4. Determining the Chemical Formula of an Ionic Compound

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1 4. Determining the Chemical Formula of an Ionic Compound What you will accomplish in this experiment You ve learned that compounds are a chemical combination of elements, meaning that they re created when two or more elements chemically react with one another. The driving force for that reaction is that the elements are trying to achieve the stability of a noble gas electron configuration (for most elements, this means an s 2 p 6 octet of electrons. When the reacting elements are a metal and a nonmetal (elements with an electronegativity difference greater than 2.0), the metal atom achieves its octet when it loses one or more electrons (to become a positively-charged ion, a cation ), and the nonmetal atom gains those electrons (to become a negatively-charged ion, an anion ). The chemical formula for the ionic compound that s formed describes the fixed proportion in which vast numbers of the oppositely-charged ions combine in order to balance out the positive and negative charges. This idea is illustrated below for the chemical combination of the elements Sodium and Chlorine to make the ionic compound: Sodium Chloride, NaCl. The ionic compound you ll work with in this experiment is the chemical combination of the metal element Copper (Cu) and the same nonmetal element as above, Chlorine (Cl). When these elements chemically combine, the fixed proportion of ions in the compound can be expressed by the chemical formula, CuCl 2, and by the compound s name, Copper (II) Chloride. Your job will be to chemically separate these two elements, and then to determine the respective masses of Cu 2+ and Cl - ions that were in your original sample of Copper (II) Chloride. By converting these masses of each element into moles (using the molar masses of Copper and Chlorine), and then determining the mole ratio of the two elements in the compound, you should be able to prove experimentally that the chemical formula for Copper (II) Chloride is, in fact, CuCl 2. C. Graham Brittain Page 1 of 9 9/28/2010

2 Concepts you need to know to be prepared A compound s chemical formula is made up of the symbols (from the Periodic Table) that represent each of the elements in that compound. And the numerical subscripts following these symbols indicate the fixed proportion in which the elements combine when that compound is formed by a chemical reaction. For an ionic compound (such as sodium chloride, NaCl), you should NEVER think of that compound as just a PAIR of oppositely-charged ions, as the chemical formula seems to imply. If you look closely at the illustration of a sodium chloride (NaCl) salt crystal to the right, you ll see that it is definitely NOT just a pair of Na + and Cl - ions, but an extensive, orderly, threedimensional network of charged particles. Each Na + ion is attracted to the Cl - ions immediately above, below, to the right, to the left, in front, and behind it. And each Cl - ion is attracted to the Na + ions surrounding it. So it s important to realize that the chemical formula for ANY ionic compound does NOT describe a small grouping of ions, it describes the fixed proportion in which vast numbers of the oppositely-charged ions combine in order to appropriately balance out the positive and negative charges. So one way to think about the chemical formula for Copper (II) Chloride, CuCl 2, the ionic compound in this week s experiment, is that it expresses the ION ratio of the elements in the compound: The compound CuCl 2 contains 2 Cl - IONS for every 1 Cu 2+ ION. But a bigger picture way to think about that chemical formula is that it expresses the MOLE ratio of the elements in the compound: The compound CuCl 2 contains 2 MOLES of Cl - ions for every 1 MOLE of Cu 2+ ions. One mole of any pure substance (element or compound) is really just an Avogadro s number of particles of that substance (where particles = atoms, ions, or molecules). For example: Or in the case of Copper (II) ions: 1 mole of Copper metal = 6.02 x Cu atoms 1 mole of Copper (II) ions = 6.02 x Cu 2+ ions An Avogadro s number of atoms for a particular element, one MOLE of the element, will ALWAYS have a mass in GRAMS equal to the Atomic Weight listed for that element on the Periodic Table: Or for Copper (II) ions: 1 mole of Copper (Cu) = 6.02 x atoms of copper = grams of Copper metal 1 mole of Copper (II) ions = 6.02 x Cu 2+ ions = grams of Copper (II) ions NOTE: The mass in grams of 1 mole of any pure substance is commonly referred to as the Molar Mass. A similar unit relationship is easily determined for any other element on the Periodic Table. The first part of this unit relationship is always the same: 1 mole of any element = 6.02 x atoms of that element C. Graham Brittain Page 2 of 9 9/28/2010

