Experiment 3: Empirical Formula PDF

Experiment -Three Empirical Formula Objectives:  To determine the empirical formula a compounds by combination reaction.  Determine the percent yield of the reaction. Introduction: The empirical formula of a compound is the simpl...

Experiment -Three Empirical Formula Objectives:  To determine the empirical formula a compounds by combination reaction.  Determine the percent yield of the reaction. Introduction: The empirical formula of a compound is the simplest whole-number ratio of moles of elements in the compounds. The experimental determination of the empirical formula of a compound from its elements requires three steps:  Determine the mass of each element in the sample.  Calculate the number of moles of each element in the sample.  Express the moles of each element in the ratio of whole numbers. For example: an analysis of a sample of table salt shows that 2.75 g of Na and 4.25 g Cl are present. The numbers of moles of these elements are: mol of Na = 2.75 g/23 g/mol = 0.12 mol mol Cl = 4.25g /(35.5 g/mol) = 0.12 mol The mol ratio of Na to Cl is 0.12 to 0.12. As the empirical formula must be expressed in a ratio of small whole numbers, the whole-number ratio is 1:1 and the empirical formula is NaCl. The empirical formula also provides a mass ratio of the elements in the compound. The formula NaCl states that 23.0 g (1 mol) of Na combine with 35.5 g (1 mol) of Cl to form 56.5 g (1 mol) of NaCl. The empirical formula of a compound can be determines by:  Combination reaction: two elements combine to form a compound. A known mass (limiting reactant) of one reactant and the mass of the product are determined. Example: 0.54 g of Al ignited in air to produce a compound of Al and O that weighed 1.02g. What is the empirical formula of the compound. 36  Mol of Al = 0.54 g/(27g/mol) = 0.02 mol Mass O = 1.02 g – 0.54 g = 0.48 g mol O = 0.48 g /(16g/mol) = 0.03 mol mol Al : mol O 0.02 : 0.03 empirical formula is Al2O3  Decomposition reaction: a compound decomposes into two or more elements or simpler compounds. The initial mass of the compound used for the analysis and the final mass of at least one of the products are determined. An example would be the decomposition of HgO to Hg metal and O2 gas, the initial mass of HgO and the final mass of Hg metal are determined. The difference between the measured masses is the mass of O in HgO. The moles of Hg and O in the original compound are the determined to provide a whole-number mole ratio of Hg:O. In this experiment a combination reaction of magnesium and oxygen is used to determine the empirical formula of magnesium oxide. When magnesium is heated to a high temperature, it reacts with oxygen in the air to form magnesium oxide. x Mg(s) + (y/2) O2(g) → MgxOy The masses before and after the oxidation is measured, the resulting masses are used to calculate the experimental empirical formula of magnesium oxide which is then compared to the theoretical empirical formula. Although, there is a higher percentage of N2 gas in the atmosphere than O2, O2 is more reactive and magnesium oxide is formed in greater amount than the nitride. 3 Mg(s) + N2(g) → Mg3N2(s) In this experiment, we are interested in making only the oxide and not the nitride. To accomplish this, we add water to the mixture of magnesium oxide and magnesium nitride. Upon heating, the nitride reacts with the water to form magnesium hydroxide and ammonia gas. The ammonia is driven off in the heating. The magnesium hydroxide is heated to high temperature and decomposes to magnesium oxide and water. Thus, at the end of the experiment, all of the magnesium has been converted to the desired product, magnesium oxide. Mg3N2(s) + 6 H2O(l) → 3 Mg(OH)2(s) + 2 NH3(g) Mg(OH)2(s) → MgO(s) + H2O(l) 37 Based on the masses of magnesium and oxygen that combine, you can calculate the empirical formula of magnesium oxide. you will weigh the magnesium before it combines with the oxygen, and we will also weigh the product of the reaction, magnesium oxide. The final weighing is necessary because you need to subtract the original weight of magnesium from this weight of product. You "weigh" the oxygen in this indirect fashion because it is easier to do than to try to weigh the oxygen gas before it combines with the magnesium. Experimental procedure: A combination reaction of magnesium and oxygen 1. Prepare a Clean Crucible: a) Obtain a crucible and lid, clean the crucible thoroughly with tap water and rinse with distilled water. The crucible may have a dark stain that can not be removed, this is O.K. b) Support the crucible on a clay triangle as shown in figure (3-1). c) Fire the empty crucible with intense flame for 5 minutes to remove water, oil, or other contaminants and to make sure there are no cracks. The bottom of the crucible should be at the tip of the inner core of the flame (hottest part) and should glow red-orange. d) Allow the crucible to cool to room temperature. Be sure that the crucible is cool before weighing, if it is too hot, the mass recorded will be incorrect e) Determine the mass of the fired, cool crucible to the nearest 0.01g and record the mass on the data sheet. Note: - Use only crucible tongs to handle the crucible and the lid for the remainder of the experiment as shown in figure 3-2; do not use your fingers as the oil and dirt from your fingers can contaminate it and add to its mass and /or you could be severely burned. - Do not place a hot crucible on the lab bench, the temperature difference may cause it to break. The hot crucible may cause the bench to burn. 38 2. Prepare the Sample: a) Polish 0.2-0.3 g of magnesium ribbon with steel wool to remove any oxide coating it until it become shiny b) Curl the ribbon to lay in the crucible. Note: do not coil the ribbon so tightly that it will be difficult for Mg to react completely. c) Measure the combine mass of the crucible and magnesium ribbon to nearest 0.01g and record the mass on the data sheet. 