Moles and Chemical Formulas 11 LABORATORY GOALS Determine the simplest formula of a compound. Calculate the percent water in a hydrate. Determine the formula of a hydrate. LAB INFORMATION Time: Comments: Related Topics: 2-2 1/2 h Tear out the report sheets and place them beside the matching procedures. Use steel wool to remove any coating on the magnesium ribbon until it is shiny. Check the crucible for cracks before you start to heat it. When you set a hot object aside to cool, remember that it is hot. Dispose of all chemicals as directed by your instructor. Formulas, moles, molar mass, calculating moles from grams, calculating grams from moles, hydrates, dehydration CHEMICAL CONCEPTS A. Finding the Simplest Formula The simplest formula of a compound is the lowest whole-number ratio of the atoms in the formula. For example, the compound benzene, with molecular formula C 6H6, has the simplest formula CH. Some molecular formulas and their simplest formulas are shown in Table 11.1. TABLE 11.1 Examples of Molecular and Empirical Formulas Name Molecular Formula Simplest Formula Acetylene C2 H2 CH Benzene C6H6 CH Ammonia NH 3 NH 3 Hydrazine N2H4 NH 2 The simplest formula of a compound is determined by converting the number of grams of each element to moles and finding the lowest whole-number ratio to use as subscripts. For example, in an experiment it was determined that 0.040 mole of Zn had combined with 0.080 mole of Cl to form a compound. To calculate the simplest formula we proceed as follows: 1. Divide the moles of each element by the smaller number of moles (0.040) and round to the nearest whole number. 0.080 mole Cl 0.040 mole Zn = 2 moles of Cl =1 mole of Zn 0.040 0.040 121
122 Laboratory Manual for General, Organic, and Biological Chemistry 2. Use the whole numbers as subscripts to write the formula of the compound. ZnC1 2 (The subscript 1 for Zn is understood.) B. Formula of a Hydrate A hydrate is an ionic compound that is combined with a specific number of water molecules. The number of water molecules is fixed for each hydrate, but differs from one hydrate to another. The number of water molecules is written after the ionic formula and separated by a large, raised dot. CaSO4 2H20 CuSO4 5H20 Na2CO3 10H20 The water molecules in the hydrate can be removed by heating. When all the water is removed, the remaining ionic compound is called an anhydrate. For example, when one mole of copper(ii) sulfate pentahydrate is heated, five moles of H 2O are removed. The water removed is also called the water of hydration. CuSO 4 5H 20 CuSO4 + 5H 20(g) hydrate anhydrate water of hydration The amount of water in a hydrate is experimentally determined by measuring the mass of the hydrate before heating and the mass of the anhydrate after heating. The difference in mass is due to the water of hydration that is lost. The percent water is calculated by dividing the grams of water by the mass of the hydrate and multiplying by 100%. For example, if 2.00 g of CuSO 4 5H20 is heated and the mass of the anhydrate CuSO 4 is 1.28 g, we would calculate the grams of H 2O as the difference. 2.00 g of hydrate 1.28 g of anhydrate = 0.72 g of H 2O in hydrate Then the percent water in the hydrate is calculated as 0.72 g H2O x100% = 36% H 2O in hydrate 2.00 g CuSO4
EXPERIMENTAL PROCEDURES A. Finding the Simplest Formula Moles and Chemical Formulas 123 GOGGLES REQUIRED! Materials: Crucible, crucible cover, crucible tongs, clay triangle, iron ring and stand, Bunsen burner, magnesium ribbon, steel wool, eyedropper, small 100- or 150-mL beaker, heat-resistant pad In this experiment, you will heat magnesium so that it reacts with the oxygen (0 2 ) in the air and forms an oxide. The difference between the mass of the oxide compound and the initial mass of the magnesium is the mass of oxygen that combined with magnesium. When the moles of the magnesium and the oxygen are calculated, the simplest formula can be determined. 