its pre lab writing no questions to answer
I need you to write the Purpose of the experiment on your own words. just read its only on sentence. also i need you to read the background information and write it on your on words. Also read Procedure and write on your own words.
Determining the Empirical
Formula of a Compound
Prepared by David P. Dingledy, SUNY Fredonia
PURPOSE OF THE EXPERIMENT
Establish the empirical formula of a chemical compound from the
gravimetric determination of the molar ratio of the reacting masses of
The formula of a chemical compound can be established from experimental
data in accordance with the law of definite proportions. This law may be
stated as follows: In a given compound, the constituent elements are always
combined in the same proportions. This relationship is assumed to exist
regardless of the method of preparation or source of the compound.
Therefore, the formula of a compound can be determined from the
combining masses of the reactants and the mass of the product, as long as
the compound follows the law of definite proportions.
All compounds have an empirical formula, and some also have a molecu-
lar formula. The empirical formula is the simplest whole number ratio of the
relative numbers of atoms of each element in the compound. The molecular
formula gives the actual number of atoms of each element in the compound.
The empirical formula of benzene is CH, which shows that the ratio of
carbon to hydrogen atoms is 1:1. On this basis we would expect the molar
mass of benzene to be
1.01 g H
empirical mass of benzene, g = (1 mol H)
1 mol H
+ (6 mol C
= 13.01 g
1 mol C
However, the molar mass of benzene is found experimentally to be 78.07
g mol-1. This mass is six times that of the empirical formula CH. Therefore,
the molecular formula of benzene must be C6H6, for
1.01 g H
molar mass of benzene, g = (6 mol H)
1 mol H
= 78.07 g
+(1 mol C)
1 mol C
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STOI0423: Determining the Empirical Formula of a Compound
The empirical formula of a compound may be determined by first
finding the mass of each element in a sample of the compound, and then
converting the masses to the number of moles of the elements. These data
can be used to determine ratios of the number of atoms in the compound.
From these ratios, we can obtain the empirical formula.
Consider an experiment in which we are to determine the empirical
formula of the substance formed when tin (Sn) is heated in air. We will
assume that the only reaction occurring is that between the oxygen (O2) in
the air and the Sn. After heating 2.0156 g of Sn in air, we measure the mass
of the product as 2.5588 g. The mass of Sn reacting was 2.0156 8
mass of O reacting is found using Equation 1.
mass of O, g = (mass of product, g) – (mass of Sn, g) = 2.5588 g -2.0156 g = 0.5432 g
The number of moles of Sn reacting is found using Equation 2.
number of moles of Sn, mol
mass of Sn reacting, g
gram – molar mass of Sn, g mol-1
1.698 x 10-2 mol
118.7 g mol-1
The number of moles of O is found using Equation 3.
mass of O reacting, g
number of moles of O, mol =
gram – molar mass of O, g mol-1
3.395 x 10-2 mol
The simplest ratio (x/y) of the number of moles of Sng to the number of
moles of O, is determined using Equation 4.
number of moles of Sn, mol 1.698 x 10-2 mol of Sn 1
number of moles of O, mol 3.395 x 10-2 mol of O 1.999
Therefore, the empirical formula of the product is Sn 1.00001.999 or, rounding
off to the nearest whole number, SnO2.
In this experiment, you will heat in air an accurately measured mass of
magnesium (Mg), forming magnesium oxide and magnesium nitride
(Mg3N2). After you add a small amount of water, the Mg3N2 will convert to
solid magnesium hydroxide, Mg(OH)2, and gaseous ammonia (NH3). As
you heat the Mg(OH)2 further, a solid oxide of magnesium and gaseous
water (H2O) will form.
From the reacting mass of magnesium and its gram-molar mass, you will
calculate the number of moles of Mg reacting. In like fashion, you will cal-
culate the number of moles of O that reacted. You will then reduce the ratio
of the number of moles of Mgr to the number of moles of O, to a small,
whole-number ratio (x/y), which you will express in the empirical formula
Wear departmentally approved eye protection while doing this experiment.
Place a clean, dry crucible with a crucible cover on a Nichrome or clay triangle,
positioning the crucible so that it extends down into the triangle, as shown in
Figure 1. For most efficient heating, position the iron ring so that the bottom of
the crucible is about 1 cm above the inner blue cone of the burner flame.
Observe laboratory regulations concerning the use of an open flame.
Do not touch the heated ring. Handle crucible and cover with crucible tongs.
Heat the crucible and cover for 2 min. Allow the crucible and cover to
cool in air to room temperature.
NOTE: Use crucible tongs to handle the crucible and cover. Do not touch either with
your fingers or hands to prevent transferring skin oils to them, which would affect your
Weigh the crucible and cover to the nearest 0.001 g. Record this mass on
your Data Sheet.
Reheat the crucible and cover, cool, and weigh again. Record this mass
on your Data Sheet.
1992 Cengage Learning
Heating crucible with crucible and cover tilted
Add about 0.3 g of Mg ribbon to the crucible. Crumple the ribbon. Pack
it into the bottom of the crucible using a clean, glass stirring rod. Weigh the
crucible, crucible cover, and Mg to the nearest 0.001 g. Record this mass on
your Data Sheet.
Do not look directly at the burning Mg at the beginning of heating. Avoid
inhaling any fumes or smoke formed during heating.
Place the crucible on the clay triangle as shown in Figure 1, with the crucible
cover tilted slightly open to allow O2 in the air to react with the Mg. Heat gently
at first, then more briskly. If the Mg starts to burn and smoke is given off, cover
the crucible completely and remove the flame for a short time. When the
contents no longer glow bright red on heating and smoke is not given off, the
reaction is complete. At this point, briskly heat the crucible for several minutes
with the cover tilted open, until the bottom of the crucible glows a dull red.
Avoid inhaling any dust from the residue formed in this experiment.
Allow the crucible and contents to cool. Carefully push the fluffy
product down into the bottom of the crucible with a clean, glass stirring rod.
Scrape the particles adhering to the rod into the crucible. Add 6 to 8 drops of
distilled or deionized water from a medicine dropper. Note the odor of NH3.
Heat the crucible with the cover in place, gently at first, and then more
briskly until the crucible glows dull red. Continue heating for 5 min.
Cool and weigh the crucible and contents. Record this mass on your
Reheat, cool, and reweigh until the mass is constant to within 0.02 g.
Record this constant mass on your Data Sheet.
Repeat the entire procedure twice more with new samples of Mg, if
there is time.
Wash your hands thoroughly with soap or detergent before leaving the