Electrochemistry
Electrochemistry is somewhat of a step-child in the general chemistry
curriculum. Often left for quick treatment at the end of the semester,
“hands-on electrochemistry labs are thought difficult and expensive.
This need not be true. Here are some new tools that will make
electrochemistry understandable, affordable, and easy and fun to teach –
and learn!
● Oxidation – Reduction
● Electroplating
MicroLab’s inexpensive Model 232 Electroplating Module provides three volts
DC battery power from two AA batteries and a lamp to monitor electron flow
in basic electroplating experiments. It can also measure ionic conductivity
and differentiate strong from weak acids.
●
Half Cell Reactions
●
The Electrochemical Series
●
The Nernst Equation
MicroLab’s Model 152 Mult-EChem Half Cell Module has space for eight metal / ion
electrochemical half cells, each equally accessing a central aqueous salt bridge through a
replaceable porous cylinder. Small ( 3 mL) samples provide extremely stable electrochemical cell
voltage measurements with the MicroLab FS-522/FS-524 Lab Interface or a digital voltmeter.
Getting Started with Electrochemistry
Because many of us work with limited equipment funds, or would like to explore a new lab topic before investing much in equipment,
here’s how the five introductory topics illustrated above can be demonstrated and taught with inexpensive apparatus. Exploration of
concepts of oxidation, reduction, and electroplating require only a battery, a flashlight bulb to monitor current, and two clip leads.
The electrochemical series and the Nernst Equation may be developed experimentally using small, safe, and low-cost chemical
samples using our inexpensive Model 152 half-cell module and a common digital voltmeter.
Development of these concepts from observation and experimental data rather than memorization is consistent with new teaching
strategies encouraged for AP and college/university general chemistry by the College Board and the National Research Council.
Computer-based Data Acquisition Tools and Software for Chemistry
2
Forced Electrochemical Reactions: Electroplating
Introducing Oxidation and Reduction: Electroplating
Electrons are the currency in electrochemical reactions. The experiments discussed in this mini-catalog all depend on transfer of electrons from one
atom, ion, or molecule to another. Sometimes these electron transfers can be externally forced, and sometimes they are spontaneous.
Each field of study or work develops its own language. Chemistry is no exception. We have words that we use to tell whether an atom, ion, or
molecule steals or accepts an electron (thus reducing its charge, or moving the overall charge in the negative direction), or if it gives up an electron,
thus moving its overall charge in the positive direction. These words are ...
Reduction and Oxidation:
Reduction:
Gain of an electron, moving the overall charge in the negative
direction, or reducing it.
Oxidation:
Loss of an electron, moving the overall charge in the positive
direction.
When something reacts with oxygen gas – burning is a good example - it gives up
electrons to the oxygen gas molecule. Oxygen is the most common element in the
Earth’s crust and reactions with oxygen are so common that loss of electrons is
generically called “oxidation”, even if an oxygen molecule did not take the electron.
How can a beginning chemistry student watch this trading of electrons happen? It is
easy: Electroplating. Electroplating experiments are safe, fast, inexpensive, and produce a visible, useful result.
An Electroplating Experiment:
There are three experiments one can do with this system:
The MicroLab Model 232 Electroplating Module uses two AA batteries
to provide voltage to force electrons to participate in electrochemical
reactions. A simple flashlight bulb monitors the flow of electrons
from the batteries to the reaction and back.
1.
Partially immerse the key in the Copper (II) Sulfate solution for
ten or twenty seconds. Observation: Nothing happens.
2.
Connect the clip lead holding the key to the positive battery
terminal and the lead connected to the copper wire to the
negative battery terminal. Immerse part of the key and copper
wire in the solution. The light bulb will show passage of current.
After 20 seconds, remove and look at the key. Observation:
Probably some corrosion or removal of material is visible on the
immersed part of the key.
3.
Reverse the clip leads so the key is connected to the negative
battery lead and the copper wire is connected to the positive
battery lead. Immerse key and copper wire in the solution. The
light bulb will show passage of current. After 20 seconds, look at
the key. Observation: There will be a metallic copper deposit on
the key. Continue electroplating for a minute or two. If you
immerse more of the key, the current will increase and the bulb
will shine brighter. Buff with steel wool to brighten.