1. Obtain a clean, dry crucible and its cover. Set the crucible, and cover, slightly offset, on a clay triangle and place on an iron ring attached to a ring stand (see Figure 11.1). Heat the crucible and cover for about one minute. Cool until they are at room temperature. Using crucible tongs, carry the crucible and cover to the balance. Do not place hot objects on a balance pan. Weigh the crucible and cover and record the mass. FIGURE 11.1 A crucible and cover, slightly offset, are heated on a clay triangle. 2. Obtain a piece of magnesium ribbon that has a mass of 0.15-0.30 g. If there is tarnish on the ribbon, remove it by polishing the ribbon with steel wool. Describe the appearance of the magnesium ribbon after polishing. 3. Twist the ribbon into a coil and place it at the bottom of the crucible. Weigh the crucible, cover, and magnesium ribbon and record the mass. Heating the Magnesium Ribbon Do this part of the experiment in a fume hood: Place the crucible with the magnesium ribbon on the clay triangle. Keep the cover and a pair of tongs nearby. Begin to heat the crucible making sure that the tip of the inner blue flame touches the bottom of the crucible. The bottom of the crucible will become red hot. Watch for smoke or fumes, which indicate that the magnesium and oxygen are reacting. As soon as the magnesium bursts into flame, use the tongs to place the cover on the crucible. The cover should be slightly offset to allow oxygen to react with the magnesium. Caution: Avoid looking directly at the bright flame of the burning magnesium. When the magnesium no longer produces smoke or a flame, remove the cover and set it on a heat-resistant surface.
124 Laboratory Manual for General, Organic, and Biological Chemistry Continue to heat the crucible strongly for another five minutes. Then turn off the burner and allow the crucible and its contents to cool to room temperature. During heating, some magnesium reacts with nitrogen in the air to form magnesium nitride. 3Mg(s)+ N2 (g) Mg3N2 (s) To remove this nitride product, carefully add 15-20 drops of water to the cooled contents. Mg3N2 (S)+ 3H20(1) A 3Mg0(s) + 2NH 3 (g) Caution: Avoid breathing fumes from the crucible because ammonia may be released. 4. Cover the crucible and heat gently for five minutes to drive off any excess water. Then heat strongly for five minutes. Allow the crucible to cool completely. Reweigh the crucible, cover, and oxide contents and record the mass. Remove the solid in the crucible and dispose of it as directed by your instructor. Calculations 5. Determine the mass of the magnesium (3 1). 6. Calculate the mass of the magnesium compound (4 1). 7. Calculate the mass of oxygen that combined with the magnesium (6 5). 8. Determine the number of moles of magnesium by dividing the mass of magnesium (5) by its molar mass. 1 mole Mg moles of Mg = x 24.31 y,-mg 9. Determine the number of moles of oxygen by dividing the mass of the oxygen (7) by its molar mass. x 1 mole 0 moles of 0 = 16.00 *CC 10. Divide the moles of Mg (8) and the moles of 0 (9) by the smaller number of moles. Round each of the results to the nearest whole number. 11. Using the whole number values obtained in 10 as subscripts, write the simplest formula of the magnesium compound. B. Formula of a Hydrate Materials: Crucible, clay triangle, crucible tongs, hydrate of MgSO 4, iron ring and stand, Bunsen burner, heat-resistant pad, laboratory balance 1. Obtain a clean, dry crucible and its cover. Set the crucible and cover, slightly offset, on a clay triangle and place on an iron ring attached to a ring stand (see Figure 11.1). Heat the crucible and cover for about one minute. Cool until they are at room temperature. Using crucible tongs, carry the crucible and cover to the balance. Do not place hot objects on a balance pan. Weigh the crucible and cover and record the mass. 2. Fill the crucible about 1/3 full with the hydrate of MgSO 4. Weigh the crucible with the MgSO 4 hydrate and record the mass. 3. Set the crucible and hydrate on a clay triangle that is set on an iron ring (see Figure 11.1). The cover should be slightly offset so that water vapor can escape. Heat gently for five minutes; increase the intensity of the flame and heat strongly for another 10 minutes. The bottom of the crucible should become a dull red color. Turn off the burner. Allow the crucible to cool to room temperature.
