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LAB EQUIVALENT MASS BY ELECTROLYSIS
Adapted from Chemical Principles in the Laboratory, by Slowinski, Wolsey, and Masterton
A careless AP Chem teacher placed a large number of electrodes into storage without labeling them. The
electrodes are made of shiny, malleable silver-colored metal, and are either Sn or Zn. Your assignment is to
determine the identity of the metal electrodes.
Experimentally we find that the equivalent mass of a metal electrode can be determined by electrolysis. As
you know, some liquids, because they contain ions, will conduct an electric current. If the two terminals on a
storage battery, or any other source of DC voltage, are connected through metal electrodes to a conducting
liquid, an electric current will pass through the liquid and chemical reactions will occur at the two metal
electrodes; in this process electrolysis is said to occur, and the liquid is said to be electrolyzed.
At the electrode connected to the cathode (the negative pole of the battery in an electrolysis), a reduction
reaction will invariably be observed. In this reaction electrons will usually be accepted by one of the species
present in the aqueous solution. The species reduced will ordinarily be a metallic cation or the H+ ion or possibly
water itself; the reaction that is actually observed will be the one that occurs with the least expenditure of
electrical energy, and will depend on the composition of the solution. In the electrolysis cell we shall study, the
reduction reaction of interest will occur in a slightly acidic medium; hydrogen gas will be produced by the
reduction of hydrogen ion:
2 H+(aq) + 2 e- → H2(g)
(1)
In this reduction reaction, which will occur at the negative pole, or cathode, of the cell, the production of
one mole of H2(g) requires two moles of electrons. A mole of electrons is called a faraday, after Michael
Faraday, who discovered the basic laws of electrolysis. The amount of a species which will react with a mole of
electrons, or one faraday, is equal to the equivalent mass of that species. To form one mole of H2(g) one would
have to pass two faradays through the electrolysis cell.
In this electrolysis experiment we will measure the volume of hydrogen gas produced under known
conditions of temperature and pressure. By using the Ideal Gas Law we will be able to calculate how many
moles of H2 were formed, and hence how many faradays of electricity passed through the cell.
At the positive pole of an electrolysis cell (the metal electrode that is connected to the (+) terminal of the
battery), an oxidation reaction will occur, in which some species will give up electrons. This reaction, which
takes place at the anode in the cell, may involve again an ionic or neutral species in the solution or the metallic
electrode itself. In the cell that you will be studying, the pertinent oxidation reaction will be that in which a
metal under study will participate:
M(s) → Mn+(aq) + n e
−
(2)
During the course of the electrolysis the atoms in the metal electrode will be converted to metallic cations and
will go into the solution. The mass of the metal electrode will decrease, depending on the amount of electricity
passing through the cell and the nature of the metal. To oxidize one mole, or one molar mass, of the metal, it
would take n faradays, where n is the charge on the cation that is formed. By definition, one faraday of electricity would cause one equivalent mass, EM, of metal to go into solution. The molar mass, MM, and the equivalent mass of the metal are related by the equation:
MM = EM × n
(3)
In an electrolysis experiment, since n is not determined independently, it is not possible to find the molar mass
of a metal. It is possible, however, to find equivalent masses of many metals. In our situation, since we know
the metal is either Sn or Zn, we know that n = 2.
The general method we will use is implied by the discussion. We will oxidize a sample of an unknown
metal at the positive pole of an electrolysis cell, weighing the metal before and after the electrolysis and so
determining its loss in mass. We will use the same amount of electricity, the same number of electrons, to
reduce hydrogen ion at the negative pole of the electrolysis cell. From the volume of H2 gas that is produced
under known conditions we can calculate the number of moles of H2 formed, and hence the number of faradays
that passed through the cell. Knowing that the metal is either Sn or Zn, the molar mass of the metal is then
calculated.
Experimental Procedure
Obtain a eudiometer and a sample of metal unknown. Lightly sand the metal to clean it. Rinse the metal with
0.1 M HC2H3O2 (acetic acid), then water, and then acetone. Let the acetone evaporate. When the sample is dry,
weigh it on the analytical balance to 0.001 g.
eudiometer
Set up the electrolysis apparatus as indicated to the left.
There should be about 100-125 mL 0.5 M HC2H3O2 in 0.5 M
Na2SO4 in a 250-mL beaker. This will serve as the conducting
solution. Completely fill a 50-mL eudiometer with the
HC2H3O2/Na2SO4 solution, insert a rubber stopper, and invert
inside the beaker. Remove the stopper. Insert the bare coiled end
of the heavy insulated copper wire (be sure to sand it first) up into
the end of the eudiometer; all but the tip of the wire should be
covered with watertight insulation. Record the level of solution
in the eudiometer (if you have a trapped air bubble).
The metal unknown will serve as the anode in the electrolysis
cell. Connect the metal to the (+) pole of the power source with
an alligator clip and immerse the metal but not the clip in the
conducting solution. The copper electrode will be the cathode in
the cell. Connect that electrode to the (−) pole of the power
source. Hydrogen gas should immediately begin to bubble from
250-mL beaker
the copper cathode. Collect the gas until about 50 mL have been
produced. At that point, stop the electrolysis by disconnecting the
copper electrode from the power source. Wait until any stray
hydrogen bubbles have risen, then record the level of the liquid in the eudiometer. Measure and record the
temperature and the barometric pressure in the laboratory. In some cases a cloudiness may develop in the
solution during the electrolysis. This is caused by the formation of a metal hydroxide, and will have no adverse
effect on the experiment.
Take the alligator clip off the metal anode and wash the anode with 0.1 M HC2H3O2 (acetic acid). Rub off
any loose adhering coating with a paper towel. Rinse the electrode in water and then in acetone. Let the acetone
evaporate. Weigh the dry metal electrode to the nearest 0.001 g.
If time permits, do a second trial. You may reuse the same solution, but be sure to sand both electrodes.
Determine whether your metal is Sn or Zn. Calculate the percent error for your molar mass calculation.
DISPOSAL OF REACTION PRODUCTS. When you are finished with the experiment, return the metal
electrode to the teacher and discard the conducting solution down the drain.