ng=0 width=100%>
Yahoo! is not affiliated with the authors of this page or responsible for its content.
419 Studying Electrochemistry and
Establishing the Relative Reactivities
of a Series of Metals
prepared by J. N. Spencer, Franklin and Marshall College,
and H. A. Neidig, Lebanon Valley College
Purpose of the Experiment
Establish the relative electromotive forces, or reactivities, of a series of
metals by observing the direction of electron exchange when pairs of the
metals, each suspended in a solution of its ions, are connected to form a
closed electrical system, or electrochemical cell.
Background Information
Chemical reactions that proceed by electron transfer
from one reactant to another are called oxidationre-
duction reactions, or redox reactions. The reactant
that gives up electrons is said to be oxidized, and the
reactant that gains electrons is reduced.
For example, when we place a strip of zinc (Zn)
metal in an aqueous copper(II) sulfate solution
(CuSO
4
), we immediately begin to see evidence that a
chemical reaction is occurring. The Zn begins to dis-
solve, and a deposit of copper (Cu) metal appears on
the Zn strip. We can write an overall equation for this
reaction once we understand each of the two compo-
nent processes separately.
In this reaction, Zn atoms lose electrons to be-
come zinc ions (Zn
2+
) as the Zn strip dissolves. We
can express the dissolution of the Zn as shown in
Equation 1, where e represents an electron.
Zn(s) Zn
2+
(aq) + 2 e (Eq. 1)
This equation also demonstrates that Zn is the reac-
tant being oxidized. At the same time, copper(II) ions
(Cu
2+
) in the solution gain electrons to become Cu at-
oms. We can express the formation of Cu, as shown in
Equation 2.
Cu
2+
(aq) + 2 e
Cu(s)
(Eq. 2)
The Cu
2+
ion is the reactant being reduced.
We can write the overall equation for this oxida-
tionreduction reaction by combining Equations 1 and 2.
Zn(s) + Cu
2+
(aq) Zn
2+
(aq) + Cu(s)
(Eq. 3)
The electrons appearing in Equations 1 and 2 cancel
each other when we add the two equations. In order for
a redox reaction to occur, the number of electrons pro-
vided by the oxidized reactant must be equal to the
number of electrons accepted by the reduced reactant.
There can be no net gain or loss of electrons in the
overall equation.
E L E C
419
m o d u l a r l a b o r a t o r y p r o g r a m i n c h e m i s t r y
program editor : H. A. Neidig
Copyright © 1993 by Chemical Education Resources, Inc., P.O. Box 357, 220 S. Railroad, Palmyra, Pennsylvania 17078
No part of this laboratory program may be reproduced or transmitted in any form or by any means, electronic or mechanical, including photo-
copying, recording, or any information storage and retrieval system, without permission in writing from the publisher. Printed in the United
States of America
97
96
95
94
93 12
11
10
9
8
7
6
5
4
3
2
1 We refer to the reactions in Equations 1 and 2 as
half-reactions, because each reaction constitutes half
of the overall chemical reaction. We refer to the combi-
nation of the oxidized and reduced substances in a half
reaction as a redox couple. Thus, as shown in Equa-
tions 1 and 2, Cu/Cu
2+
ion and Zn/Zn
2+
ion are exam-
ples of redox couples in their respective half-reactions.
If we place a Cu strip in an aqueous solution of a
zinc(II) salt such as zinc(II) sulfate (ZnSO
4
), we see no
visible evidence of a reaction. This experiment shows
that indeed the reverse reaction of Equation 3 does not
occur. Thus, we conclude that it is easier for Zn atoms
to give up electrons to Cu
2+
ions than it is for Cu atoms
to give up electrons to Zn
2+
ions.
A remarkable feature of oxidationreduction reac-
tions is their ability to spontaneously proceed even if
we confine the redox couples to separate containers,
called half-cells. In order to do so, we must connect
the half-cells with a wire, and, to maintain a charge bal-
ance in each half-cell, we must also connect the two
half-cells with a device that permits ionic flow into and
out of the cells. We call such a device a salt bridge.
In our example, we place a Zn strip in a container
with an aqueous Zn
2+
ion solution, and a Cu strip in a
separate container with an aqueous Cu
2+
ion solution.
When we connect the two metal strips with a wire and
add a salt bridge between the two solutions as shown
in Figure 1, electrons will pass through the wire from
the Zn strip to the Cu strip. We call these metal strips
electrodes. These are the components of the respec-
tive half-cells on the surface of which the actual
half-reactions occur. They are also the portions of the
half-cells that directly connect to the external circuit
or wire. We call the electrode in the half-cell that
receives the electrons the cathode, and the electrode
in the half-cell that supplies the electrons the anode. In
our example, Cu is the cathode, and Zn is the anode.