Here’s the experimental set-up:




Put about 30 mL of approximately 1.0 M CuSO4 solution in a
50 mL beaker. Copper (II) Sulfate usually is provided as the
pentahydrate, CuSO4 ● 5 H2O. It molecular mass is 159.6
g/mole, so 16 grams dissolved in 100 mL of water will
produce a solution approximately 1.0 M in Cu2+. Slightly
acidify with 2 mL of 0.1 M H2SO4 to keep Cu (II) hydroxide
from forming and precipitating.
Make a copper electrode from about six inches of bare
copper wire, making it into a coil and hanging the straight
end of the wire as a hook over the side of the beaker.
Connect a clip lead to the copper wire.
Connect another clip lead to a key, coin, or other clean
metallic object.
What happened during Experiment 3?
When the negative half of the battery forces electrons out onto the key, aqueous Cu 2+ ions in contact
with the key will pick up the electrons, becoming metallic copper:
Cu2+ + 2 electrons  Cu 0
This gain of electrons is reduction.
The positive battery terminal is connected to the copper wire. It pulls electrons away from the metallic
Cu0 atoms in the wire, making Cu2+ ions which float out into the solution to replace the Cu2+ ions that
were removed at the negative electrode. Oxidation occurs at the positive electrode. Removal of the
copper atoms from the surface of the wire leaves an etched look.
Electron
motion
Spontaneous Electrochemical Reactions: Half Cells and the Electrochemical Series
3
3
Spontaneous Electrochemical Reactions: Half Cells and the Electrochemical Series
3
Electrochemical reactions are of two types: Reactions in which the electron transfer is
forced, and reactions in which the electron transfer is spontaneous.
Electroplating is an example of a forced electrochemical reaction. The battery in the
electroplating module provided the energy to force electrons from the oxidation to the
reduction reaction.
Spontaneous transfer of electrons will occur between two atoms or ions if one has a
greater attraction for electrons than the other, and they come in contact with each other.
These spontaneous reactions always release energy and are responsible for corrosion of
metals, for photosynthesis, and they provide the power for flashlights and electric
automobiles.
It is easy to predict if an electron will transfer occur between any pair of reactants if one
has a list that ranks these elements and ions in order of their ability to steal electrons from
one another. Chemists call this list the “Electrochemical Series”. There is nothing
theoretical or magic about it. It is experimentally-based and can easily be developed in an
hour in the lab.
The electrochemical series creates problems for students because its scale is relative instead
of absolute, and because the sign conventions are often memorized instead of
being thought out. Let’s try to make this easy to understand.
Experimental Development of the Electrochemical Series:
Ordering Elements and Ions in terms of their Reduction Potential
1.10
Voltmeter
Electrons flow from the zinc half-cell to the copper
half-cell. Zn is oxidized, Cu 2+ is reduced. Copper
has a greater attraction for electrons than Zinc. It
is higher on a relative scale of reduction potential.
Consider the experiment illustrated to the upper right. The MicroLab half-cell module contains eight small wells, each touching a central porous
cylinder which is saturated with aqueous potassium nitrate from the center well. This cylinder acts as a “salt bridge”, connecting the several “halfcells” around its perimeter. [More on the salt bridge on page 5). A milled overflow area in each half-cell cylinder prevents spills and mixing of
solutions.
The left “half-cell”, Number 7, contains a solution of 0.1 M Copper (II) nitrate and a copper wire (CuO). The right “half-cell”, number 3, contains a
solution of 0.1M Zinc (II) nitrate and a strip of metallic zinc (ZnO). When a voltmeter is connected between the two metallic electrodes, electrons
spontaneously move from the metallic zinc electrode (black lead) through the voltmeter to the metallic copper electrode (red lead). The
voltmeter reads + 1.10 volts. (Voltmeters read + when electrons move into their black lead and out their red lead. You can use this to determine
which way the electrons are moving.)
What can we learn from this experiment?
1.
2.
Each “half-cell” contains materials for both a reduction reaction and an
oxidation reaction. These reactions are illustrated to the right, with arrows
pointing both directions. These are reversible reactions. By convention they are
always written in the forward direction (to the right) as reduction reactions.
