Experiment 1
Determination of the Empirical Formula of a Metal Iodide
Objectives
1. To prepare a metal iodide compound
2. To use the method of the gravimetric analysis in the determination the percent
composition and empirical formula of a compound
Theory
Atoms of different elements chemically combine to form different compounds following
two fundamental laws. The Law of constant composition states that a simple of a pure
compound is always composed of the same element combined in the same proportion by mass.
If we analyze sample of water, we will find that the ratio of the masses of hydrogen and oxygen
in each sample is the same, regardless of the source of the water. The law of the multiple
proportions states that if two elements combine to form a series of compounds, the mass ratio of
the element are small whole numbers. For example, hydrogen forms two compounds with
oxygen: water (H2O) and hydrogen peroxide (H2O2). In the first compound, 16.0 grams of
oxygen (one mole of O) combine with 2.016 grams of hydrogen (two moles of H); and in the
second compound, 32.0 grams of oxygen (two moles of O) combine with 2.016 grams of
hydrogen (two moles of H). The ratio of masses of oxygen that combine with the hydrogen in
these compounds is 16.0 g/32.0 g or 1:2, which supports the law of multiple proportion.
When a new compound is prepared, its formula must be determined. This task is most
commonly achieved form the percent composition of the compound-the percent by mass of
each element relative to the total mass of the compound. It is obtained by dividing the mass of
each element in the compound by the molar mass of the compound, and multiplying by 100. In
one mole of water (MW = 18.02 g/mol) are two moles of H (2.016 g) and one mole of O (16.0 g).
Therefore, %H = (2.016/18.016)100 = 11.19 percent and %O = (16.0/18.016)100 = 88.81
percent.
The simplest whole-number ratio of the number of atoms of each element in a
compound is called the empirical formula of the compound. The molecular formula shows the
actual number of each type of atom present in a compound. For example, suppose that the
empirical formula CH is obtained experimentally from the percent composition of a compound.
This formula may represent a molecule having the molecular formula form the empirical formula
one must know the molar mass of the compound.
The empirical formula of the compound can be determined either by chemical analysis
or by synthesis. In chemical analysis, a known mass of the compound is decomposed to obtain
the masses of the elements themselves or some of their known derivatives (for carbon, the
derivative is CO2 and for hydrogen, it is water). From these masses, the percent composition of
the compound is determined. In synthesis, known masses of the elements are allowed to form a
compound and the mass of the compound is determined. Form these data, it is possible to
derive the empirical formula of the compound.
Example 1.1 In an experiment, a mixture of copper and sulfur was heated to produce a sample
of copper sulfide. The following data were obtained:
Weight of the empty crucible
2.077 g
Weight of the crucible + copper
2.289 g
Weight of the crucible + copper sulfide
2.396 g
Calculate the percentages of copper and sulfur in copper sulfide and derive its empirical
formula.
Answer
Copper mass
= 2.289- 2.077 = 0.212 g
Copper sulfide mass
= 2.396- 2.077 = 0.319 g
Sulfur mass
= 2.396- 2.289 = 0.107 g
%Copper in copper sulfide
= (0.212/0.319)100 = 66.5%
% Sulfur in copper sulfide
= (0.107/0.319)100 = 33.5%
Moles of Cu
= 66.5 g (1 mol/63.55 g) = 1.05 mol Cu
Moles of S
= 33.5 g (1 mol/32.0 g) = 1.05 mol S
Empirical formula = CuS
Procedure
Accurately weight the beaker containing the micro stir bar. Place 200 mg of zinc powder
in the beaker, reweigh it, and record the mass. Transfer about 120 to 130 mg of I2 crystals to the
beaker and reweigh the beaker accurately. Calculate the exact amount of iodide taken.
Add 2 to 3 ml of distilled water to the mixture. Place the beaker on a magnetic stirring
hot plate and cover it with a watch glass. With continuous stirring, warm the contents of the
beaker (do not boil). Make sure that iodine does not sublime (vaporize). With in a minute or two,
the aqueous layer will become brown or yellow as a result of the reaction between zinc and
iodine. The reaction is exothermic and may become brisk. If that happens, remove the beaker
form the heat and cool it in an ice-water bath. Continue to stir the mixture (heat if necessary) for
10-15 minutes until the brown color of the aqueous layer almost disappears. Add water to make
up the loss due to evaporation, washing down the side with a few drops of water.
When the aqueous layer become colorless, discard the solution by decantation or using
a dropper to a wasted container No. 4 (making sure that no zinc powder is transferred). Rinse
the unreacted zinc powder in a beaker five or six times with 2 ml of distilled water, decanting the
aqueous layer to a waste container No. 4. While transferring the wash liquid, you must not lose
any zinc. Finally, wash the zinc metal with acetone four or five times, using 1 ml of acetone each
time, decant the solvent, and keep all the zinc in the beaker (acetone solvent discard in waste
container No 4). Place the beaker with zinc in an oven. Heat slowly to dry the powder (to prevent
spattering). When it is dry, cool the beaker in the desiccator. Weight the beaker with the dried
zinc in it. Empty in the beaker by transferring the unreacted zinc powder into a waste container
No. 8 and rinse the beaker.
Do not throw away the leftover zinc; it must be recovered, recycled and reused.
Calculate the amount of zinc that has reacted with the given amount of iodine.
Determine the percent zinc and percent iodine form the data and obtain the empirical formula of
zinc iodine.
Prelaboratory Problems
1. Balance the following reactions:
Fe(s) + Br2(g)
Fe2Br3(s)
C6H6(l) + O2(g)
CO2(g) + H2O(g)
2. A 1.275 g sample of Al was heated strongly in pure oxygen and the combustion product
weighed 2.409 g. Calculate the percent O in the product and find the empirical formula
of the compound formed.
3. Cobalt forms to different chlorides. One contained 54.7 percent Cl and the other 64.4
percent Cl. Determine the empirical formulas of the two compounds.
Postlaboratory Problems
1. A compound contains 90.6 percent lead and 9.40 percent oxygen.
Determine the empirical formula of the compound.
2. Calculate the percent H and percent O in hydrogen peroxide (H2O2)
Reference
Sing, M.M., Pike, R.M., and Szafran, Z., Microscale & Selected Microscale Experiments for
General & Advanced General Chemistry : An Innovative Approach. 1st ed., John Wiley &
Sons, USA, 1995.