1. Electrochemistry Part III: Reduction Potentials
Jespersen Chap. 20 Sec 2 & 3
Dr. C. Yau
Fall 2014 1 1 1 1

2. 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.
(Often "red" is omitted and given as Eo)
Units are V (volts)
1V=1J/C where C is coulomb, a unit of charge. 2 2

3. Standard Reduction Potentials (E0red)
E0 are tabulated for nearly every known substance.
A high value of E0 (E0 > 0) means that the substance is easily reduced
E0 is a relative number, arbitrarily determined.
Figure 19.5 The hydrogen electrode. The half-reaction is 2H+(aq) + 2e– ↔ H2(g). Table 19.1 gives E0red
2H e−  H2
All substances are compared to H+, which has a E0 of 0.00 V. 3 3

4. Table 20.1 p.928 4

5. Section of Reduction Potential Table
Half -Reaction
Eo (volts)

6. Cell Potentials The standard cell potential is calculated as:
E0cell= E0cath-E0anode
If the cell is non-standard:
Ecell=Ecath-Eanode
In spontaneous redox reactions, the cathode portion of the reaction has a higher reduction potential than that of the anode (Ecath>Eanode)
So Ecell is positive.
ChemFAQs: How can we compute the cell potential from reduction and oxidation potentials at the electrodes?
How can I predict whether a redox reaction occurs spontaneously as written, using reduction potentials?
How can I predict the cell potential of a galvanic cell?
How can I predict whether a redox reaction is spontaneous, using the calculated cell potential? 6 6

7. Making Use of Standard Reduction Potentials Table (Table 20.1 p. 928)
Example 20.3 p. 931
What spontaneous reaction occurs if Cl2 and Br2 are added to a solution that contains both Cl- and Br-? Assume standard states.
What are the possibilities?
2Cl− + Br2  Cl2 + 2Br−
or 2Br− + Cl2  Br2 + 2Cl−
Cl2 (g) + 2e Cl- (aq) Eo = +1.36
Br2 (g) + 2e Br- (aq) Eo = +1.07
What is Eo for Cl2 + Br-?
what is Eo for Br2 + Cl-? 7 7

8. Pract Exer 20.7 p. 932
Predict whether the following reaction is spontaneous if all the ions are 1.0 M at 25oC. If it is not, write the equation that represents a spontaneous reaction.
Ni2+(aq) + 2 Fe2+(aq) Ni(s) + 2Fe3+ (aq)
Which reaction is at the cathode?
Do Pract Exer 5, 6 p. 932 8 8

9. Helpful Tips For a reaction to be spontaneous,
the half-rxn with the larger (more positive) standard reduction potential Eo takes place as REDUCTION…
the other half-rxn will be reversed and occur as oxidation (change the sign of Eo)
and ADD the two Eo together.
This is the same as saying "to subtract one Eo from the other". 9 9

10. Example 20.4 p. 932
A typical cell of a lead storage battery of the type used to start automobiles is constructed using electrodes made of lead and lead(IV) oxide PbO2 and with sulfuric acid as the electrolyte. The half-rxns & their standard reduction potentials in this system are shown below. What is the cell rxn and what is the standard potential of the cell?
PbSO4 (s) + H+(aq) + 2 e Pb(s) + HSO4 (aq) Eo= -0.36V
PbO2(s)+3H+(aq)+HSO4(aq)+ 2e PbSO4 (s)+2H2O (l)
Eo = 1.69V 10 10

11. Note: We do not multiply the Eo by two. Do Pract Exer 8, 9, 10 p. 934
Example p. 933
A galvanic cell is constructed by placing 1.0 M aluminum nitrate in one beaker with an Al electrode and a second beaker with 1.0 M copper(II) nitrate and a Cu electrode. When a saltbridge is inserted to connect the 2 beakers, what cell potential do we expect to measure between the Cu and Al electrodes? Which electrode is the anode, and what is the spontaneous chemical reaction?
Note: We do not multiply the Eo by two.
Do Pract Exer 8, 9, 10 p. 934 11 11

12. Example 20.6 p. 935
Determine whether the following rxns, at standard state, are spontaneous as written. If not, write the eqn for the rxn that is.
1) Cu (s) + 2H+ (aq)  Cu2+ (aq) + H2 (g)
2) 3Cu(s)+ 2NO3(aq)+ 8H+(aq)  Cu2+(aq)+ 2NO(g) + 4 H2O(l)
Do Pract Exer 11, 12 p. 936 12 12