Cu2+ + 2 electrons
Cu O
Zn2+ + 2 electrons
Zn O
Electrons are observed to spontaneously move from the Zinc half -cell to the
Copper half cell. This means that a reduction reaction was taking place in the
copper half-cell, and an oxidation reaction was taking place in the zinc half-cell.
The copper half-cell reaction appears to have a greater attraction for electrons
than the zinc half-cell reaction. The red arrows show electrons moving from the
zinc half-cell reaction to the copper half-cell reaction, passing through the
voltmeter.
2+
Cu
+ 2 electrons
Zn2+ + 2 electrons
Cu
O
Zn O
Reduction Reaction
Oxidation Reaction
When Cu 2+/ Copper metal and Zn 2+ / Zinc metal half cells are
coupled with a salt bridge, electrons will flow from the zinc to the
copper in an outside wire. Note that both reactions are written to
the right as reduction reactions, the red arrow shows the actual
direction of each reaction. Copper has greater attraction for
electrons than zinc.
4
Spontaneous Electrochemical Reactions: Half Cells and the Electrochemical Series
Two generalizations can be drawn from this experiment:

We can order the half-reactions in terms of electron-attracting ability. Copper (II) ion can take electrons away from Zinc (O) metal.
We can assign a voltage difference to the half-cell pair. The electrical force moving electrons from one half-reaction to the other is
measured as 1.10 volts. This means that we could assign a value of + 1.10 volts to the Cu half-cell and a voltage of O to the Zn half-cell.
Or a voltage of O to the Cu half-cell and a voltage of – 1.10 volts to the Zn half-cell. It is completely arbitrary how this assignment is
done, but it is easier to talk about positive voltages than negative voltages, so the copper half-cell could be assigned a voltage of + 1.10
volts. But notice: This is + 1.10 volts with respect to the zinc half-cell, which now is arbitrarily assigned a value of 0 volts. It has become
a “zero voltage reference”.
Increasing tendency
to gain electrons

4.
1.10 volts
3.
Cu2+ + 2 electrons
Cu O
Reduction Reaction
Zn2+ + 2 electrons
Zn O
Oxidation Reaction
There are two take-home lessons in this simple experiment:

It always takes two half-reactions to move an electron. One will be a reduction reaction, taking the electron, and the other will be
an oxidation reaction, supplying the electron. The combination of the two half-cells is called an “electrochemical cell”.

The force with which electrons are moved from one half-reaction to the other is determined by the difference in electron-attracting
ability (reduction potential) of the two half-reactions.
Now, if Copper and Zinc were the only two elements in the world, our
job would be done. We have a simple “electrochemical series” or scale
that says, if copper ion and zinc metal ever touch, the copper ion will
take two electrons away from the zinc metal atom. Depending on the
concentration of the ions and the temperature, you will see a voltage of
about 1.10 volts moving the electrons. You could make a simple
flashlight battery with these two half-cells.
There is more to it, however. In the next page we will consider
movement of the associated negative ions that are always present. But
for now, consider the complication that arrives when you add more
elements to the electrochemical series:
The table to the right shows five metallic elements and their associated
ions, and a reaction in which hydrogen gas breaks into two hydrogen
ions and two electrons.
This electrochemical series, developed by using the Model 152 half-cell
module and half-cells of each of the five metals and their associated
ions, lists silver ion as the best electron-attractor, and zinc metal as
having the least hold on its electrons. These are experimentally ordered
as were the copper and zinc half-reactions, in order of each halfreaction’s ability to take an electron from another.
How was the data for this series collected? Five of the metal / ion halfcells were set up in the half-cell module and its salt bridge filled with
potassium nitrate solution. The black lead of the voltmeter was
arbitrarily connected to the lead electrode, and then the red lead moved
to the other half-cells to see which way the electrons moved and with
how much force (voltage). This data was recorded on the right-hand
scale on the table to the right.
What’s special about the lead/lead ion half-cell? Nothing. It was
arbitrarily chosen as a “reference” and assigned a value of O. All of the
other half-cell reactions were compared to it. Some were better
0.45 volts
0.45 volts
1.10 volts
1.10 volts
Arbitrary Hydrogen Reference
Arbitrary Lead Reference
electron attractors, Cu and Ag, for example. These reduction potentials
were given positive signs because electrons moved toward them from
the lead. Two (Fe and Zn) gave up electrons to the lead half-cell, so they
received negative signs.