The passage of electrons from the anode to the cath-
ode through the wire is called an electrical current.
We measure current in amperes (A).
Electrons are driven from the Zn/Zn
2+
ion half-cell
to the Cu/Cu
2+
ion half-cell by the difference in electri-
cal potential, called the electromotive force or EMF,
between the two half-cells. We measure this potential
difference in volts (V). The assembly of two half-cells
connected to form a circuit such as that shown in Fig-
ure 1 is called an electrochemical cell, or a galvanic
or voltaic cell. A spontaneous redox reaction results in
an electrochemical cell with a positive voltage.
We can measure the net EMF for a pair of
half-reactions in an electrochemical cell by connecting
a voltmeter between the half-cells. To do so, we attach
the anode to the negative voltmeter terminal and the
cathode to the positive voltmeter terminal. Thus, for the
cell shown in Figure 1, we attach the wire from the
Zn/Zn
2+
ion half-cell to the negative voltmeter terminal
and the wire from the Cu/Cu
2+
ion half-cell to the posi-
tive voltmeter terminal.
The magnitude of the cell EMF and the direction
of electron flow through the external wire depend on
the relative abilities of the metal/metal ion couples to
give up or accept electrons. These abilities depend on
the inherent chemical natures of the specific
metal/metal ion couples we use to construct the cell.
For instance, suppose that we replace the Zn/Zn
2+
ion half-cell in Figure 1 with a strip of silver metal (Ag)
suspended in an aqueous silver ion solution (Ag
+
).
When we connect the Ag/Ag
+
ion half-cell to the
2
ELEC 419/Studying Electrochemistry
Figure 1
An electrochemical cell +
voltmeter
filter paper saturated with
KNO
3
(aq) used as salt bridge
Cu electrode
CuSO
4

solution
Zn electrode
ZnSO
4

solution
e e
positive voltmeter terminal and the Cu/Cu
2+
ion
half-cell to the negative terminal we observe a positive
voltage reading. In practice, without knowing before-
hand which electrode will be the anode and which the
cathode in the spontaneous redox reaction, the wires
are connected to the voltmeter by trial-and-error so as
to yield a positive voltage. This reading tells us that the
Cu/Cu
2+
ion half-cell gives up electrons more readily
than the Ag/Ag
+
ion half-cell does. Hence, the electron
flow is from the Cu anode to the Ag cathode in the ex-
ternal circuit. This observation also indicates that if we
place a Cu strip in an aqueous Ag
+
ion solution, the Cu
will dissolve, and Ag will deposit on the Cu strip.
By testing different pairs of metals in solutions of
their salts in electrochemical cells, we can rank the
various metal/metal ion couples in order of their rela-
tive abilities to accept electrons. This type of listing is
called an electromotive force series. Table 1 shows
the reduction half-reactions for a series of common
metal/metal ion half-cells, arranged in order of increas-
ing EMF. As the half-cell EMF becomes increasingly
positive, the ability for the half-cell to accept electrons,
or to be reduced, increases. Accordingly, these
half-cell EMFs are sometimes called the reduction po-
tentials for the half-reactions. From Table 1, we can
predict the direction of electron flow for a cell com-
posed of any two metal/metal ion combinations listed
in the table. We can also predict the chemical reactions
that might occur when various redox couples are
mixed in aqueous solution.
Table 1
Reduction half-cell reactions arranged
in order of increasing ability to accept electrons.*
reduction half-cell reaction
EMF, V
Mg
2+
(aq) + 2 e
Mg(s)
2.37
Al
3+
(aq) + 3 e
Al(s)
1.66
Cr
2+
(aq) + 2 e
Cr(s)
0.91
Fe
2+
(aq) + 2 e
Fe(s)
0.44
Ni
2+
(aq) + 2 e
Ni(s)
0.26
Sn
2+
(aq) + 2 e
Sn(s)
0.14
Cu
2+
(aq) + 2 e
Cu(s)
+0.34
Ag
+
(aq) + e
Ag(s)
+0.80
* The EMF can also be defined as the standard reduction
potential for that half-reaction.
As an illustration of how we can use Table 1, con-
sider the Ag/Ag
+
ionCu/Cu
2+
ion cell just described.
The EMF of the Ag/Ag
+
ion half-cell, whose reduction
half-reaction is shown in Equation 4,
Ag
+
(aq) + e
Ag(s)
(Eq. 4)
is +0.80 V. The EMF of the Cu/Cu
2+
ion half-cell,
whose reduction half-reaction is shown in Equation 5,
Cu
2+
(aq) + 2 e
Cu(s)
(Eq. 5)
is +0.34 V. The more positive EMF of the Ag/Ag
+
ion
half-cell indicates that this half-cell has a greater ability
to accept electrons than does the Cu/Cu
2+
ion
half-cell. Th