Mini-Lab 6-03: Going my way?
Purpose:
Electrolytic cells are very useful for electroplating, recharging batteries, etc. Can you
build one and figure out which direction the electrons are flowing?
Think about this:
An electrolytic cell is just the opposite of a galvanic cell. Electrons are forced into it by
an outside source, and the chemical reactions that are forced to occur are nonspontaneous.
Pencil “lead” (graphite) makes a good electrode in aqueous solutions and the graphite
does not participate in any chemistry.
If aqueous iodide is oxidized, it will form aqueous I2, which will look kind of brown.
Safety
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Don’t eat or drink anything in the lab.
Always wear eye protection.
Wear protective clothing (lab coats, etc.).
Don’t play around – treat the lab with respect.
Questions
1) Using a small container (like a 24-well tray), set up an electrolytic cell using
potassium iodide solution as the electrolyte and a 9-volt battery as the power
source. Remember to conserve materials and use as little as possible.
2) Write down your observations.
3) Write the balanced equations for the half-reactions at each electrode.
4) Draw a picture of the cell and indicate which graphite electrode is the anode and
which is the cathode.
WAIT! Do not write down an answer to the Final question until your Instructor
tells you to.
5) FINAL: Which pole of the battery is producing the electrons? Explain how you
know this.
Instructor’s Page
6-03: Going my way?
Source: SCALE-UP (“KI Electrolysis”)
Concepts: reduction/oxidation potentials, electrolysis
Materials: 1.0 M KI solution, 9V battery, clip wires, pencil lead, phenolphthalein,
saturated sodium chloride or potassium nitrate, strips of filter paper
Hints: Their first idea of what is happening is to say that iodide is being oxidized at the
anode and potassium is being reduced at the cathode. It may be easier to see the colors
form if they set this up in two different wells on a 24-well tray and use a piece of filter
paper as a salt bridge (see photos, below).
Some students will immediately point out that the “negative” pole of the battery produces
the electrons. “Fine”, you might say, “Now I’d like you to prove it!” Draw their
attention to their observations and help them think through what they see. For example:
- “What does the brownish color near that electrode tell you?”
o that iodine is being produced, which means iodide is being oxidized
- “If iodide is being oxidized at this electrode, what kind of reaction MUST be
happening at the other electrode?”
o Reduction
- “What do you see forming at the electrode where oxidation is taking place?”
o Gas bubbles
- “So, what kind of gas could be formed from a reduction reaction in a water-based
solution?”
o The only possibility is Hydrogen gas. If oxygen gas were being formed, it
would be from the oxidation of O2-.
- They might initially think that potassium is being reduced, but if so, you can ask them
this: “Wouldn’t you see the formation of potassium metal, if this was true?”
Ultimately, it comes down to the fact that splitting water to produce hydrogen gas and
hydroxide ions takes a lower voltage than producing potassium metal.
Once they realize that the oxidation is occurring at the electrode where hydrogen gas is
produced, then they should understand that the battery is providing the electrons for this
to happen. Indeed, it is the pole marked “negative” that produces the electrons.
Correspondingly, the other pole of the battery is “sucking up” electrons from the
iodide  iodine reaction.
Discussion ideas: This is a good follow-up to the Activity of Metals Mini-Lab, since
you could use it to discuss why potassium metal is not produced in this process. Also,
it’s a good chance to emphasize that the electron-producing pole of a battery is
ALWAYS called the cathode, and the other one is ALWAYS called the anode, and
electrons ALWAYS flow from anode to cathode, whether in a galvanic cell or
electrolytic cell.