Chemical Formulas
Name: Andrew Yaksic
Purpose: To become familiar with chemical formulas and how they are obtained.
Equipment: analytical balance, Bunsen burner, 50-mL graduated cylinder, wire gauze, crucible, crucible cover, glass beads, 250-mL beaker, evaporating dish, ring stand, iron ring, stirring rod, clay triangle, crucible tongs, beaker tongs
Materials: granular zinc, powdered sulfur, copper wire, 6 M HCl, water
Introduction: Chemical formulas are used to indicate how new compounds are formed by chemical reactions. These formulas are composed of chemical symbols. One or two letters (with the first letter capitalized) symbolize a chemical element.
Many elements are found in nature in molecular form. Two or more of the same type of atom (e.g. two oxygen atoms) are tightly bound together. The molecule behaves in many ways as a single distinct unit. A diatomic oxygen molecule would be represented by the chemical formula O2. The subscript in the formula represents the two oxygen atoms present in each oxygen molecule.
Chemical formulas that indicate the actual numbers and types of atoms in a molecule are called molecular formulas, whereas chemical formulas that indicate only the relative numbers of atoms of a type in a molecule are called empirical formulas. The subscripts in an empirical formula are always the smallest whole number ratios.
Some other important terms for this experiment are atomic mass, atomic weight, formula weight, molecular weight, and mole. Atomic mass defines the masses of individual atoms. Atomic mass units are used instead of grams. One atomic mass unit is 1.66054x10-24 grams. Atomic weight is the average atomic mass of each element. The formula weight of a substance is the sum of the atomic weights of all atoms in its chemical formula. If the chemical formula of a substance is its molecular formula, then the formula weight is also called the molecular weight. Ionic substances have only formula weights. A mole represents 6.02x1023 atoms. The mass of 1 mole (abbreviated mol) of a substance is called its molar mass. The molar mass of any substance is numerically equal to its formula weight.
Taking a ratio of molar masses will yield an empirical and molecular formula. For example, if there are two moles of hydrogen per mole of oxygen in a given compound, then the empirical formula is H2O. The experimentally determined molecular weight can then be used to determine how many times the empirical formula must be multiplied. For example, if the experimentally determined molecular weight for H2O is 36.0 g, then the molecular formula of the substance in question must be H4O2, because the molecular weight of H2O is 18.0 g.
Figure 5.1 depicts a steam bath. A steam bath will be used to heat the contents of an evaporating dish. This indirect method of heating is better than a direct method of heating because it prevents the contents of the evaporating dish from getting too hot and spattering out of the dish. This method is as effective as direct heat without any of the risks. The steam bath is depicted below:
Figure 5.2 depicts the setup for the crucible heating. The crucible is placed onto a clay triangle and heated. This holds the crucible securely while still allowing for the flame to directly contact the crucible. This setup will be used to determine the chemical formula of copper sulfide. The crucible heating setup is depicted below:
In this experiment, two compounds will be analyzed for chemical composition: zinc chloride and copper sulfide. The compounds will be broken down into their base components and molar ratios will provide the chemical formulas.
Procedure:
A. Zinc Chloride
1. The mass of a clean, dry evaporating dish was measured and recorded.
2. A sample of granular zinc was obtained. Approximately 0.5 g was added to the weighed evaporating dish. The mass of the evaporating dish and zinc was measured and recorded.
3. 15 mL of 6 M HCl was slowly added to the evaporating dish containing the zinc. The
mixture was stirred during the addition of the acid. No flames were permitted in the laboratory during this part of the procedure. An additional 5 mL of acid was added to the mixture if any zinc remained undissolved.
4. A steam bath was set up by placing a 250-mL beaker on an iron ring and wire gauze. Two glass beads were placed into the beaker. The evaporating dish was placed on top of the beaker.
5. The beaker was heated to create steam. The beaker was heated until most of the liquid from the evaporating dish had disappeared. The evaporating dish was then removed from the steam bath and placed directly on the wire gauze. The evaporating dish was then heated until all of the liquid appeared to be gone.
