Module 1 - Electrochemistry (Part 1)
Module 1 - Electrochemistry (Part 1)
Module 1 - Electrochemistry (Part 1)
MODULE 1
ELECTROCHEMISTRY
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OVERVIEW:
In this module, you will study the theoretical aspects and other applications
of electrochemical reactions. The discussions begin with a review of redox reactions
and where you will learn how to balance equations describing these processes.
(Lesson 1.1)
Next, you will examine the essentials of voltaic cells. Here you will learn to
determine the standard reduction potentials based on the standard hydrogen
electrode reference and use them to calculate the emf of a cell and hence the
spontaneity of a cell reaction. (Lesson 1.2 – 1.3)
Consequently, you will learn to calculate the emf of a cell under nonstandard
conditions using the Nernst equation. (Lesson 1.4)
Then you will see that a significant relationship exists between an emf of a
cell, the change in the standard Gibb’s free energy and the equilibrium constant for
the cell reaction. (Lesson 1.5)
OBJECTIVES:
At the end of this module, you should be able to:
1. explain the concept of electrochemistry;
2. use the terminology of electrochemistry (terms such as “cell,” “electrode,”
“cathode,” “anode”);
3. differentiate the types of electrochemical cells;
4. recognize oxidation and reduction half-reactions, and know at which
electrode each occurs;
5. describe the construction of simple voltaic cells from half-cells and a salt
bridge, and understand the function of each component;
6. write half-reactions and overall cell reactions for voltaic cells;
7. compare various voltaic cells to determine the relative strengths of
oxidizing and reducing agents;
8. interpret standard reduction potentials;
9. use standard reduction potentials, E0, to calculate the potential of a
standard voltaic cell, E0 cell;
10. use standard reduction potentials to identify the cathode and the anode
in a standard cell;
11. use standard reduction potentials to predict the spontaneity of a redox
reaction;
12. use standard reduction potentials to identify oxidizing and reducing
agents in a cell or in a redox reaction;
13. Use the Nernst equation to relate electrode potentials and cell potentials
to different concentrations and partial pressures;
14. Relate the standard cell potential (E0 cell) to the standard Gibbs free
energy change (ΔG0) and the equilibrium constant (K);
15. Distinguish between primary and secondary voltaic cells
16. distinguish between primary and secondary voltaic cells;
17. describe the compositions and reactions of some useful primary and
secondary cells (batteries);
18. describe the electrochemical processes involved in discharging and
recharging a lead storage (automobile) battery;
19. Describe some corrosion processes and some methods for preventing
corrosion;
20. predict electrode products of a given electrolysis process;
21. write half-reactions and overall cell reactions for electrolysis processes;
22. explain Faraday’s Law of electrolysis;
23. solve problems on Faraday’s Law as applied to electroplating;
24. describe the refining and plating of metals by electrolytic methods; and
25. give some applications of electrolysis.
Prepared by: E. V. SILFAVAN•C.D.SACDALAN•M.S.P.RODIL•M.C.T.CABILDO•E.S.CAPINDING•G.A.ERGINO
Chemistry Dept. / COS, TUP Manila
SY: 2020 - 2021
3
CHEM ENG
MODULE 1: ELECTROCHEMISTRY
2. Those which use electricity to cause the reactions to occur. The process
takes place in an electrolytic cell wherein the positive electrode is the
anode and the negative electrode is the cathode.
Whether the cell is voltaic or electrolytic, oxidation always occurs at the
anode, while the reduction always occur at the cathode.
ions, the positive ions towards the CuSO4 and the negative ions towards the
ZnSO4.
6. The over-all reaction is as follows:
Zn(s) → Zn+2(aq) + 2e-
Cu+2(aq) + 2e- → Cu(s)
Over-all reaction: Zn(s) + Cu+2(aq) → Zn+2(aq) + Cu(s)
The electrical current flows from the negative electrode (anode) to the
positive electrode (cathode) because there is a difference in the electrical potential
between the two electrodes. The difference in electrical potential is called an
electromotive force, emf (cell potential or cell voltage), which is measured by a
voltmeter.
The standard potential values are intensive properties, hence are not
affected by the change of stoichiometric coefficients during balancing.
For example: Suppose the reaction is multiplied by 2, it becomes:
2 Zn(s) → 2 Zn+2(aq) + 4e- E0 = + 0.76
The E0 values are useful for evaluating redox reactions. Substances with
high positive reduction potential are strong oxidizing agents (greater tendency to be
reduced) and substances with high positive oxidation potential or low reduction
potentials are strong reducing agents (greater tendency to be oxidized). As can be
seen in the table, the substances are arranged from least active metal to most active
Prepared by: E. V. SILFAVAN•C.D.SACDALAN•M.S.P.RODIL•M.C.T.CABILDO•E.S.CAPINDING•G.A.ERGINO
Chemistry Dept. / COS, TUP Manila
SY: 2020 - 2021
7
CHEM ENG
MODULE 1: ELECTROCHEMISTRY
metal. Therefore, less active metal has greater tendency to be reduced, and more
active metal has greater tendency to be oxidized.
