Ch 21. Electrochemistry

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Ch 21. Electrochemistry

• Brady Senese 4th Ed.

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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

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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)
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• 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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Recognizing The Cathode

• Given the cell reaction the cathode is the
  substance that is reduced (gains electrons)
• The anode is the substance that creates electrons
• LEO the lion says GER
• the loss of electrons is oxidation and
• the gain of electrons is reduction
• Spontaneously the cathode has the most Ered of
  the choices and the anode has the most Ered.

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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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• 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

• 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

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• 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

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Learning Check

• Write the half reactions and identify the cathode
  and anode
• Cu2(aq) Zn(s) Cu(s) Zn2(aq)
• HCl(aq) Zn(s) H2(g) ZnCl2(aq)

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Electrical Potential

• Every substance has the potential to gain
  electrons or be reduced in oxidation state
• The relative ease of gaining electrons is termed
  the reduction potential and is symbolized Ered
• If the matter being observed is in standard state
  then E is termed the standard reduction potential
  and is symbolized as E0red

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Standard Reduction Potentials (E0red)

• E0red are tabulated for nearly every known
  substance.
• A high value of E0red (E0red 0) means that the
  substance is easily reduced
• E0red is a relative number arbitrarily
  determined.
• All substances are compared to H which has a
  E0red of 0.00 V.

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• 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

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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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Cell Potentials

• The standard cell potential is calculated as
• E0cell E0cath-E0anode
• If the cell is non-standard
• EcellEcath-Eanode
• In spontaneous redox reactions the cathode
  portion of the reaction has a higher reduction
  potential than that of the anode (EcathEanode)

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Predicting The Cells In A Spontaneous Reaction.
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Learning Check

• Calculate E0cell. Which are spontaneous
• Cu(s) Ag(aq) Cu2(aq) Ag(s)
• Pb Cu2 Pb2 Cu

.799-.337V0.462V
.337-(-.126)V0.463V
Cu2/Cu 0.337V Ag/Ag 0.799V
Pb2/Pb -0.126V
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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 96485 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 expressions are forms 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

at 25 oC
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• 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

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Learning Check

• Calculate G0 in kJ. Which are spontaneous under
  standard conditions
• Cu(s) Ag(aq) Cu2(aq) Ag(s)
• Pb Cu2 Pb2 Cu

G0 -2mol x 96485C/mol x 0.462J/C -89.1 kJ
G02mol x 96485C/mol x 0.463J/C -89.3 kJ
Cu2/Cu 0.337V Ag/Ag 0.799V Pb2/Pb -0.126V
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Learning Check

• Calculate K for the following cells.
• Cu(s) Ag(aq) Cu2(aq) Ag(s)
• Pb Cu2 Pb2 H2O

E00.462V
E00.463V
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Learning Check Calculate Ecell
Al3/Al -1.662V Zn2/Zn -0.763V

• AlAl3(aq) (0.5M)Zn2(aq) (0.2M)Zn
• AlAl3(aq) (0.5M)Zn2(aq) (1M)Zn

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Learning Check (Cont.)

• AlAl3(aq) (0.5M)Al3(aq) (0.2M)Al
• AlAl3(aq) (0.5M 250 C )Al3(aq) (0.5M 500 C
  )Al

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Learning Check Find The Unknown

• CrCr3(M) Ni2(0.5M)Ni
• Ecell .494 V
• PtH2(1atm) H( pH)H(1M)H2(1 atm)Pt
• Ecell .000571 V

0.353Mx
0.978Mx pH0.00965
H/H2 0.00V Cr3/Cr -0.744V Ni2/Ni -0.25V Cu2/Cu
0.337V
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Learning Check

• Consider the following reaction. Is it favored
  at high or low temperatures
• PtH2(1 atm)H(pH4)Cu2(0.1M)Cu

When Qtemperature
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Batteries

• 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

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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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• 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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The Electrolysis Of Aqueous K2SO4
The products of the electrolysis are H2 and O2
gas not the expected products solid K and
S2O82-. Why
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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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