6. The dish was placed on a cooling pad and allowed to cool to room temperature. The mass of the evaporating dish was measured and recorded. The dish was then reheated gently. The mass was measured and recorded again. The dish was reheated gently for a third time. The mass was measured and recorded.
B. Copper Sulfide
1. The mass of a clean, dry crucible with cover was measured and recorded.
2. Approximately 1.5 g to 2.0 g of tightly wound copper wire was placed in the crucible. The mass of the copper, crucible, and cover was measured and recorded.
3. In a fume hood, enough powdered sulfur to cover the copper wire was placed into the crucible. The crucible and cover was placed onto a clay triangle supported by an iron ring and ring stand.
4. The crucible was heated until the sulfur stopped burning. The crucible was heated further for about 5 minutes to ensure that all of the sulfur had been burned. The crucible was
removed from the clay triangle and was placed onto a cooling pad. The crucible was allowed to cool to room temperature. The mass of the crucible with cover was measured and recorded.
5. Steps 3 and 4 were repeated twice.
Observations:
A. Zinc Chloride
Mass of evaporating dish and zinc: 41.1807 ± .0001 g
Mass of evaporating dish: 40.7205 g ± .0001 g
Mass of evaporating dish and zinc chloride: Trial 1: 41.7034 ± .0001 g
Trial 2: 41.7091 ± .0001 g
Trial 3: 41.7132 ± .0001 g
B. Copper Sulfide
Mass of crucible, cover, and copper: 18.0721 ± .0001 g
Mass of crucible and cover: 16.5402 ± .0001 g
Mass of crucible, cover, and copper sulfide: Trial 1: 18.4708 ± .0001 g
Trial 2: 18.4987 ± .0001 g
Trial 3: 18.5132 ± .0001 g
Results:
A. Zinc Chloride
Mass of zinc (mass of evaporating dish and zinc – mass of evaporating dish) =
41.1807 ± .0001 g – 40.7205 ± .0001 g =
.4602 ± .0002 g
Mass of zinc chloride (trial 3 mass of evaporating dish and zinc chloride – mass of evaporating dish) =
41.7132 ± .0001 g – 40.7205 ± .0001 g =
.9927 ± .0002 g
Mass of chlorine in zinc chloride (mass of zinc chloride – mass of zinc) =
.9927 ± .0002 g – .4602 ± .0002 g =
.5325 ± .0004 g
Empirical formula for zinc chloride:
Moles of zinc (mass of zinc * (moles zinc per gram zinc)) = (.4602 ± .0002 g Zn * (1 mol/65.4 g)) =
(.4602 ± .04% g Zn * (1 mol/65.4 g)) = (.007037 ± .04% mol Zn) =
.007037 ± .000003 mol Zn
Moles of chlorine (mass of chlorine *(moles chlorine per gram chlorine))= (.5325 ± .0004 g Cl * (1 mol/35.5 g)) =
(.5325 ± .08% g Cl * (1 mol/35.5 g)) = (.01500 ± .08% mol Cl) =
.01500 ± .00001 mol Cl
Moles chlorine to moles zinc ratio: .0150/.007037 : .007037/.007037
2.13:1 ratio
Balanced chemical equation for the formation of zinc chloride from zinc and HCl: Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
B. Copper Sulfide
Mass of copper (mass of crucible, cover, and copper – mass of crucible and cover) =
18.0721 ± .0001 g – 16.5402 ± .0001 g =
1.5319 ± .0002 g
Mass of copper sulfide (trial 3 mass of crucible, cover, and copper sulfide – mass of crucible and cover) =
18.5132 ± .0001 g – 16.5402 ± .0001 g =
1.9730 ± .0002 g
Mass of sulfur in copper sulfide (mass of copper sulfide – mass of copper) =
1.9730 ± .0002 g – 1.5319 ± .0002 g =
.4411 ± .0004 g
Empirical formula for copper sulfide:
Moles of copper (mass of copper * (moles copper per gram copper)) = (1.5319 ± .0002 g Cu * (1 mol/63.5 g)) =
(1.5319 ± .01% g Cu * (1 mol/63.5 g)) = (.023423 ± .01% mol Cu) =
.023423 ± .000002 mol Cu
Moles of sulfur (mass of sulfur *(moles sulfur per gram sulfur))= (.4411 ± .0004 g S * (1 mol/32.1 g)) =
(.4411 ± .09% g S * (1 mol/32.1 g)) = (.01374 ± .09% mol S) =
.01374 ± .00001 mol S
Moles copper to moles sulfur ratio: .023423/.01374 : .023423/.023423
1.70:1 ratio
Balanced chemical equation for the formation of copper sulfide from copper and sulfur:
16Cu(s) + S8(s) → 8Cu2S(s)
Discussion: The true value of the ratio of moles chlorine to moles zinc ratio is 2. The true value of the ratio of moles copper to moles sulfur is 2. Thus, two percent error calculations are in order.