3. Heat the Sample in Air: a) Place the crucible containing the Mg-ribbon and lid on the clay triangle. b) Heat slowly, occasionally lifting the lid slightly to allow air to reach the Mg- ribbon (figure 2-1). If too much air comes in contact with the Mg-ribbon, rapid oxidation of Mg occurs and it burns brightly. (Caution: you don not want this to happen. If it does, it will cause some of the metal to vaporize. Do not watch the burn; it may cause temporary blindness). d) Immediately return the lid to the crucible to slow down the reaction by limiting the supply of O2 that react with Mg. 4. Heat for Complete Reaction: a) Continue heating the crucible and lifting the lid until no visible change is apparent in the magnesium ash at the bottom of the crucible (sample no longer glow brightly when you lift the lid). b) Remove the lid; continue heating the open crucible and ash for 2-3 minutes. c) Remove the heat and allow the crucible to cool to room temperature on the wire gauze. d) Add a few drops of water (about 10 drops) to decompose any magnesium nitride form during combustion. (Caution: cool water added to the hot crucible may cause the crucible to break). e) Dry the ash with a low flame for 2 minutes and then strongly for about 1 minute. Allow the crucible to cool to room temperature. f) Measure the mass of the open crucible on the same balance as was used earlier and records the mass on your data sheet. 39 5. Test for Complete Reaction: a) Reheat the sample for 1 minute, but do not intensify the flame. b) Again measure the mass of the crucible and ash. If this second mass measurement is greater than ± 1% from that recorded in step 4, repeat the heating process until you obtain a "constant mass" by successive measurements. 6. Calculations: a) Determine the mole ratio of magnesium to oxygen in the compound. b) Calculate the theoretical maximum mass of magnesium oxide that would formed from the mass of magnesium. c) Calculate the percent yield of magnesium oxide percent yield = actual yield MgO x 100% theoretical yield MgO d) Calculate the percent error for your experiment percent error = │theoretical mass MgO – actual mass MgO│ x 100% theoretical mass MgO 40 Figure 3-1: Heating a crucible Figure 3-2: Using crucible tongs to hold a crucible. 41  EXPERIMENT-Three Pre-laboratory Assignment Empirical Formula Name: ………………………….. Lab Section: ……………….Date: ………………….. 1. When 3.20 g of sulfur is burned, 6.37 g of the oxide is obtained. a) Write the balanced chemical equation for the reaction b) Determine the empirical formula of the compound c) Determine the percent yield of the reaction 2. Explain how the mass of oxygen that combine with magnesium is determined in today's experiment. 3. In today’s experiment, when magnesium burns in air, in addition to the oxide being formed, what other product will be formed? How do you convert this other compound to oxide? 42  EXPERIMENT-Three Report Sheet Empirical Formula Name: …………………………..Lab Section: ……………….Date: ………………….. DATA SHEET and CALCULATIONS: Trial (1) Trial (2) Mass of empty crucible (m1) g g Mass of empty crucible and magnesium before heating (m2) g g Mass of empty crucible and magnesium after heating (first g g heating) (m3) Mass of empty crucible and magnesium after heating (second g g heating) (m4) Average mass of empty crucible and magnesium after heating g g (m5) [(m3 + m4)/2] Mass of magnesium (m2 – m1) g g Mass of magnesium oxide (actual yield) (m5 – m1) g g Mass of oxygen in the oxide (m5 – m2) g g Moles of magnesium (n1) mol mol Moles of oxygen (n2) mol mol Formula of magnesium oxide (Mgn1On2) Empirical Formula of MgO Theoretical yield of MgO (calculated from MgO) g g [Mass MgO=moles of Mg x(1mol MgO/1mol Mg)xmolar mass MgO] %yield of the reaction [(actual yield / theoretical yield) x 100%] % % Percent error of your experiment % % 43  EXPERIMENT-Three Post-laboratory Report Empirical Formula Name: ………………………….. Lab Section: ………………. Date: ………………….. 1. What evidence do you have that a chemical reaction took place 2. How would your determination for the empirical formula be affected by the following sources of error? (a larger or smaller ratio of magnesium to oxygen? Explain your answer completely. a. The magnesium was not allowed to react in sufficient air. b. During the reaction, the magnesium metal burst into white flame and white smoke escaped from the crucible. 3. A sample of a compound containing C, O, and silver (Ag) weighed 1.372 g. On analysis it was found to contain 0.288g O and 0.974 g Ag. The molar mass of the compound is 308.8 g/mol. Determine a)the molecular formulas of the compound. b) the percent by mass of each element 44