Video Transcript
In this video, we will learn how to
describe primary cells and explain how they produce electrical energy. In the first section of this video,
we’ll see how electrical energy comes from cells used in portable electrical devices
that we see in everyday life.
Portable electric devices like
torches or flashlights get their energy from chemical reactions that take place
inside electrical cells. Electrical cells are often referred
to as batteries, although a battery is strictly several cells connected
together. A primary galvanic cell is a
single-use electrochemical cell where electrons are generated through a spontaneous
redox reaction. When the circuit inside the
flashlight is completed by closing the switch, electrons flow around the circuit as
a result of spontaneous chemical reactions in the cells. The flow of electrons results in an
electric current. Since electrons carry an electric
charge, an electric current is a measure of the rate of flow of electric charge.
The type of chemical reactions that
occur in each cell are called redox reactions. In a redox reaction, electrons are
transferred from one chemical species to another. Redox reactions therefore involve
species being either oxidized or reduced. If a species is oxidized, it loses
electrons. If a species is reduced, it gains
electrons. Oxidation and reduction are easily
remembered using the mnemonic “OIL RIG.”
Quite often, flashlights are
powered by relatively cheap zinc–carbon cells. These cells have a zinc casing. And in the reactions that drive
these cells, the zinc loses electrons. The zinc casing is therefore the
site of oxidation. In a galvanic cell, the site of
oxidation is the negative electrode. Sometimes this is also referred to
as the anode. Electrons are transferred from the
negative electrode to a carbon electrode, which is the positive electrode in this
cell. In a galvanic cell, the positive
electrode is the site of reduction. It’s also known as the cathode
sometimes.
The carbon cathode collects
electrons, and it is surrounded by a paste containing manganese dioxide. Manganese dioxide accepts the
electrons in a reduction reaction. So overall in these cells,
electrons flow from the zinc anode to the carbon cathode. Of course, electrons also flow
through the light bulb, which causes it to light up. The chemical reactions in the cell
continue until all of the reactants are consumed. The zinc casing will become thinner
as the zinc turns into zinc ions in the electrolyte in the cell. The chemical composition of the
manganese dioxide paste will also change in the reduction reaction. Over time, the bulb will become
dimmer, and the cell is eventually spent or discharged. It can only be used once. It’s then discarded or recycled to
recover any contents of value.
Primary galvanic cells operate in
this way. The chemical reactions that take
place spontaneously inside them are not reversible. So when the flashlight stops
working, the cells inside it cannot be recharged.
In the next section of this video,
we’ll take a detailed look of a specific type of primary galvanic cell called the
mercury cell to understand how the electrons flow within it.
The mercury cell consists of a
steel can that acts as a container, but it’s also in contact with mercury two oxide
that has the formula HgO. This part of the cell forms the
positive electrode or cathode, which is where a reduction reaction will take
place. The cap of this cell is in contact
with zinc metal, and this forms the negative electrode of the cell. The negative electrode or anode in
this cell is where an oxidation reaction will take place. Use the mnemonic CROA to help
remember where reduction and oxidation take place. The cathode is the site of
reduction, and oxidation takes place at the anode.
A porous separator keeps the
chemicals in the anode and cathode apart. The porous separator contains
electrolytes that allow for the movement of ions, and hence charge, to complete the
circuit when the battery or cell is operating. To prevent a short circuit in the
cell, an insulating seal separates the anode and the cathode. It also contains the chemicals to
stop them leaking out.
Now let’s take a look at the
chemical reactions occurring in this cell and the voltage that it generates. At the anode or negative electrode
in this cell, electrons are lost. An oxidation reaction occurs. Zinc metal is oxidized to zinc
oxide. At the positive electrode or
cathode in this cell, mercury two oxide is involved in a reduction reaction. Mercury two oxide is reduced to
mercury, which is a liquid metal in this cell. By adding the two half-equations
together, we get the overall reaction that occurs in this cell. Zinc metal reacts with mercury two
oxide to produce zinc oxide and mercury.
When we added the two
half-equations together, the two moles of hydroxide ions on each side of each
half-equation canceled each other out. This also happened for the one mole
of liquid water molecules on each side. The electrons also canceled out, as
there were two moles of electrons involved in the oxidation process and two moles of
electrons involved in the reduction process.
To calculate the 𝐸 standard for
the cell, that is, the cell emf or the voltage produced by this cell, we need to
know the standard reduction potential for the cathode and the standard reduction
potential for the anode. We simply subtract the standard
reduction potential for the anode from the standard reduction potential from the
cathode. Standard reduction potentials are
measured against the standard hydrogen electrode under standard conditions. The standard hydrogen electrode is
by reference given a voltage of zero volts.
The standard reduction potential
for the anode in this cell is negative 1.25 volts. In this cell, the standard
reduction potential for the cathode, where mercury two oxide turns into mercury, is
positive 0.0977 volts. The voltage produced by this cell
under standard conditions is therefore positive 0.0977 volts subtract negative 1.25
volts. This voltage has been rounded to
two decimal places.
Since mercury cells contain toxic
mercury compounds, they’re not produced in large quantities anymore. Because the cell could be
manufactured to be very thin, these batteries or cells were known as button cells,
and they were used widely in watches, hearing aids, and calculators. Since the chemical reactions in
this cell are not reversible, the cell cannot be recharged. This type of cell would need to be
either disposed of carefully or recycled due to the toxic mercury content.
In the next part of this video,
we’ll take a look at hydrogen fuel cells.