The voltages on the right-hand scale of the table above were those
measured with respect to the arbitrary Pb / Pb2+ half-cell reference.
Those on the left-hand scale were measured with respect to a hydrogen
half-cell. The hydrogen half-cell is hard to implement experimentally,
but is defined by the IUPAC as an international standard or “zero”
position in the electrochemical series.
Textbooks give electrochemical half-cell voltages (reduction potentials)
with reference to hydrogen as a zero. The voltage produced by any two
half cells is just the absolute difference between their positions in the
electrochemical series. If they are on the same side of zero, subtract. If
they cross zero, add. The hydrogen half cell reference is not magic. It is
just an arbitrary (but internationally-agreed upon) zero reference.
The Role of the Salt Bridge / Which way do the Electrons Go?
5
The Role of the Salt Bridge
It has to do with maintaining electrical neutrality in each half-cell.
An electrochemical cell produces energy by moving electrons
from one half-cell that gives up electrons (oxidation) to
another half-cell that grabs the electrons (reduction).
The reduction reaction starts with a positive metal ion teamed up with one
or more negative ions to preserve electrical neutrality. In this example, the
left-hand cell starts with each Cu 2+ balanced by two (NO3) - ions. When
the Cu 2+ ion accepts electrons and is reduced to metallic Cu 0, there are
left-over negative nitrate ions. This will produce an overall negative
change in the half cell, and repel any additional electrons coming in.
The electrons won’t move in the outside circuit unless a salt
bridge — an aqueous or gel connection containing mobile
ions -- between the two half-cells is present. Why is this?
Likewise, the oxidation reaction produces positive metal ions, and needs
negative ions to balance the additional positive charge in this half-cell.
aqueous or gel connection between the two half-cells — is
present. Why is this?
Half-Cell contact with the porous cylinder:
Make sure each half-cell solution level is at least to the top of its
inside well to make good contact with the porous cylinder. The
KNO3 level should be high enough to cover the central post in the
center well. This is just a reservoir of mobile K+ and NO3- ions. The
post reduces the volume of KNO3 required. The cell should be
soaked overnight before use in approximately 0.1 M KNO3 to
saturate the porous cylinder. Rinse with DI water after use. Store
in a zip-lock bag between labs.
The extra nitrate ions produced in the Copper half-cell can move to the
Zinc half- cell through a “salt bridge”, which in this case is a porous cylinder
saturated with aqueous potassium nitrate. Extra aqueous potassium
nitrate is provided in the center well to enhance the movement of nitrate
ions. Because the porous cylinder and salt bridge stay wet, ions are free to
move, and the electrochemical cell voltage remains very stable.
Which way do the electrons go?
For a beginner, probably the most difficult part of an electron-transfer
experiment is figuring out which way the electrons are moving. In the old
days we had center-zero moving needle galvanometers, and you could see
which way the needle leaned. Today we have digital voltmeters that
protest if the electrons are going the “wrong” way by putting a negative
sign in front of the voltage number. The convention is that electrons
running into the black (negative) lead of a voltmeter produce a positive
voltage display.
MicroLab’s FS-522/524 electrochemistry half-cell software display can help
with this. It gives three types of display:
 If electrons go in the black voltage lead, the meter reads to the right and
the voltage sign in the digital display is positive. (Figure on Page 1.)
The analog scale will shift to 4 volts if the input voltage is over 2 volts.
However, if the position of oxidation and reduction reactions is
reversed, the display will respond as shown to the right.
 Blue arrows show the direction of electron motion, and a light bulb lights
up to show the motion of electrons though a load.
 Generic oxidation and reduction reactions are
written in the correct direction for each half-cell.
If the position of the half-cells are reversed compared to the
illustration on Page 1, placing the oxidation reaction on the
right, the direction of electron motion will reverse, the
meter will reverse, the voltage sign will turn negative, the
electron motion arrows will reverse, and the position of the
reactions will reverse. Students can easily identify which is
the oxidation and which is the reduction reaction.