3 The unit relationship is then completed by taking the Atomic Weight of the desired element from the Periodic Table. For example: Or for Chloride ions: 1 mole of Chlorine (Cl) = 6.02 x atoms of chlorine = grams of chlorine 1 mole of Chloride ions = 6.02 x Cl - ions = grams of chloride ions These unit relationships can be used to convert the mass in grams of any substance to the number of moles represented in that mass. And the unit conversion from grams of an element to moles of an element is essential to this week s experiment. You ll begin by obtaining a sample of copper (II) chloride and recording its mass to the maximum number of significant figures provided by the balance. You ll then dissolve this compound in water. In dissolving, ionic compounds separate into their component ions. So in water, copper (II) chloride dissociates into copper (II) ions (Cu 2+ ) and chloride ions (Cl - ). The hydrated Cu 2+ ions have a brilliant blue color, while the Cl - ions are colorless. Into this solution, you ll place a coiled piece of aluminum (Al) wire. This is aluminum metal: an element, not a compound. The aluminum wire will spontaneous react with the Cu 2+ ions in the solution, and the copper will be converted into its elemental form: copper metal. Think about what this means: The Cu 2+ ions will be gaining electrons from the Al atoms, and thus converted into atoms of Cu metal. The atoms of Al metal are losing electrons to the Cu 2+ ions, and are converted into colorless, water-soluble Al +3 ions. And you ll be able to observe this chemical transformation: The brilliant blue color of the aqueous Cu 2+ ions will eventually fade away, and brown copper metal will collect at the bottom of your reaction beaker. The aluminum wire will become smaller as the Al atoms are converted into aqueous Al +3 ions. This chemical reaction is slightly exothermic, meaning that heat is released as the reaction progresses. You can observe this heat release in two ways: qualitatively by simply noting the warmth of the reaction container by touching it carefully with your fingertips, and quantitatively by measuring the actual temperature change of the aqueous mixture. The assumption that you ll make in this experiment is that every Cu 2+ ion in solution will react with the aluminum wire and be converted into an atom of copper metal. After the reaction goes to completion, you are to filter, dry, and weigh this copper metal. At this point, realize that you ve chemically separated the two elements in the compound (copper and chlorine), and you ve determined the mass of copper that was in your original sample. By subtracting this mass of copper from the mass of your original copper (II) chloride sample, you ll know the mass of chlorine that was in that sample. Now remember your mole unit relationships: You can use the molar masses of copper and chlorine from the Periodic Table to convert the mass in grams of each element to the number of moles of that element. So your final steps are to convert your experimental masses of copper and chlorine into moles, and then to determine the mole ratio of chlorine to copper in the compound. As discussed, the chemical formula of CuCl 2 indicates that two moles of chloride ions combine with one mole of copper (II) ions when the compound copper (II) chloride is formed: a mole ratio of two-to-one. Ideally, your experimental result should be consistent with this fixed proportion. C. Graham Brittain Page 3 of 9 9/28/2010

4 If it s not, you ll need to think back through the various steps of the procedure, identify the possible sources of experimental error, and consider how each source of error might have contributed to making your experimental result too high or too low. You can see that it s essential to do a good job of recording your data and observations, so that you ll be able to make a thorough assessment as to why your experimental result is higher or lower than the true value. Procedure that you will follow 1. Measure a known mass of copper (II) chloride into a 250-mL beaker (approximately 3 grams, but weighed to the maximum number of significant figures provided by the balance). 2. Add approximately 60 ml of distilled water to the beaker and stir until the copper (II) chloride has completely dissolved. The resulting solution should be the brilliant blue color that is characteristic of Cu 2+ ions in aqueous solution. 3. Your lab instructor will provide you with an aluminum wire and demonstrate how to shape it into a coil. Take a moment to measure the initial mass of the wire. You ll then immerse the coiled end of the wire in the copper solution. Maintain your grasp on the other end of the rod, so that you can use it to stir the reaction mixture. 4. Watch closely to observe the chemical reaction that occurs, and make detailed notes of your observations in your lab notebook. IMPORTANT: Be sure to review the Skills in Recording Observations provided in the Remedies document: When recording observations of a substance, describe everything that you see (or perhaps smell): Is the substance a solid, liquid, or gas? Is it a solution? What color is it? Is it opaque, translucent, or transparent? If you are instructed to waft the vapors toward your nose, does the substance have an odor? When recording observations of a reaction, first write down what you did (heat a solution, or add one chemical to another and mix thoroughly). Then write down what you saw, heard, smelled, or felt with your fingertips: Was there a color change? Did a gas evolve? Did a solid form? Was an odor emitted? Was there any sound? Was any heat evolved? (If yes, monitor the temperature change with your thermometer, and record it as an observation.) How long did it take for the reaction to occur? A Note on Common Sense and Experimental Technique: Because the chemical reaction between the Cu 2+ ions and the Al atoms is taking place on the surface of the aluminum wire, the copper metal that s produced by the reaction will accumulate on that surface. So you ll need to use your metal spatula or glass stirring rod to scrape or knock the copper metal off the wire as the reaction progresses. The goal is to keep exposing the aluminum surface to the solution so that the reaction can continue all the way to completion (until all of the copper ions in the solution have been converted to copper metal). 5. When the reaction appears to be complete, remove the aluminum wire, rinse and dry it thoroughly, determine its final mass, and then return it to your lab instructor. 6. Now you ll need to separate the copper metal from the reaction mixture by vacuum filtration. In preparation for this procedure, pre-weigh a piece of filter paper (that fits in the Buchner funnel) with your watchglass. Set up the vacuum filtration apparatus as illustrated in the LabCam video and described in the Remedies document. Be sure to follow the steps carefully especially securing the Buchner funnel to the filter flask with the rubber adaptor, and clamping the neck of the filter flask to a ring stand. Ask your lab instructor to inspect your vacuum filtration set-up before you begin filtration. 7. As shown in the video, you ll need to place your pre-weighed filter paper in the Buchner funnel, pre-moisten it with water, and make certain that the wet paper forms a seal against the funnel when the vacuum is turned on. C. Graham Brittain Page 4 of 9 9/28/2010

