Experiment 9 Empirical Formula of Zinc Iodide Objectives Upon completion of this experiment, students should have learned: 1. The law of conservation of mass. 2. How to calculate an empirical formula. 3. The concept of limiting reagents. Introduction Synthesis and the determination of empirical formulas are two extremely important parts of chemistry. In this experiment, you will synthesize zinc iodide and determine its empirical formula. The molecular formula gives the actual number of atoms of each element in the molecule and the empirical formula gives the lowest whole number ratio of the atoms in the molecule.
For instance, the molecular formula of glucose is C6H12O6 and its empirical formula is CH2O. For some molecules such as water, the empirical and molecular formulas are the same. One of the common techniques used to help determine the identity of a substance involves combustion of very small amounts of the substance. By measuring the amounts of water, carbon dioxide, and other gases produced, it is possible to determine the percentage by mass of each element in the compound. From the percentages, it is possible to calculate the empirical formula of the compound.
Suppose, for example, that the combustion experiment for a compound that contains carbon, hydrogen, and oxygen gives the compositions as 40. 00% carbon, 6. 71% hydrogen, and 53. 29% oxygen by mass. The goal is to convert mass percentages to mole or atom ratios. Whenever dealing with percentages in a calculation of this type, one of the easiest ways to proceed is to assume that you have 100 g of the substance. Thus 100 g of the substance contains 40. 00 g of carbon, 6. 71 g of hydrogen and 53. 29 g of oxygen. As is common for many chemical calculations, the calculation of the number of moles is one of the first steps.
Thinking ahead, once the number of moles is known, the ratio of moles is easily calculated. 40. 00 g C x 1 mol C ————- = 3. 33 moles C 12. 011 g C 6. 71 g H x 1 mol H ————- = 6. 66 moles H 1. 008 g H 53. 29 g O x 1mol O ————- = 3. 33 moles O 15. 999 g O Dividing by the lowest number of moles gives a C:H:O ratio of 1:2:1 so the empirical formula is CH2O. This means the compound could be glucose but it could also be many other compounds including other sugars or formaldehyde.
The determination of the molecular formula from the empirical formula requires an additional measurement of the molecular mass. If a molecular mass measurement for the above compound results in 150 ± 2 g/mole, the molecular formula would be calculated as C5H10O5. This means the compound could be a five carbon sugar such as ribose. Molecular formula = molecular mass ————————– x subscripts of empirical formula = empirical form. mass 150 x CH2O = C5H10O5. — 30 The empirical and molecular formulas can go a long way in helping to identify the compound.
Almost all ionic compounds are solids at room temperature. The crystal structure of the ionic solid consists of a repeating array, and it is not straightforward to define the smallest unit (a molecule) of the substance that still would have the properties of the substance. Thus it is better to refer to the formula mass of an ionic substance rather than molecular mass. For the same reasons, the formula calculations for ionic compounds should be reduced to the lowest whole number ratio or the empirical formula. Except for a few unusual cases, the term molecular formula is not applicable for ionic compounds.
For metals with only one common oxidation state, it is usually possible to predict the empirical formula that will result from the reaction with a non-metal with a high degree of confidence. For instance, sodium chloride and sodium oxide should be and are NaCl and Na2O respectively. Aluminum chloride and aluminum oxide should be and are AlCl3 and Al2O3. When there are multiple oxidation states of a metal, more than one compound is possible. For instance, copper and oxygen form Cu2O and CuO and more information are needed to distinguish between the two.
When naming a copper oxide, it is very important that the name distinguish between the two possibilities [copper(I) oxide or cuprous oxide and copper(II) oxide or cupric oxide]. Care must be exercised as chemistry often contains wonderful surprises. Suppose the percent by mass experiment for an iron oxide results in 72. 36% Fe by mass and 27. 64% O by mass. Following the procedure above, results in a molar ratio of 1. 296:1. 727. Division by the lowest number, results in a mole ratio of 1:1. 33. For cases like this, the next step is to multiply by the lowest whole number that converts each of the subscripts to a whole number.
Thus the formula is Fe3O4. As iron has two oxidation states (+2, +3), the expected formula is either FeO of Fe2O3. Anytime a result that is inconsistent with expectations is obtained, the first and immediate step should be to recheck all the data and calculations. In this case, the results are real and Fe3O4 is a known compound often called magnetite. Understanding the oxidation state of iron in this compound will be left to you or your instructor but it is results like this that should not be ignored and often lead to important and exciting discoveries. Today, iodine will be reacted with an excess of zinc.