6
Nernst equation: Predicting the effect of changing ion concentration on cell voltage
The Nernst Equation: Predicting the Effect of changing ion concentration on Electrochemical Cell Voltage:
Walther Nernst (1864-1941) was a German physical chemist who
received the Nobel Prize in Chemistry in 1920 for his work in
thermochemistry. One of Nernst’s contributions was an equation
that describes the change in an atom or ion’s reduction potential (its
ability to take electrons from an atom or ion of lower electronattracting ability) when the ionic concentration, the temperature or
the number of electrons transferred during the reaction are
changed.
Nernst’s equation has a lot of practical application in chemistry. It
explains how common pH meters and ion-specific electrode meters
work. It explains why winter wildlife photographers keep their
camera batteries inside their coat until they are ready to shoot a
photo. And it gives a way of predicting the transport of ions across
cell membranes in nerve conduction.
Ecell
=
Eo
+
The voltage calculated for the
two half cells when the
temperature is 25 o C and the
ion concentrations are 1.0 M
..t.electrochemical cell
The sign for the correction factor is additive (+) if the ion is being
reduced, or subtractive (-) if the ion is being oxidized.
The voltage produced by
the electrochemical cell
Nernst’s equation describes only a single half-cell reaction – a
reaction involving only one ion. The voltage produced by a
complete electrochemical cell (a battery) can be predicted by
combining the Nernst equation statements for the two half cells
that make up the battery.
Nernst’s equation has been a thorn in the side of many generations
of chemistry students because they are immediately presented the
equation for the full electrochemical cell. In desperation they
memorize it rather than breaking it into its understandable parts.
Let’s make it easy to understand by considering only one of the two
half-reactions at a time. MicroLab’s Model 152 half-cell module
separates these two reactions into two visually and chemically
separate chambers, and connects them with aqueous salt bridge to
transfer negative ions. Now let’s look at the equation part-by-part:
(RT/nF)
*
log [Ion concn]
Correction Factor
The constant (RT/nF) compensates for changes in the absolute temperature T, and
the number of electrons transferred, n. It also includes the universal gas constant
R, and the Faraday, F, that relates moles of electrons to coulombs of charge. It is
multiplied by the log of the ion concentration to generate the correction factor.
Predicting Changes in Cell Voltage
Consider a silver / silver ion and copper / copper ion
electrochemical cell: Silver is higher on the reduction potential
scale than copper. Silver ion will take electrons from copper metal.
Copper metal will oxidize and silver ion will be reduced.
 What will happen if [Ag ] is increased? It should increase the
forward tendency of the reaction and increase the reduction
potential. Voltage measured between the Ag and Cu half cells
should increase. [Conceptual]
+
 Increasing [Ag+] will increase the correction factor, adding to the
overall cell voltage. [Quantitative]
Voltage vs Cu Reference
[Ag +] Half Cell Concentration
0.55
0.508
0.5
0.45
0.448
0.4
0.388
0.35
y = 0.0598x + 0.5676
Increasing concentration
0.3
-4
-3
-2
-1
Log [Ag+] Concentration
Cell voltage increases with increasing Ag+ concentration. It
changes 0.0598 volts per decade change in ion concentration.
 What would happen if [Cu 2+ ] were increased? It should
make the oxidation of the Cu more difficult. This shifts
the equilibrium more toward reduction and moves the
Cu half-cell upward on the reduction potential scale.
The overall Cell voltage should decrease because the
reduction potentials of the two half cells will become
closer together.
Cell voltage decreases
with increasing Cu 2+
ion concentration. The
.
slope is half that
of the
Ag+ experiment.
Increasing concentration
Here is the combined reaction and a simplified model. Increasing a reactant ion
concentration [Ag+] increases rate and will increase cell voltage. Increasing the product
ion concentration [Cu2+] will reduce the rate and will reduce the cell voltage.
0
The Nernst Equation / Getting Started with Electrochemistry
7
7
7
o
7 ● The value of [RT/F] is 0.0592 for base ten logs, at 25 C. For a one-electron change (Ag+), the voltage should change 0.0592 volts for
each power of ten change in [Ag+] concentration. From the experiment, the slope is 0.0598 volts/decade. This is how pH meters [H+]
work.
