Lab #19 – EM of a metal by Electrolysis
Name
Background
There are precious few ways to count moles in the chemistry lab, and counting electrons is probably the
cheapest and easiest. By recognizing that an amp is a coulomb/second and a mole is 96,485 coulombs of e-‘s,
one can count electrons with a stopwatch and an ammeter. Of course, this technique is limited to
electrochemical (redox) reactions in which electrons are transferred.
In this lab, though, instead of counting electrons directly, we will track the transfer of e-‘s by measuring the
quantity of one of the products produced in the reaction . You will use a power supply to force the oxidation of
an unknown metal at the anode of an electrochemical cell, and H+ ions will be reduced at the cathode. As you
know, the reduction of H+ produces H2 gas, so we will collect the gas produced and use the ideal gas law to
calculate moles of gas. By considering the hydrogen half-reaction we can then determine the number of e-‘s
that have been transferred.
Since our metal is unknown, we cannot say with certainty how many electrons are lost by each atom of metal.
We encountered this problem in acid-base titration when we were uncertain how many moles of electrons were
provided by our unknown acid. We will solve the problem the same way we did with the unknown acid;
instead of determining the molecular weight of our unknown metal we will determine its equivalent mass, the
mass of metal oxidized during the loss of 1 mole of electrons. As before, the molar mass is an integral multiple
of the equivalent mass, depending on the charge of the ion formed by the oxidation of the metal.
Procedure
Obtain an unknown metal, a piece of sandpaper, about 150 mL of the electrolyte (0.1 M acetic acid in 0.5 M
Na2SO4), an eudiometer, a copper wire electrode, 2 alligator clip wires, and a power supply. Record the UK
number. Use the sandpaper to lightly sand the surface of the metal. Do NOT sand the surface of the benchtop,
or any other inappropriate surface. Use paper towel to wipe off any loose material from the metal and weigh it
on the analytical balance. Record the mass of the metal before reaction.
Set up the apparatus as demonstrated by the instructor. Attach the power supply so that the unknown metal is
the anode. Gradually increase the power until a reasonable rate of bubbling occurs in the eudiometer. Continue
the reaction until just under 50 mL of gas have been produced. Equalize the pressure inside and outside of the
eudiometer by moving the eudiometer until the liquid level inside the eudiometer is about even with the liquid
level outside. Read and record the volume of gas produced. Record the temperature of the electrolyte
solution. Record the barometric pressure. If time permits, refill the eudiometer with fresh electrolyte and
repeat the electrolysis.
Once electrolysis is complete, clean the unknown metal electrode: Rinse with 0.1 M acetic acid, then use your
fingers to rub off any loose material. Rinse your hands and the electrode thoroughly with water. Rinse the
metal with acetone and allow it to dry completely. Weigh the dry metal on the analytical balance and record
the mass.
Prelab summary
Procedural Tips: Clean the electrode thoroughly before and after, following directions.
Do not go past the bottom graduation on the eudiometer.
Read the eudiometer when the liquid level inside is nearly even with the liquid level outside.
Safety: Blown fuses will be replaced at the technician’s rate of $90/hour.
Disposal: Return the clean and dry metal unknowns.
The electrolyte solution can be rinsed down the sink.
Data and Calculations
Mass of UK metal before oxidation
g
Mass of UK metal after oxidation
g
Mass of metal oxidized by reaction
g
Total volume of H2 produced
mL
Temperature
°C
Barometric pressure
torr
Calculate moles H2 produced (show work)
moles H2
Calculate moles of electrons transferred (show work)
moles electrons
Equivalent mass of the UK metal #
g
Prelab Questions
1. Write the anode half reaction and the cathode half-reaction, assuming the metal forms a +1 cation.
2. Write the balanced equation for the reaction.
3. Considering the half-reactions in #1 above, what is the equivalent mass of H2, the mass of H2 produced by
the transfer of 1 mole of electrons? Justify your answer.
4. In an electrolysis experiment similar to this one, a student collects 47.22 mL of H2 gas that is formed at the
cathode, while the metal anode loses 0.1175 g of mass. The lab was run at 25.0°C and barometric pressure is
745 mm Hg. The vapor pressure of water at 25°C is 23.8 mm Hg. Use this data to determine the equivalent
mass of the unknown metal. Show all work.
5. If the metal forms a +2 cation, what is the molar mass of the metal?
6. If the metal forms a +3 cation, what is the molar mass of the metal?