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CENTRE CHE 131 - Experiment 3A: Formula of a Compound

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Experiment 3A: Formula of a CompoundIn classifying matter, it is important to recognize the distinction between a chemical compound (a puresubstance) and a mixture. Mixtures may have variable composition, whereas a chemical compound hasa fixed composition. In a mixture the separate components retain their own identities and properties,whereas a compound has a definite set of properties different from those of the constituent elements.For example, the compound water is a liquid at room temperature, even though the constituent elementshydrogen and oxygen are both gases at the same temperature. Also, every sample of pure water hasthe elements hydrogen and oxygen in a 1:8 mass ratio. An important stage in the history of chemistry was the investigation of the proportions in which elementscombined into compounds. The recognition that there was a definite mass ratio between any twoelements in a compound was only arrived at after many careful preparations and analyses ofcompounds. This principle was discovered by the French scientist Joseph Proust and came to be knownas the Law of Definite Proportions (or Law of Constant Composition). In this exercise, the student will determine the formula of a metal sulfide formed by the reaction of thepure metal with elemental sulfur. A known amount of the metal is combined with an excess of sulfur,and the mixture is heated in a crucible to form the compound (equation 1). x M(s) + y S(s) 6 MxSy(s) (1) Any excess uncombined sulfur that may remain after the reaction then proceeds to react withatmospheric oxygen to produce gaseous SO2, which is driven off. S(s) + O2(g) 6 SO2(g) (2) Note that the metals used are nonvolatile under the conditions of the experiment. The mass of the sulfurthat combined with the metal may be found by weighing the compound formed. In order to assure thatcomplete reaction has occurred, it is generally advisable to conduct a second heating with additionalsulfur. An empirical formula is simply the ratio of the combining masses re-expressed in moles. The studentmay then calculate the empirical formula of the compound by converting the mass of each element inthe compound to moles and expressing the ratio in whole numbers. PROCEDURE1. Clean the porcelain crucible by scrubbing with soap and water. Rinse the crucible and lid withdistilled water. Some stains may remain in the porcelain after this treatment, but they will notinterfere with the experiment. 2. Support the clean crucible and lid on a clay triangle with the lid slightly ajar (Figure 7) and heat overa Meker burner until the bottom of the crucible is red hot. Allow the crucible and lid to cool to roomtemperature and weigh them to the nearest 0.001 gram. Handle the crucible and lid with crucibletongs at all times. (Fingerprints add mass). 3. Obtain from your laboratory instructor a metal sample and record the identification code number, ifprovided. Place approximately 1 gram of metal sample into the crucible and weigh the crucible, lid,and metal to the nearest 0.001 gram. 4. Weigh 1 gram of sulfur into a plastic weighing boat. Add the sulfur to the metal in the crucible andmix the two elements well.2Figure 7. Heating crucible and crucible cover5. Place the covered crucible and contents on the clay triangle and begin heating gently with the gasburner. Continue heating gently for approximately 15 minutes. This heating and the subsequentoperation should be conducted in the hood due to the obnoxious nature of the sulfur dioxide that willbe formed. 6. When the reaction is complete (sometimes signified by a "poof" and the cover of the cruciblejumping), move the flame around the crucible and lid to ensure complete oxidation and volatilizationof the excess sulfur. Finally, with the lid slightly ajar, increase the heat until the bottom of the cruciblebecomes a dull red. Maintain this heat for approximately five minutes, remove the heat and allow thecrucible and contents to cool to room temperature. There should be no evidence of elemental sulfurin the crucible (yellow or gummy red material). 7. Add an additional gram of sulfur, mix this with the reaction mixture, and repeat the heating steps (5and 6). Cool to room temperature and weigh the crucible and contents. Using the mass data recorded, calculate the mass of sulfur that reacted with the mass of metal used inyour experiment. Using the molar masses for these elements, calculate the empirical formula of themetal sulfide that you prepared. Express the numbers of moles to the appropriate number of significantfigures (consistent with your data), then convert to a whole-number ratio.If time permits, the student may want to heat the metal and sulfur separately to observe what happens toeach when heated. This may explain some observations of the reaction mixture. POST-LAB QUESTIONS: After discussing your result, answer these questions.1. Does the amount of sulfur in your compound appear to be too low or too high? By what percentage? Postulate reasons for this systematic error. 2. A few binary (i.e., two element) compounds have more than one ratio of elements. An examplewould be water (H2O) and hydrogen peroxide (H2O2). What experimental steps would you suggest todetermine if the metal sulfide you prepared forms more than one compound? 3. Suppose that you were trying to prepare a binary metal oxide and determine its empirical formula. Outline in reasonable detail an experimental procedure to form a metal oxide and determine itsformula.3Experiment 3B: Formula of a Hydrate In this experiment you will again use gravimetric methods (weighing) to determine a chemical formula. Many inorganic salts exist as hydrates, in which a specific number of water molecules is present. Forexample, gypsum is calcium sulfate dihydrate, CaSO4A2H2O. The water of hydration can be removed byheating strongly in a burner flame; the remainder of the compound does not volatilize. After removal ofthe water, the salt is referred to as an anhydrous salt. The mass difference before and after heating canbe related to the amount of water of hydration. The compound to be studied is hydrated copper(II) sulfate (or cupric sulfate), CuSO4AxH2O. You willdetermine the number of water molecules x. This reaction is especially convenient to study, since thehydrated and anhydrous forms have distinctly different colors. CuSO4AxH2O 6 CuSO4 + x H2O (3) Experimental 1.


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