| When an electrolyte is dissolved in water it will cause an increase in the
electrical conductivity. In this beaker an electrolyte has been added to the water.
Perhaps you can see that the gas bubbles are more abundant here than they were in the
picture of the electrolysis of water. |
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| What happens during the electrolysis of a solution depends on what is in
the solution and even materials that electrodes are made of, as you will see. In this case
silver nitrate has been added to the water. The electrodes are made of copper. |
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| After a short period of time, you can see that silver crystals have formed
on the cathode (left side, black wire) and have begun to fall off. On the right side,
around and below the anode (connected to the red wire), I hope you can see a light blue
color. Let's investigate the reactions that caused these changes. |
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| In example 2 in your workbook you will find a diagram similar to this. I
recommend that you use it for notes. As we work through this example, I will allude
to various possible reactions that might occur and expand upon those that
do occur. The reason for this is to emphasize that generally more than one
reaction might occur even though one might be favored over the others. For now, realize
that there are a variety of reactions that might occur. Later in this lesson we will talk
about how you can decide which of the possibilities is most reasonable. |
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| In this particular case both of the electrodes were made of copper. The
solution was, of course, primarily water. It also contained Ag+ and NO3-.
|
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Cathode Reaction
The cathode reaction in this case is the simplest so let's deal with it first. At the
cathode, electrons from the negative pole of the power supply are coming to the solution.
These negative electrons attract silver ions (Ag+) and combine with them to
make silver metal. The process is reduction because the charge
(oxidation state) of the silver decreases (Ag+1  Ag0). The equation for the reaction is
Ag+ + e- Ag. |
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Anode Reaction
At the anode we have several possibilities to consider. At the anode, electrons are
being taken away and moved to the positive pole on the power supply. But what are they
being taken away from? They could be taken away from the negatively charged NO3-
ions, but they're not. They could be taken away from water to make O2, but
they're not. Instead they are taken away from the copper atoms that make up the anode. |
| Copper atoms react by giving up electrons and becoming copper(II) ions.
The equation for the reaction is Cu 2 e- + Cu2+. One
bit of evidence for this is the blue color that appeared in the solution near the anode.
Cu2+ ion is blue in water. More evidence to verify this would be to weigh the
electrode before and after the reaction. If the reaction runs long enough, there would be
a detectable weight loss in the anode itself. |
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Questions to Ponder
Let's refer back to the electrolysis of water and consider a few questions.
Why didn't the anode oxidize? This question brings up the idea
of a different ease of oxidation for different
chemicals. Later in this lesson we will work with a list of oxidation abilities.
Why didn't something plate out on the cathode? This question brings up
the issue of availability. Only what is there can react. If more than one chemical is
available, do they all react? Just one? Which one(s)? Later in this lesson we will also
work with a list of reduction abilities.
What about the voltage needed for electrolysis? Will any amount do? Is
it different for solutions? Later in this lesson we will also work with a list of voltage
potentials.
For now, move on to the electrolysis of a molten salt.
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Eden Francis
Clackamas Community College
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