Chapter 17: Electrochemistry

1
Chapter 17 Electrochemistry

• Electrochemical Cell Contains electrodes
  dipping into electrolyte solutions. The
  half-reactions occur at two different electrodes
  and a flow of electrons (current) is involved.
• Galvanic (voltaic) cells a spontaneous redox
  reaction occurs and causes a flow of electrons.
• Electrolytic cells an external source of
  electric current is used to drive an otherwise
  nonspontaneous redox reaction.

2
Galvanic Cells

• Consider the spontaneous redox reaction that we
  observed in class (occurring in one beaker)
• Cu(s) 2 Ag(aq) ? Cu2(aq) 2 Ag(s)
• No useful work was done in this process (e.g., no
  usable electric current). Some heat was released
  to the surroundings.
• How can we harness this spontaneous reaction in
  such a way as to yield useful work?
• Answer Separate the half-reactions in two
  containers.

3
Figure 17.2Galvanic Cells
4

• Ex Suppose we have a piece of Zn metal dipping
  into a 1 M solution of ZnSO4 in the left-hand
  compartment and a piece of Cu metal dipping into
  a 1 M solution of CuSO4 in the right-hand
  compartment.
• If we connect the metal electrodes with a wire,
  the following observations are made
• The piece of Zn decreases in mass over time.
• The piece of Cu gains mass over time.
• The concentration of the Zn2(aq) in the
  left-hand compartment increases over time.
• The concentration of Cu2(aq) in the right-hand
  compartment decreases over time.

5

• What is the salt bridge?
• It contains a solution of an electrolyte such as
  KCl in a gel, such that the ions are mobile but
  will not mix rapidly with the other solutions.
• If we place a voltmeter between the two
  electrodes, we obtain a reading of 1.10 V.
• Now, lets analyze more closely what is happening
  in the cell.

6
Figure 17.6 A Zn/Cu Galvanic Cell
7

• At the left-hand electrode (anode half-reaction)
• Zn(s) ? Zn2(aq) 2 e oxidation
• At the right-hand electrode (cathode
  half-reaction)
• Cu2(aq) 2 e ? Cu(s) reduction
• Add them to get the overall reaction
• Zn(s) Cu2(aq) ? Zn2(aq) Cu(s)
• But now the electrons lost by the Zn flow through
  the external circuit from the anode to the
  cathode, i.e., an electric current which we can
  make use of!

8

• Note By convention the anode is always placed
  on the left and the cathode on the right when
  picturing a galvanic cell.
• You should be able to describe the half-reactions
  at each electrode and the direction of movement
  of all species (including electrons and the ions
  of the salt bridge).
• Sometimes a half-reaction will not involve a
  solid substance, e.g., Fe3(aq) e ?
  Fe2(aq)
• In such cases, an inert metal such as Pt is used
  as an electrode to which the wire is attached. A
  example is shown in the following figure.

9
Figure 17.7A Schematic of a Galvanic Cell
10

• Shorthand (Line) Notation for Cells to save
  time we sometimes eliminate the beakers, wires,
  and salt bridge, and simply write the components
  of the cell in order from left to right with
  vertical lines representing phase boundaries.
• For the zinc-copper cell previously described
• Zn(s) Zn2(aq) Cu2(aq) Cu(s)
• Be certain that you dont confuse this with the
  cell reaction! Know how to relate one to the
  other.

11

• A note on nomenclature the terms half-reaction
  and half-cell are synonymous (e.g., Cu2Cu).
• Cell Potential (or EMF or voltage)
• Think of this as the push that sends electrons
  from the anode to the cathode.
• We measure (with a voltmeter) the potential
  difference between the anode and cathode.
• ?Ecell Ecathode Eanode
• The superscript refers to standard-state
  conditions and reversible cell operation.

12

• ?Ecell must be positive in order for a
  spontaneous reaction to occur. Thus, the cathode
  must be at a higher potential than the anode
  (think of electrons rolling downhill).
• Ex for the Zn-Cu cell previously described, ?E
  1.10 V.
• By definition, 1 joule of work is done when a
  charge of 1 coulomb is moved through a potential
  difference of 1 volt.
• Therefore, 1 V 1 J/C
• The electrical work performed is then given by
• welec q ?E

