• A reference electrode, called the standard
  hydrogen electrode, has been assigned the
  potential of exactly 0 V

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The standard hydrogen electrode. Hydrogen gas at
1 atm is passed over finely divided platinum. The
solution contains 1.00 M hydrogen ion. The
reduction potential is exactly 0 V at 298 K
(25oC).
12

• Using a hydrogen half-cell, other reduction
  potentials can be measured

A galvanic cell comprised of copper and hydrogen
half-cells. The reaction is Cu2(aq)H2(g)?
Cu(s)2H(aq)
Cell notation Pt(s), H2(g)H(aq)Cu2(aq)Cu(s)
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• Table 21.1 and Appendix C lists the standard
  reduction potentials for a number of species
• Calculated cell potentials can be used to
  identify spontaneous reactions
• For any pair of reduction potentials ordered most
  reactive to least reactive
• The higher (more positive) reduction potential
  will occur as a reduction
• The lower (more negative) reduction potential
  will be reversed and occur as an oxidation

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• In a galvanic cell, the calculated cell potential
  for the spontaneous reaction is always positive
• If the calculated cell potential is negative, the
  cell is spontaneous in the reverse direction
• The free energy change for a system can also be
  used to predict if a reaction is spontaneous
• Free energy changes and cell potentials are
  related

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• The maximum useful work that can be obtained from
  a reaction is
• In electrical systems, work is supplied by the
  current as it is pushed along by the potential of
  the cell
• maximum work nF
• n number of moles of electrons transferred
• F Faraday constant 96,485 C/mol e
• cell potential in volts

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• Equating
• The cell potential can be related to the
  equilibrium constant K
• Cell potentials depend on concentrations

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• The last expression is a form of the Nernst
  equation which relates ion concentrations to the
  cell potential
• Use molar concentrations (M) for ions and partial
  pressures of gases in atmospheres when
  calculating Q

18

• Example In a certain zinc-copper cell,
• Zn(s)Cu2(aq)?Zn2(aq)Cu(s)
• the ion concentrations are Cu20.0100 M and
  Zn21.0 M. What is the cell potential at 298
  K?
• From Table 21.1
• For this two electron change at 298 K

19

• Galvanic cells, commonly called batteries, can
  be classified as either primary or secondary
  cells
• Primary cells are not designed to be recharged
• Secondary cells are able to be recharged
• A battery is usually a collection of cells
  connected in series
• When connected in series, the voltage of each
  cell is added to provide the total voltage of the
  battery

20

• Lead Storage Battery

Three 2 V cells connected in series. The cell
voltages are added. The total voltage provided is
6 V.
A 12 V lead storage battery, like those used in
automotive applications, consist of six cells
like the one shown here. Heavy duty plates are
used so the current produced is sufficient to
start a car.
21

• Fuel cells are electrochemical cells in which
  electrode reactants are continuously supplied

Early hydrogen-oxygen fuel cell. The net reaction
is the conversion of hydrogen and oxygen into
water. Typically, a gas or diesel engine is only
about 20 30 efficient. Fuel cells have
efficiencies as high as 75.
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• Electricity can be used to make nonspontaneous
  redox reactions to occur
• The process is called electrolysis
• Electrolysis occurs in an electrolysis or
  electrolytic cell
• These cells require a source of direct current,
  possibly one of the batteries just discussed, to
  provide electrical energy

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Electrolysis of molten sodium chloride. The
passage of an electric current through molten
sodium chloride decomposes the material into
molten sodium and chlorine gas. The products must
be kept separated because they react on contact
to re-form NaCl.

24

A microscopic view of the changes at the anode in
the electrolysis of molten NaCl. The positive
charge of the electrode attracts Cl- ions. At the
surface of the electrode electrons are pulled
from the ions yielding neutral Cl atoms with
combine to form Cl2 molecules that rise to the
surface as a gas.
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• Sodium is a liquid at the melting point of sodium
  chloride (801 oC)
• Reaction at the cathode
• Na(l) e- ? Na(l)
• Reaction at the anode
• 2Cl-(l) ? Cl2(g) 2e-
• Cell reaction
• 2Na(l) 2Cl-(l) ? 2Na(l) Cl2(g)
• The electrode polarities are reversed in
  electrolytic cells relative to those in galvanic
  cells

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• In both types of cells oxidation occurs at the
  anode and reduction occurs at the cathode
• Electrolysis in aqueous solutions often involves
  water molecules
• This is often unintended and called a competing
  reaction

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Electrolysis of an aqueous solution of potassium
sulfate. The products of the electrolysis are H2
and O2 gas, not the expected products solid K
and S2O82-.
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• The competing reactions at the cathode are
• K(aq)e- ? K(s) EoK
  -2.92 V
• 2H2O(l)2e- ? H2(g)2OH-(aq) EoH2O -0.83 V
• Water has a less negative reduction potential
  than the potassium ion so it it easier to reduce
• When electrolysis is performed, the more easily
  reduced substance is reduced and H2 is observed
  to form
• A similar situation occurs at the anode

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• The competing reactions at the anode are
• The standard cell potential tells us that S2O82-
  is more easily reduced than O2
• This means that the product SO42- is harder to
  oxidize than water
• When electrolysis is performed the more easily
  oxidized substance is oxidized and O2 is observed
  to form at the anode

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• The net change is then
• 2H2O?2H2(g)O2(g)
• This is the sort of analysis that needs to be
  done to anticipate the products of electrolysis
• Sometimes the predicted products are still not
  obtained
• The reasons are beyond the scope of this book but
  serve as a reminder to be cautious about
  predicting the products of electrolysis solely on
  the basis of reduction potentials

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• Michael Faraday discovered that the amount of
  chemical change is directly proportional to the
  amount of electrical charge passed through an
  electrolysis cell
• For example, the equation Cu2(aq)?Cu(s)2e-
• says that to deposit 1 mol of Cu, 2 mol of
  electrons are required
• The SI unit of electrical current is the ampere
  (A) and the SI unit of charge is the coulomb (C)

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• A coulomb is the amount of charge that passes a
  given point in a wire when an electrical current
  of 1 ampere flows for 1 second
• The charge carried by 1 mol of electrons is the
  basis for the Faraday constant (96,485 C/mol e)
• This provides a way to measure the amount of
  chemical change in the laboratory

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• Example What current must be supplied to deposit
  3.00 g Au from a solution of AuCl3 in 200.0 s?
• Electroplating is a procedure in which
  electrolysis is used to apply a thin coating of
  one metal over another
• There are many industrial applications

34

Apparatus for electroplating silver. Silver
dissolves at the anode where it is oxidized to
Ag. Silver is deposited at the cathode (the
fork) where Ag is reduced.
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• Aluminum does not occur as a free metal in nature
  because it is so reactive
• It is obtained from the electrolysis of Al2O3 at
  high temperatures
• Al3 3e- ? Al(l)
  (cathode)
• 2O2- ? O2(g) 2e-
  (anode)
• 4 Al3 6O2- ? 4Al(l) 3 O2(g) (net cell)
• Production of Al requires lots of electricity and
  heat
• Recycling of Al costs only a small fraction of
  what it costs to produce Al from the ore bauxite