Fuel cells are a type of galvanic
cell where the chemical reactants are stored separately outside the cell. The reactants are fed into the fuel
cell where the chemical reaction takes place. The redox reactions inside the fuel
cell can provide an electric current on demand. This might happen when a device
like an electric motor is connected to the fuel cell and is doing work. The fuel cell will produce a
constant voltage provided the reactants are fed into the fuel cell at a constant
rate.
The most common type of fuel cell
is the hydrogen fuel cell. Hydrogen gas is fed into the cell
from an outside store, where it reacts with oxygen gas, which is drawn in from the
outside air. Inside the fuel cell. The anode and the cathode are
separated by an electrolyte-filled membrane. In the acidic hydrogen fuel cell,
the membrane is filled with an acidic electrolyte. At the anode, hydrogen gas is
oxidized, and electrons are released. These can be made to flow around an
external circuit through an electric motor for example.
The remaining hydrogen ions or
protons migrate through the electrolyte membrane to the cathode. In the acidic hydrogen fuel cell,
the membrane is frequently referred to as a proton-exchange membrane. At the cathode in the fuel cell,
oxygen gas is reduced. Oxygen molecules gain electrons
returning from the external circuit. They react with hydrogen ions, and
they form water. Whilst producing electricity, the
hydrogen fuel cell makes water as the only waste product. This type of fuel cell has been
used to produce electricity in space travel, and it also produces electricity for
some electric cars.
Let us take a look at the reactions
occurring at each electrode in more detail. At the anode, hydrogen gas is
oxidized. Hydrogen ions and electrons are
released. The standard reduction potential
for this process is zero volts. At the cathode, we saw that oxygen
gas gains electrons and reacts with hydrogen ions to produce water. The standard reduction potential
for this reaction is positive 1.23 volts. We can balance the electron
transfer in each half-equation by multiplying the top equation by two. The overall equation for the
reaction occurring inside the fuel cell is then obtained by adding the two
half-equations together.
The equation simplifies to two
moles of hydrogen gas reacts with one mole of oxygen gas to make two moles of
water. The standard cell potential for
this fuel cell can be calculated using the equation standard cell potential equals
the reduction potential of the cathode minus the reduction potential of the
anode. For this type of fuel cell, the
standard cell potential is positive 1.23 volts. The cell voltage remains constant
as long as the reactants are supplied to the cell at a constant rate.
Since water is the only waste
product, the hydrogen fuel cell provides a source of electricity that does not
directly produce carbon dioxide. Carbon dioxide is emitted by
combustion engines in cars. One idea is to replace combustion
engines in cars with fuel cells to power electric motors. We must consider, however, that the
hydrogen gas for the fuel cell must be manufactured, since it’s not available in our
atmosphere.
A hydrogen fuel cell may also
operate in alkaline conditions. In the alkaline fuel cell, oxygen
gas is fed to the cathode, where a reduction reaction takes place. Oxygen molecules gain electrons and
react with water molecules to produce hydroxide ions. The hydroxide ions move through an
alkaline electrolyte membrane. Hydrogen gas is fed to the negative
electrode or the anode. At the anode, hydrogen reacts with
the hydroxide ions that have moved through the alkaline electrolyte membrane. Since electrons are produced at the
anode, and the anode has a more negative electrode potential than the cathode, the
electrons can be made to flow around an external circuit where they can do work.
The electrodes are often made from
ceramic materials coated in a thin layer of platinum, which helps to catalyze the
reactions. It may come as a surprise to see
that the cell potential for the alkaline fuel cell is the same as the cell potential
for the acidic fuel cell when under standard conditions. However, when we add the two
half-equations together, the overall reaction in the alkaline fuel cell is the same
as the overall reaction in the acidic fuel cell, so the voltage produced is the
same.
Now it’s time to look at a question
to test your understanding of galvanic cells.
Which of the following reactions is
the half-equation for the reaction that occurs at the cathode in a mercury cell.
To answer this question, we need to
remember that the cathode is the site of reduction in the galvanic cell. In a galvanic cell, oxidation
happens at the anode. In a reduction reaction, electrons
are gained. So we’re looking for a
half-equation where the species on the left, that is, the reactants, gain
electrons. The electrons have to appear on the
left side of our half-equations.
We can rule out answer (C) as the
electrons are being lost in this half-equation. This half-equation represents an
oxidation process. We can also rule out answer (D), as
this is not a half-equation; it’s a full redox equation. In (D), zinc loses electrons, its
oxidation state increases, and it’s being oxidized. Hg or mercury is being reduced. It gains electrons as its oxidation
state decreases. For the same reason that we ruled
out answer (C), we can also rule out answer (E). Although it appears to be a
half-equation, we can see that the electrons appear on the right-hand side. They’re being lost, and it’s an
oxidation reaction. Remember, we’re looking for a
reduction process.
On closer inspection of answer (B),
we can see that there appears to be no redox processes taking place at all. The oxidation state of mercury
remains at plus two throughout the equation. Since we’re looking for a reduction
reaction, it can’t be the right answer. Mercury cells contain mercury two
oxide or HgO at the cathode. In answer (A), we see mercury two
oxide gaining electrons. Mercury two oxide is the correct
species for the cathode in our cell, and it’s involved in a reduction reaction
here. This is the correct answer.
Now it’s time to review the key
points from this video. A primary galvanic cell is a
single-use electrochemical cell where electrons are generated through a spontaneous
redox reaction. Reduction happens at the
cathode. Oxidation occurs at the anode. The standard cell potential can be
calculated by taking the standard reduction potential of the cathode and subtracting
the standard reduction potential of the anode. Fuel cells are a type of galvanic
cell where the reacting components are constantly supplied.