Moles and Chemical Formulas 125 Caution: Allow heated items to cool to room temperature. Do not place a hot object on the balance pan. Weigh the crucible and its contents. Record the mass. 3a. Repeat step 3. If more water has been lost, use this final mass in the calculations. Remove the solid in the crucible and dispose of it as directed by your instructor. Calculations 4. Calculate the mass of the hydrate (2 1) 5. Calculate the mass of the anhydrate after heating (3 1) 6. Calculate the mass of water lost from the hydrate sample (4 5) 7. Calculate the percent H 2O in the hydrate by dividing the mass of H 2O lost (6) by the mass of the hydrate (4) and multiply by 100%. Percent water in hydrate = g waterx 100% g hydrate 8. Calculate the moles of H 2O in the hydrate by dividing the mass of the H 2O (6) by its molar mass. 1 mole water Moles of water = pwatei x 18.02 gwatei-- 9. Calculate the moles of anhydrate by dividing the mass of the anhydrate by its molar mass. Moles of anhydrate = ate x 1 mole anhydrate 120.4 j...anhydrile 10. To determine the ratio of moles of water to 1 mole of anhydrate, divide the moles of water (8) by the moles of anhydrate (9). Round off the value for moles of H 2O to the nearest whole number. moles of water moles of H 2O moles of MgSO 4 1 mole of MgSO 4 11. Complete the formula of the hydrate by writing in the number of moles of water for each mole of anhydrate.
Date Section Instructor Name Team 1. What is meant by the simplest formula of a compound? Pre-Lab Study Questions 11 2. How does a hydrate differ from an anhydrate? 3. What happens when a hydrate is heated? 4. A hydrate of CoC1 2 with a mass of 6.00 g is heated strongly. After cooling, the mass of the anhydrate is 3.27 g. a. How many grams of H 2O were lost from the hydrate? b. What is the % water in the hydrate? c. What is the formula of the CoC1 2 hydrate? d. Write the equation for the dehydration of the CoC1 2 hydrate. 127
Date Section Instructor Name Team REPORT SHEET LAB Moles and Chemical Formulas 11 A. Finding the Simplest Formula 1. Mass of empty crucible + cover 2. Initial appearance of the magnesium 3. Mass of crucible + cover + magnesium 4. Mass of crucible + cover + oxide product Calculations 5. Mass of magnesium 6. Mass of magnesium compound 7. Mass of oxygen in the product 8. Moles of Mg (Show calculations.) g g g g g g mole 9. Moles of 0 mole (Show calculations.) 10. Which number of moles (Mg or 0) is smaller moles of Mg = moles of Mg (rounded to a whole number) moles 0 = moles of 0 (rounded to a whole number) Formula: Mt 0 subscripts 129
130 Laboratory Manual for General, Organic, and Biological Chemistry Questions and Problems Q1 Using the rules for writing the formulas of ionic compounds, write the ions and the correct formula for magnesium oxide. Q2 Write a balanced equation for the reaction of the magnesium and the oxygen (0 2 ), including their physical states. Q3 Calculate the simplest formula for each of the following compounds: a. 0.200 mole of Al and 0.600 mole of Cl b. 0.080 mole of Ba, 0.080 mole of S, 0.320 mole of 0 Q4 When 2.50 g of copper reacts with oxygen, the copper oxide compound has a mass of 2.81 g. What is the simplest formula of the copper oxide? B. Formula of a Hydrate 1. Mass of crucible 2. Mass of crucible and hydrate 3. Mass of crucible and anhydrate 3a. Mass of crucible and anhydrate (second heating) Calculations 4. Mass of hydrate 5. Mass of anhydrate 6. Mass of water lost
Moles and Chemical Formulas 131 7. Percent water % (Show calculations.) 8. Moles of water moles (Show calculations.) 9. Moles of salt (anhydrate) moles (Show calculations.) 10. Ratio of moles of water to moles of hydrate (Show calculations.) 11. Formula of hydrate MgSO4 H2O Questions and Problems Q5 Using the formula you obtained in B11, write a balanced equation for the dehydration of the MgSO 4 hydrate used in the experiment. (