Gram Equivalent Weight of an Unknown Metal

Discussion

Many reactions occur in which the oxidation states of the reactants and products change. These are known as “redox” reactions. An electrochemical cell is a device constructed to separate a redox reaction into two half reactions, the oxidation and reduction half reactions. The reaction can still occur, but the electrons will be transferred indirectly from the reducing agent to the oxidizing agent via an external circuit. For a thermodynamically favored reaction, i.e. spontaneous, the reaction will proceed of its own accord. These types of cells are called galvanic or voltaic cells but are more commonly known as batteries. The natural thermodynamic driving force of the reaction transferring electrons can be harnessed to provide energy and do work.

Reactions which do not spontaneously occur can be “forced” to occur in a related type of electrochemical cell called an electrolytic cell. In this case external electrical energy is used to “force” an electron transfer reaction to occur which would not proceed on its own, i.e. thermodynamically unfavored. Work is done by the supplied electrical energy to force a nonspontaneous reaction (composed of two half reactions) to occur.

The transfer of electrons through an external circuit is common to both types of electrochemical cells. The number of electrons transferred while the reaction is proceeding is directly related to the amount of reactants converted to products. For redox reactions involving metals, the gram equivalent weight (GEW) of a metal is the weight of metal which reacts with or releases one mole of electrons (a Faraday). Therefore, the GEW is the mass of metal per mole of electrons transferred. Notice that one Faraday is also used to reduce one mole of hydrogen cations to half a mole of hydrogen gas according to the following reaction:

H⁺(aq) +  e ^- -----> 1/2 H ₂(g)

An unknown metal can be identified by its GEW as long as the amount of metal reacted and the moles of electrons transferred are known. By measuring the volume of hydrogen gas generated in one half reaction, the moles of electrons transferred from the unknown metal in the corresponding oxidation half reaction can be determined. (Question: How can this be done?) The two half reactions are interrelated by the Faradays of electricity passing from the metal to the hydrogen cations. For a generic, unknown metal the oxidation half reaction can be written as follows:

M(s) -----> M ⁿ⁺(aq) + n e ^-

The following electrochemical apparatus can be used to find the GEW of an unknown metal. At the anode, oxidation of the unknown metal occurs; at the cathode, hydrogen cations (hydronium ions) are reduced to hydrogen gas. Note that the copper strip joining the two beakers functions as an “electron bridge”. Normally, a salt bridge which allows passage of ions between the two solutions is used to maintain charge balance. However, an “electron bridge” incorporates a second redox reaction which effectively accomplishes the same result. Copper(II) ions are produced (by oxidation) from the bridge to replace the hydrogen cations being converted to hydrogen gas as evidenced by the blue color generated in this solution. Simultaneously, water is initially being reduced at the opposite end of the copper bridge to form hydroxide ions and hydrogen gas. The hydroxide ions balance the charge of the unknown metal cations being generated and often cause precipitation of an insoluble, white metal hydroxide. Record observations specific to your experiment.

Equipment :

Procedure

1. Students will work in pairs. Obtain the buret (an expensive item) and your unknown metal sample from the Stockroom. Weigh the metal and accurately record its mass in your notebook.
   
2. Set up the apparatus as illustrated in the diagram. Ensure that all of the exposed copper wire (wire without insulation) is inside the buret. Also make sure that the buret stopcock valve does not leak. Otherwise, some hydrogen gas will escape from the buret and hinder accurate analysis.
   
3. Connect the anode (unknown metal) to the positive terminal at your desk. Connect the the wire going inside the buret to the negative terminal. Have your instructor check your apparatus and connections.
   
4. Put 180 mL of distilled water in each of two 250 mL beakers. Add 60 mL of 2M KNO3 to one beaker of water and 60 mL of 2M H2SO4 to the other beaker of water. Sulfuric acid will burn your skin if not washed off with water immediately. Always wear eye protection. Lower the buret into the beaker containing the sulfuric acid solution, place the beaker of potassium nitrate solution next to it, and put one end of the copper strip into each beaker. Place the unknown metal electrode in the potassium nitrate solution so that it does not touch the copper strip and the electrode clip does not contact the solution.
   
5. You are now ready to fill the buret with sulfuric acid solution using the aspirator and to begin your electrolysis. Your instructor will describe the proper way to fill the buret and use the direct current (DC) power source. Extreme caution should be exercised not to touch the wires, the copper strip bridge or the electrolyte solutions.
   
6. You will need to collect approximately 150 mL of hydrogen gas for each trial. This is accomplished by collecting a series of three accurately known (and recorded) volumes of hydrogen gas successively in the buret without disturbing the unknown metal electrode. You will need to turn the power off before you reach the 50 mL mark of gas in the buret, accurately read the volume of gas generated, refill the buret with acid, and repeat this procedure until a total of approximately 150 mL of gas has been generated. Calculate the total volume of hydrogen gas produced under your experimental conditions by combining your three measured volumes. (What experimental conditions are important to consider when collecting a gas over an aqueous solution? Recall what you are trying to determine.)
   
7. With the power off, remove the metal unknown, wipe off any deposits, wash it well with distilled water, and finally rinse it with acetone. Allow the metal to dry for five minutes and then weigh the metal. Record the mass and calculate the mass loss for the trial. At least three trials should be completed for a reasonable statistical analysis, but complete as many as time allows.
   
8. Calculate the GEW of your unknown metal from the data you have gathered. The moles of electrons transferred in your redox reaction can be determined from the volume of hydrogen gas generated. You will need to consider any barometric pressure correction and water vapor pressure correction at the temperature of your experiment (measured room temperature).  The mass of unknown metal reacted is obtained by the difference in mass you measured.
   
9. To identify your unknown metal, examine the table of standard reduction potentials listed below. This table provides the common oxidation states of these possible metals. Calculate the GEW for one mole of each of the metals listed and report their values. You can use the periodic table to find the molar mass of each metal.  Now, compare the experimentally determined GEW of your unknown metal to the actual values you calculated and  identity your unknown.
   
10. If a salt bridge had been used instead of an “electron bridge” (copper strip bridge) would your unknown metal spontaneously reduce hydrogen cations to hydrogen gas under standard conditions? Calculate this cell potential at standard conditions and report your findings. Show the half reactions and the net cell reaction. Would this be a galvanic or electrolytic cell?
    
11. Next, consider your experimental electrochemical cell including the “electron bridge” at standard conditions.  Based on your experimental observations, write all four half reactions and the net cell reaction.  Then using the standard reduction potentials provided in the table below, calculate the standard potential for this cell. Does it differ from the cell with a salt bridge (step 10)? How?  

Standard Reduction Potentials

Revised by Shane Phillips (12/2/98)