Percent error, moles chlorine to moles zinc
(observed – true)/true * 100 (2.13 – 2)/2 * 100 =
6.50% error.
Percent error, moles copper to moles sulfur
(observed – true)/true * 100 (1.70 – 2)/2 * 100 =
-15.00% error.
These are fairly high values for percent error. There are many sources of error in this experiment. If any liquid remained in the zinc chloride sample, then the mass reading would have been inaccurate, as water would be factored into the mass. If the evaporating dish was not completely dry, then water would be factored into the mass. If the sulfur had not been fully removed from the crucible before weighing it, then the mass of the sulfur would have changed the reading. If any of the samples were spilled or allowed to evaporate, then mass would have been lost. Anything that could have caused the masses of the samples to be too high or too low is a source of error in this experiment.
The theory associated with this experiment is the atomic theory. This experiment proves Dalton’s law of definite proportions, which states, “In a given compound, the relative numbers and kinds of atoms are constant.” Every sample of copper sulfide or zinc chloride will have a constant molar ratio due to the atomic makeup of the compound. Every molecule contains a definite number of atoms, whether they are 2 copper atoms and 1 sulfur atom or 2 chlorine atoms and 1 zinc atom.
The ramifications of this experiment are far-reaching. Familiarity with formulas was gained. Laboratory skills were improved. A steam bath was used for the first time. Proper use of a crucible was practiced. In industry, the procedures used in this experiment can be used to determine the chemical composition of any compound. If it can be decomposed into its base parts, then the chemical formula for any compound can be determined. A simple ratio of molar content yields information about any compound. For example, scientists at NASA can analyze compounds taken from objects in space. Marine biologists can determine what elements are necessary for the diet of whales so that whales in captivity can be fed and cared for properly. Scientists from all industries or causes can
use these procedures for practical applications.
Questions:
1. The molecular formula of a substance can be determined from its percent composition. The percents can be used to determine the number of moles in a 100-gram sample, which can be used to determine the molecular formula.
2. The formula of zinc chloride is ZnCl2. (65.4 g + 35.5 g + 35.5 g ≈ 136.28 g)
3. The atomic weights of zinc or copper could be determined by the methods used in this experiment. The number of moles of substance could be determined from the mass of the substance, and the number of moles could be used to determine grams per mole.
4. How many grams of zinc chloride could be formed from the reaction of 9.76 g of zinc with excess HCl?
9.76 g Zn (1 mol Zn/65.4 g Zn)(1 mol ZnCl2/1 mol Zn)(136.4 g ZnCl2/1 mol ZnCl2) =
20.4 g ZnCl2
5. How many kilograms of copper sulfide could be formed from the reaction of 2.70 mol of copper with excess sulfur?
2.70 mol Cu (1 mol Cu2S/2 mol Cu)(159.1 g Cu2S/1 mol Cu2S)(1 kg/1000 g) =
.215 kg Cu2S
6. First reaction: 2Cu2S + 3O2 → 2Cu2SO3
Second reaction: Cu2SO3 → Cu2O + SO2
Conclusion: The experiment was completed successfully. Familiarity with chemical formulas and how they are obtained was gained. The results were not as accurate as possible, but they were close enough to substantiate the theory associated with this experiment.