5 Then pour the mixture into the center of the funnel. Remember that you can use small squirts of distilled water from your plastic wash bottle to rinse the copper from the beaker into the funnel. 8. As the final step to the vacuum filtration, rinse the copper metal in the funnel with several squirts of acetone from the plastic squeeze bottle provided in the hood. Acetone is a volatile organic solvent, so it will help the copper dry more quickly. You should also allow the vacuum to pull air through the copper for several minutes, to remove excess liquid and hasten drying. 9. Then carefully transfer the wet filter paper and copper metal from the funnel to the pre-weighed watch lass. Be sure to spread the copper across the surface of the watchglass to enable its drying. Mark the watchglass with your initials, and then place it in the heated oven to dry. When the drying and cooling is complete, weigh and record the mass of the watchglass (containing the dried filter paper and copper). The mass of the copper metal can be determined by difference. 10. Then by difference, determine the mass of chlorine in your original sample of copper (II) chloride. 11. Use Atomic Weights from the Periodic Table to convert your masses of copper and chlorine to moles, then determine your experimental mole ratio of chlorine to copper. IMPORTANT: You MUST dispose of all chemical waste as directed by your lab instructor. Do NOT put any chemical waste in the laboratory sinks or garbage cans. Use the solid and liquid waste containers in the hood as directed by your lab instructor. ALSO IMPORTANT: Please thoroughly clean your laboratory glassware before replacing it in your equipment drawer. Large Nalgene bottles of soap solution are provided near the laboratory sinks. You should notify your lab instructor if the stock of soap solution is running low. C. Graham Brittain Page 5 of 9 9/28/2010

6 Report Sheet 4: Chemical Formula of a Compound Student Lab Partner Date Lab Performed Section # Lab Instructor Date Report Received Lab Notebook: Data and Observations Experimental Data: Mass of empty 250-mL beaker Mass of 250-mL beaker with copper (II) chloride sample Mass of copper (II) chloride sample (by subtraction) Mass of aluminum wire before reaction Mass of aluminum wire after reaction Mass of watchglass and filter paper (pre-weighed) Mass of watchglass, filter paper, and dried copper metal Additional Data: Molar Mass of Copper Molar Mass of Chlorine Observations: Initial physical appearances of the copper (II) chloride solid, aqueous solution, and the aluminum wire. Chemical reaction of the aluminum wire with the copper solution, including the temperature change. Final physical appearance of the copper metal, aluminum wire, and filtered solution. C. Graham Brittain Page 6 of 9 9/28/2010

7 Formal Report: Results and Conclusions Change in mass of aluminum rod Mass of copper metal produced by the reaction Mass of chlorine in original sample of copper (II) chloride Number of moles of copper metal Number of moles of chlorine Experimentally determined mole ratio of chlorine to copper Chemical formula of copper (II) chloride as determined experimentally C. Graham Brittain Page 7 of 9 9/28/2010

8 Additional Questions: 1. One way to express the fixed proportion of elements in the compound copper (II) chloride is the chemical formula. Another way is the Mass Percent of each element in the compound. For example: The theoretical Mass Percent of copper in the compound (as predicted by the true chemical formula, CuCl 2 ), would be: Complete this calculation and determine the Theoretical Mass Percent of Copper in copper (II) chloride: 2. Now calculate the Mass Percent of Copper from your experimental data: 3. Compare your experimental mass percent of copper to the theoretical value by calculating the Percent Error. C. Graham Brittain Page 8 of 9 9/28/2010

9 4. Thoroughly explain why your experimental mass percent is different from the theoretical value. Think about what you observed: Was there any color remaining in the solution after your vacuum filtration? Was any copper metal lost in transfer or filtration? Was the copper metal still moist when it was weighed? Be specific as to how each possible source of error would have affected the outcome; that is, caused the mass percentage of copper to be higher or lower than the theoretical value. Finally, explain where you could improve your technique if you repeated the experiment. 5. Thought Question: How did the change in mass of the aluminum rod compare to the mass of copper produced by the chemical reaction? Larger? Smaller? The same? And is this result consistent with your expectation? Explain your reasoning. C. Graham Brittain Page 9 of 9 9/28/2010

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