The masses of the starting materials, products and leftover starting material will enable you to calculate the formula of the zinc iodide produced. Since zinc in compounds has only the +2 oxidation state, the product is certainly expected to be ZnI2. However, as noticed above for iron, expectations do not always agree with results so it is always necessary when performing a synthesis to verify the formula of the product. Procedure Be sure to record all masses to the limit of the balance capability. Weigh a 50 mL beaker to the nearest milligram (or if milligram balances are not available, to the nearest 0. 1 g) and record the mass. Tare bar procedure. Push the tare bar, add about 1. 2 g of 20 mesh granular zinc to the beaker and record the mass of the zinc. Again push the tare bar, carefully add about 1. 2 g of iodine [Caution: iodine is toxic and should not be touched] to the beaker containing the zinc and record the mass of the iodine. No tare bar procedure. Weigh out about 1. 2 g of 20 mesh granular zinc on weighing paper, add it to the pre-weighed beaker and record the mass of the beaker and its contents. Now carefully weigh out about 1. g of iodine [Caution: iodine is toxic and should not be touched] on a new weighing paper, add it to the beaker containing the zinc and record the mass of the beaker and its contents. Add 3 mL of 0. 2 M acetic acid (acetic acid prevents the precipitation of zinc hydroxide) to the solid mixture in the beaker, swirl and observe changes of temperature and colour. Continue to swirl until no further changes are evident. The beaker should have cooled back down to room temperature and the iodine and triiodide (iodine and iodide in aqueous solution combine to form I3-) colour should have disappeared.
This process should take about 20 minutes. Weigh an evaporating dish and record the mass. Carefully decant (pour off) the aqueous solution into the evaporating dish, being sure that all leftover zinc remains in the original beaker. Add about 1 mL of 0. 2 M acetic acid to the beaker (and contents), swirl and decant again into evaporating dish. This “washes” the remaining zinc free of residual zinc iodide and transfers it to the evaporating dish. Repeat the washing process two more times. Wash the zinc three more times with 0. 2 M acetic acid and discard these additional washings but save the beaker and the remaining zinc.
Mount the evaporating dish in a 250 mL beaker containing about 100 mL of water on a wire gauze above a Bunsen burner. Heat the beaker until the contents of the evaporating dish appear to be dry. Achieving apparent dryness with steam heating can consume considerable time. Time can be saved if you moved to the next step before dryness is completely attained. Carefully remove the beaker with beaker tongs and place the evaporating dish directly on the wire gauze. Very gently flame the dish until crackling stops and the solid in the dish turns slightly off-white.
Excessive heating results in a dark yellow colour indicating that the product is decomposing to iodine. This needs to be avoided. Allow the dish to cool and weigh it. Repeat the direct heating, cooling and weighing cycle until constant mass is achieved (within at least 0. 02 g). Now flame the beaker containing the zinc until it is dry. Allow it to cool and weight it. Repeat the heating, cooling and weighing cycle until constant mass is achieved (within at least 0. 02 g). Optional enhancement. Dissolve about 0. 1 g of the white solid that remains from the aqueous solution in 2 mL of water on a watch glass. Obtain a battery clip (e. . , Radio Shack 270-325) that has had two copper leads attached to the ends of the clip wires. Attach the clip to a 9-volt battery and insert the two copper leads into the solution. Record your observations. Results and Discussion Empirical Formula of Zinc Iodide 1. Write the expected balanced equation for the reaction between zinc and iodine ____________________________________________________________________ 2. Mass of beaker ____________ 3. Mass of evaporating dish ____________ . Mass of iodine ____________ 5. Mass of zinc ___________ 6. Mass of zinc + iodine ____________ 7. Observations after addition of water containing acetic acid 8. Mass of beaker + residual zinc __________ __________ ____________ trial 1 trial 2 trial 3 if necessary . Mass of residual zinc ____________ 10. Mass of zinc consumed in reaction ____________ 11. Mass of evap. dish + zinc iodide __________ __________ ____________ trial 1 trial 2 trial 3 if necessary 12. Mass of zinc iodide ___________ 13.
Was mass conserved within experimental error in this reaction? Explain your answer. ____________________________________________________________________ 14. Moles of zinc consumed in reaction ___________ 15. Moles of iodine consumed in reaction ___________ 16. Mole ratio of zinc to iodine ____________ 17. Formula of zinc iodide ____________ 8. Does your experimental formula agree with the expected formula? Explain your answer. 19. In terms of the chemistry and energy of the reaction, explain your observation in #7 above. 20. Explain why some zinc was left over at the end of the reaction and why you think the amounts were selected to have zinc left over. 21. (optional) What did you observe in the electrolysis reaction? 22. (optional) Were your observations on the electrolysis consistent with your expectations? Explain your answer. 23. Suggest any ways you can think of to improve any part(s) of this experiment.