13

• We desire a scale of standard electrode
  potentials E for individual half-reactions.
• Since we can measure only ?E, we must arbitrarily
  assign a potential to one half-reaction, much as
  we assign an atomic mass of exactly 12 to 12C.
• The Standard Hydrogen Electrode (S.H.E.)
• This consists of an inert Pt electrode immersed
  in a 1.00 M H solution and over which H2 gas at
  a partial pressure of 1 atm is bubbled.
• 2 H(1 M) 2 e ? H2(1 atm) E 0.00 V

14
Figure 17.5A Zn/H Galvanic Cell
15

• In the pictured cell, ?E 0.76 V with oxidation
  occurring at the Zn electrode.
• Zn(s) Zn2 (1 M) H (1 M) H2 (1 atm) Pt
  (s)
• ?E 0.76 V Ecath Eanode
• 0.00 V Eanode
• Therefore, Eanode 0.76 V.
• By convention, all half-cell reactions in the
  tables are written as reductions.
• Zn2 (aq) 2 e ? Zn (s) E 0.76 V

16

• If the half-cell containing Cu(s) and Cu2(aq)
  were connected to the S.H.E., the copper
  half-cell would have to be the cathode in order
  to get a positive ?Ecell 0.34 V.
• Thus, the copper half-cell is at a higher
  potential than the S.H.E.
• Cu2(aq) 2 e ? Cu(s) E 0.34 V
• The listing of half-cells and their potentials is
  called the table of electrode potentials or the
  table of reduction potentials or the activity
  series.

17

• Lets look more closely at a portion of the
  activity series.
• Ag(aq) e ? Ag(s) E 0.80
  V
• Cu2(aq) 2 e ? Cu(s)
  0.34 V
• 2 H(aq) 2 e ? H2(g)
  0.00 V
• Ni2(aq) 2 e ? Ni(s)
  0.25 V
• Zn2(aq) 2 e ? Zn(s)
  0.76 V
• What is the significance of these values?
• 1) The more positive the E value, the more
  favorable is a reduction half-reaction (e.g., Ag
  will be reduced more spontaneously than Cu2).

18

• In a galvanic cell, the electrode with the higher
  E value will be the cathode.
• Ex In a cell using the Ni2Ni (0.25 V) and
  AgAg (0.80 V) half-reactions, Ag will be the
  cathode
• and the spontaneous reaction will be
• Ni(s) 2 Ag(aq) ? Ni2(aq) 2 Ag(s)
• ?E 0.80 (0.25) 1.05 V
• Important Notice that the sign of the potential
  of the anode is used as found in the table (not
  reversed). This is because we subtracted the
  anode potential from the cathode potential.

19

• 3) The table also ranks substances according to
  relative oxidizing or reducing power.
• Substances on the left accept electrons and are
  oxidizing agents. Substances on the right give
  up electrons and are reducing agents.
• A higher E means that the substance on the left
  is more likely to become reduced, i.e., is a
  stronger oxidizing agent. In the last example,
  Ag is a stronger oxidizing agent than Ni2.
• A lower E means that the substance on the right
  is more likely to become oxidized, i.e., is a
  stronger reducing agent.

20

• Examples of good oxidizing agents
• F2(g) 2 e ? 2 F(aq)
  E 2.87 V
• H2O2(aq) 2H(aq) 2 e ? 2 H2O(l)
  1.79 V
• MnO4(aq) 4 H(aq) 3 e ? MnO2(s) 2
  H2O(l) 1.68 V
• Examples of good reducing agents
• Li(aq) e ? Li(s)
  E 3.04 V
• K(aq) e ? K(s)
  2.92 V
• Ca2(aq) 2 e ? Ca(s)
  2.76 V

21

• Displacement Reactions
• Ex Zn(s) Cu2(aq) ? Zn2(aq) Cu(s)
• We say that zinc has displaced copper(II) ion
  from the solution.
• The stronger reducing agent will enter the
  solution (dissolve), and the stronger oxidizing
  agent will leave the solution (be displaced).
• Spontaneous redox reactions always proceed from
  the stronger oxidizing and reducing agent to the
  weaker oxidizing and reducing agent.

22

• Ex Fe2(aq) 2 e ? Fe(s) E
  0.41 V
• Pb2(aq) 2 e ? Pb(s) E
  0.13 V
• Which metal will displace the other from
  solution?
• Is Fe2(aq) or Pb2(aq) the stronger oxidizing
  agent?
• Is Fe(s) or Pb(s) the stronger reducing agent?
• The spontaneous reaction should be
• Fe(s) Pb2(aq) ? Fe2(aq) Pb(s)
• Therefore, Fe displaces Pb2 from the solution.