13. Redox Trends in Periodic Table
Be sure to review Sec 9.7 p. 381
Reactivities of Elements & Electronegativity
Reactivities of metals and nonmetals are related to electronegativity.
Metals react by losing electrons to form cations. Is this oxidation or reduction?
Write a general eqn for metals reacting.
The most “reactive metal” or “active metal” are ones that can lose electrons most readily (the most electronegative or least?)
What is the general trend of ease of oxidation? General trend of ease of reduction? 13 13

14. Redox Trends in Periodic Table
Nonmetals react by gaining electrons to form anions. Is this oxidation or reduction? Write a general equation of a nonmetals reacting, using Group VIIA as an example. The most reactive nonmetals are ones that can gain electrons most readily (the most electronegative or least?) What is the general trend of ease of reduction? General trend of ease of oxidation? 14 14

15. What do metals tend to do?
Metals have tendencies to...
be oxidized.
be reduced.
Metals tend to be
oxidizing agents.
reducing agents.
M  M+ 15 15

16. What do nonmetals tend to do?
Nonmetals with high electronegativities have strong tendencies to...
be oxidized.
be reduced.
Which two elements are the most powerful oxidizing agent?
Ans. F2 followed by O2
X  X- 16 16

17. General Trends most active nonmetals Periodic Table
(as oxidizing agents)
easily reduced
most active metals
(as reducing agents)
easily oxidized
Arrow shows increase in “metallic character.”
Periodic Table 17

18. Protection Against Corrosion
What exactly is "corrosion" of a metal?
We often talk about a metal being "oxidized."
When an iron nail rusts, what exactly is it doing? What is rust?
A simplified view of the rusting process:
Fe (s) Fe2O3 (s) oxidation or reduction?
Why does iron rusts more quickly in contact with a metal such as Cu?
It rusts more slowly in contact with other metals…such as what? 18

19. Facts about Rusting of Iron
does not rust in dry air (moisture must be present)
does not rust in air-free water (O2 must be present)
loss of Fe and deposit of rust often occur at different places on the same object
rusts more quickly at low pH
rusts more quickly in contact with ionic solutions
rusts more quickly in contact with a less active metal such as Cu
rusts more slowly in contact with a more active metal such as Zn or Mg. 19

20. Fe acts as both anode & cathode.
Salt water acts as salt bridge.
H+ is a catalyst & speeds up process. 20

21. Elimination of Corrosive Factors
Keep dry
Rinse off road salts (remove electrolytes)
Protect with paint (keep oxygen, moisture, salts and acids out)
What about contact with other metals? 21

22. Let's examine cell potentials of some metals:
Cu2+ + 2e-  Cu Eo = V
Sn2+ + 2e-  Sn Eo = V
Fe2+ + 2e-  Fe Eo = V
Zn2+ + 2e-  Zn Eo = V
Mg2+ + 2e-  Mg Eo = V
"Active metal" refers to metals that oxidizes easily, ones that do not reduce easily.
Metals less active than Fe can oxidize it.
Fe in contact with Cu, Fe is the anode, changing to Fe2+.
less "active"
more "active" 22

23. O2 is oxidizing the Mg instead of the Fe.
Cathodic Protection
Cu2+ + 2e-  Cu Eo = V
Sn2+ + 2e-  Sn Eo = V
Fe2+ + 2e-  Fe Eo = V
Zn2+ + 2e-  Zn Eo = V
Mg2+ + 2e-  Mg Eo = V
If we place a more active metal, such as Mg, in contact with Fe, then Fe is the cathode and no Fe is lost as Fe2+.
O2 is oxidizing the Mg instead of the Fe.
e- from Mg (anode) travels thru Fe (cathode) to oxygen which is reduced to OH- 23

24. An active metal such as Mg or Al, is connected to underground Fe pipes to prevent corrosion. The active metal is "sacrificed" instead of the Fe. It has to be replaced from time to time. 24

25. What about other metals?
How do you protect other metals?
Cu2+ + 2e-  Cu Eo = V
Sn2+ + 2e-  Sn Eo = V
Fe2+ + 2e-  Fe Eo = V
Zn2+ + 2e-  Zn Eo = V
Mg2+ + 2e-  Mg Eo = V
Just remember you would use a more active metal.
What does "active" mean?
“Sacrificial” metal is the more active metal. 25