Lab #18 - Oxides of Metals and Non-Metals
|Introduction: An electrochemical cell is constructed
from two-half cells. One half cell contains both the oxidized and reduced
form of the oxidizing agent. The other half-cell contains the corresponding
forms of the reducing agent. The half-cells are connected by means
of a salt bridge or a porous container filled with an inert material through
which ions can pass. A wire completes the external circuit through
which electrons can flow. The cell voltage is calculated using 1.0
M solutions. In this experiment, you will construct several cells to
determine their Eo or electromotive force.
|Problem: Construct and then compare
several electrochemical cells for their Eo force.
|Apparatus: safety goggles, test
tubes, 100 mL beakers 250 mL beakers bunsen burner, u-tube, steel wool, galvanometer
|Materials: Strips of metals, Cu, Zn, Pb and
Mg and others. 1.0 M solutions of each of the above metals in their nitrate
form. Concentrated NH3 solution. 1.0 M NaCl solution,
|Safety Issues: Concentrated ammonia
solution is very irritating to the eyes. Wear safety goggles and avoid
inhaling the vapour. Lead compounds are poisonous. Avoid spills
and dispose of these materials properly in the hazardous waste containers.
||After reading through this procedure, create a data table to record
your results. Make sure that you cover all possible combinations of
||Clean the metal electrode with sandpaper or steel wool, and set up
the apparatus as shown above.
||Record the voltage of the cell. If the needle shows a negative
deflection in the reading, reverse the connections at the terminals.
||Write balanced equations for each of the spontaneous cell reactions.
||From a handbook of chemistry, look up the reduction potential of
each of the half cells used in your experiment. Calculate the
Eo from these handbook potentials and compare them to your actual
||List some of the possible sources of error.