Table 1.1. Standard Reduction Potentials in Aqueous Solution (1M) at 250C
E0
Half Cell Equation
Volts
F2(g) + 2e- → 2 F- (aq) +2.87
H2O2 (aq) + 2H+ (aq) + 2e- → 2 H2O +1.77
PbO2(s) + 4H+(aq) + SO4-2 (aq) + 2e- → PbSO4(s) + 2H2O +1.70
MnO4-(aq) +8H+(aq) + 5e- → Mn+2(aq) + 4H2O +1.51
Stronger Weaker
oxidizing Au3+ (aq) + 3e- → Au(s) +1.50 reducing
agent Cl2(g) +2e- →2Cl- (aq) +1.33 agent
MnO2 (aq) +4H+(aq) + 2e- → Mn+2(aq) + 2H2O +1.23
O2(g) + 4H+(aq) + 4e- →2 H2O +1.23
Br2(g) +2e- →2Br- (aq) +1.07
NO3- (aq) + 4H+(aq)+ 3e- →NO(g) + 2H2O +0.96
Ag1+ (aq) + e- → Ag(s) +0.80
Fe3+ (aq) + e- → Fe2+(aq) +0.77
O2(g) + 2H+(aq) + 2e- → H2O2(aq) +0.68
MnO4-(aq) +2H2O + 3e- → MnO2(s) + 4OH-(aq) +0.59
I2(g) +2e- →2I- (aq) +0.53
O2(g) + 2H2O + 4e- →4OH-(aq) +0.40
Cu2+ (aq) + 2e- → Cu(s) +0.34
SO42-(aq) +4H+(aq) +2e- →SO2(g) +2H2O +0.20
2H+ (aq) + 2e- → H2(g) 0.00
Pb2+ (aq) + 2e- → Pb(s) -0.13
Sn2+ (aq) + 2e- → Sn(s) -0.14
Ni2+ (aq) + 2e- → Ni(s) -0.25
Co2+ (aq) + 2e- → Co(s) -0.28
PbSO4(s) +2e- →Pb(s) +SO42-(aq) -0.31
Cd2+ (aq) + 2e- → Cd(s) -0.40
Weaker Fe2+ (aq) + 2e- → Fe(s) -0.44 Stronger
oxidizing Cr3+ (aq) + 3e- → Cr(s) -0.74 reducing
agent Zn2+ (aq) + 2e- → Zn(s) -0.76 agent
2H2O + 2e- → H2(g) +2OH-(aq) -0.83
Mn2+ (aq) + 2e- → Mn(s) -1.18
Al3+ (aq) + 3e- → Al(s) -1.66
Mg2+ (aq) + 2e- →Mg(s) -2.37
Na+ (aq) + e- → Na(s) -2.71
Ca2+ (aq) + 2e- → Ca(s) -2.87
Ba2+ (aq) + 2e- → Ba(s) -2.90
The cell potential values (Ecell) determines whether the proposed half-cell
reactions will occur or not. If the value is positive, the reaction is spontaneous and
will occur as it is written. If the value is negative, the reaction is non-spontaneous
and it will not occur as it is written, however, if the half-cell reactions are reversed,
i.e. the oxidation becomes, the reduction, the reaction will occur (if the E 0cell of the
non-spontaneous reaction is measured with a voltmeter, the reading is negative,
hence the electrode should be reversed).
The reaction is spontaneous when more reactive metal is placed in the anode
and less reactive metal is placed in the cathode. In galvanic cell, electron will always
flow from more to less reactive metal.
➢ Cell Representations
An electrochemical cell can be represented by
1. Using overall cell reaction:
Zn(s) + Cu+2(aq) → Zn+2(aq) + Cu(s)
2. Using the electrode reactions:
Zn(s) /Zn+2 // Cu+2 / Cu(s)
The “/” represents phase boundary and “//” represents a salt bridge.
EXAMPLES
Solution:
3. Al / Al+3 // Mg+2 / Mg
4. Ca / Ca+2 // Cr+3 / Cr
5. Mg / Mg+2 // Fe+2 / Fe
II. In the table of standard reduction potentials, locate the half-reactions for the
reductions of the following metal ions to the metal: Sn+2(aq), Au+(aq), Zn2+(aq),
Co2+(aq), Ag+(aq), and Cu2+(aq). Among the metal ions and metals that make up
these half-reactions:
a) Which metal ion is the weakest oxidizing agent?
b) Which metal ion is the strongest oxidizing agent?
c) Which metal is the strongest reducing agent?
d) Which metal is the weakest reducing agent?
e) Will Sn(s) reduce Cu2+(aq) to Cu(s)?
f) Will Ag(s) reduce Co2+(aq) to Co(s)?
g) Which metal ions on the list can be reduced by Sn(s)?
h) What metals can be oxidized by Ag+(aq)?