 Is the overall cell voltage more sensitive to changes in [Ag+ ] or [Cu 2+ ] ? n is 2 for the Copper reaction. The Cu 2+ slope will be
(0.0598 / 2 ) volts/decade, with the voltage decreasing with increasing concentration. So if the silver ion concentration is held
constant, you can make a p[Cu2+] meter that will read log copper ion concentration. This is how ion-specific electrodes work.
The Nernst equation corrections for each of the two half cells can be combined to create a single equation describing the complete electrochemical
cell:
Ecell = Eo - (RT/nF) * log [ Oxidized ion concn] + (RT/nF) * log [ Reduced ion concn], or algebraically rearranged ..
Ecell = Eo - (RT/nF) * log [ Oxidized ion concn ] / [Reduced ion concn]
measures voltage, or a simple digital voltmeter that you might have
on hand or purchase.
Getting Started with Electrochemistry
There are several ways to get your students started with the
electrochemistry experiments discussed in the first part of this minicatalog. The electroplating power supply and half-cell modules are
unique, rugged, require very small amounts of chemicals, and will
last a long time. They have been designed for inexpensive entry
into this field. To do electrochemical series and Nernst equation
experiments, your students must be able to measure voltage. This
can be a MicroLab FS-522/524, another brand of lab interface that
Model Number
232
152
151
233
Electroplating Module
Half Cell Module
7 element Metal kit
Digital Voltmeter
$44
$7
Price Individually
$28
Experiment
Electroplating, REDOX
Electrochemical Series, Nernst Equation
ElectroChem Kit 235
ElectroChem Kit 236
ElectroChem Kit 237
The table below shows equipment packages for these several
alternatives. The metal kit contains 3 cm lengths of wire
representing seven elements: Ag, Cu, Ni, Fe, Pb, Zn, and Al. Sample
experiments are available on the MicroLab web site. The next page
of this brochure shows additional electrochemical sensors and
modules that work with the MicroLab FS522 / 524. Please check our
web site for information concerning these units.
$ 49
$ 69
[152 + 151]
[232 + 152 + 151]
$ 94
[232 + 152 + 151 + 233]
$ 25
Voltage Measurement
More Electrochemistry with the MicroLab FS-522/524
Advantages of the
Mult-EChem design:
Our Model 274 Electrochemistry Electroplating / Coulometry Power Supply
operates with the FS-524 Lab Interface to provide a software- adjustable constant
voltage (0-5 volts, up to 1 amp) for electroplating or electrolysis experiments.
Software integrates current over time to calculate coulombs of charge or moles of
electrons delivered.
Electrochemical cells are easily constructed and clearly
visualized. These may be different metal / ion pairs
to develop the electrochemical series, or different
concentrations of the same metal / ion pair referenced
to a single half-cell to demonstrate the Nernst equation.
Our new Model 292 Electrochemistry Module operates with the FS-524 Lab Interface to
perform coulometry and voltammetry experiments in analytical and physical chemistry.
MicroLab’s unique electrically-isolated sensor amplifiers prevent coulometric cell current from
interfering with pH, Redox, ISE, and amperometric end-point measurements.
The salt bridge does not dry out. MicroLab’s
Electrochemical Cell voltages are extremely stable for
long periods of time.
It requires very small samples (3 mL). Experiments are
safe and inexpensive.
The polypropylene body of the module is rugged and
chemically resistant. Its non-skid feet and 200 g mass
keep it securely in place on the lab bench. It will have a
long service lifetime.
Web Tutorials
An excellent web tutorial on Galvanic
cells and the Nernst Equation by
Dr. Sophia Nussbaum of the University
of British Columbia and the
Chem Collective is available at ...
MicroLab ‘s Model 170 Cyclic Voltammetry Module uses
inexpensive Pine Instrument screen printed electrodes. It
connects to one of the FS-524’s multipurpose inputs, and
will scan with 1-20 mV steps.
http://collective.chem.cmu.edu/chem/ubc/exp11/step1_investigate.php
MicroLab’s NEW FS-524 FASTspec +plus™ and its integrated
360-880nm scanning spectrophotometer will make
almost every instrumental measurement required in
general and environmental chemistry and biology. It will
serve college/university chemistry courses freshman
through senior and undergraduate research.
MicroLab Inc. ● microlabinfo.com ●
(888) 586-3274 ● PO Box 7358, Bozeman, MT 59715