Laboratory 3 Development of an Equation Objectives Apply laboratory procedures and make observations to investigate a chemical reaction. Based on these observations, identify the pattern of reactivity and communicate what occurred in the reaction by writing a chemical equation and a particle diagram consistent with the data. Introduction Suppose two substances are added together and they begin burning, with smoke (and a great demonstration!), and she may describe this phenomenon in different ways. One way is to record macroscopic observations. What took place? What was measured? What was seen, heard, smelled? Or perhaps the chemist thinks about her observations and infers (concludes based on evidence and reasoning) what happened to the particles. What did the atoms, ions, or molecules do as they ist cannot actually see the atoms, but based on the experimental evidence it might be possible to provide an explanation, a theory, that is consistent with the observations. Finally, the chemist may decide to write a chemical equation to symbolize what took place during the reaction. All of these are legitimate ways to describe what transpired and in this experiment, you will utilize each one. Chemical equations are usually presented as facts in textbooks, but how does a chemist know that the reaction really took place as written? Where is the evidence? What led to the inference? In this experiment, you will investigate a particular 13
Laboratory 3 reaction, make observations, infer what you think happened to the atoms and ions, and Discussion Chemical Equations and Particle Diagrams Chemical equations are used to describe a chemical reaction with symbols. For example, nitrogen, N 2, and oxygen, O 2, can combine to form nitric oxide, NO. The chemical equation for this is: N 2 (g) + O 2 (g) $ 2 NO(g) Reactants Product In this equation, the (+) symbol indicates that nitrogen reacts with oxygen and the arrow indicates that nitric oxide is formed. The chemical formulas on the left side of the equation are collectively known as the reactants and those on the right side as the products. In this case we have one kind of product, NO(g). The letter g in parentheses is included to indicate that these are in the gaseous phase. Other phase symbols include (s) for solid, (l) for liquid, and (aq) for aqueous or in water. The subscripts in the formula tell us how many of each atom there are in a molecule. In the formula above, there are two nitrogen atoms in the gas molecule N 2, two oxygen atoms in the gas molecule O 2, and one nitrogen atom with one oxygen atom in the gas molecule NO. The number in front of the chemical formula, or, indicate how many molecules are present. In a chemical reaction atoms are not created or destroyed, thus there should be an equal number of atoms of each element on the left and right sides of the equation. An equation is balanced when there are an equal number of atoms of each element on the left and right sides of the equation. In the equation above, we can see that one nitrogen gas molecule and one oxygen gas molecule combine to form two nitric oxide gas molecules. There are a total of two nitrogen atoms on the left side of the equation and two nitrogen atoms on the right side of the equation. The same is true for the oxygen atoms, so this chemical equation is balanced. This same reaction may also be described by indicating what happens to the particles: Before After N 2 (g) + O 2 (g) 2NO(g) 14
Development of an Equation Question 3.1: Nitrogen and oxygen can combine to form other molecules containing N and O. One is nitrous oxide, or N 2 O. N 2 O is commonly used in dentistry for its anesthetic and analgesic effects and is better known as laughing gas. Write a balanced chemical equation for N 2 reacting with O 2 to form laughing gas. In the boxes below, draw a particle diagram showing the reactants (before) and the products (after) the reaction. The drawing should be consistent with your balanced equation. Reactants (Before) Products (After) Common Patterns of Chemical Reactivity Different types of common chemical reactions are discussed in your textbook. These include: 1. Combination reactions, which describe the reaction of two or more reactants to form one product. 2. Decomposition reactions, in which one substance reacts to form two or more compounds or elements. This type of reaction typically occurs when compounds are heated. 3. Combustion reactions, which often occur quickly and evolve heat in the form of a bon dioxide and water. 4. in which the positive and negative ions appear to exchange partners. Question 3.2: Is the chemical reaction N 2 (g) + O 2 (g) $ 2 NO(g) an example of a combination, decomposition, combustion, or exchange reaction? Overview of Experiment In this experiment you will investigate the reaction that takes place when an aqueous lead nitrate solution is combined with an aqueous potassium iodide solution. Your objective is to determine the appropriate pattern of reactivity and propose a chemical equation and particle diagram that is supported by your laboratory observations. To get started you must make observations. In Part A you will combine a lead nitrate solution with a potassium iodide solution and record your observations. What atoms, ions, 15
Laboratory 3 or molecules do you think are in each solution before they are combined? After they are combined, is there evidence for a chemical reaction? Is there any indication that a gas is formed? What happens to the solutions of lead nitrate and potassium iodide when they are combined? Is there a chemical reaction? To answer these questions you must be able to identify the presence of different ions. In Part B you will perform qualitative tests for the different possible ions in this experiment. Procedures are given for these qualitative tests. Make careful observations when the known ions are tested, since you will be repeating these tests with your actual sample and comparing the results. In Part C you will apply your understanding of qualitative tests to determine which ions are in the solid product that forms when lead nitrate and potassium iodide solutions are mixed. Determining the identity of ions comprising your solid product is important. After Part C you probably have enough evidence to select the appropriate pattern of reactivity. However, you need more information to write a chemical equation. The qualitative analysis determined what ions are present, but it did not provide any quantitative information. A chemical equation includes quantitative information and communicates the amounts of the different substances. To get this information you must complete a quantitative test. In Part D you will perform a quantitative test to determine the exact ratio of ions in the precipitate that formed during your reaction. This will be accomplished by mixing solutions of lead nitrate and potassium iodide (the reactants) in varying ratios and making observations. There is an exact stoichiometric ratio in which these two substances will combine with each other such that neither one remains. However, for every other ratio the reaction will take place, but one of the reactants will be leftover. The reactant that remains after the reaction was in excess. The reactant that was completely consumed is the limiting reactant. To illustrate the idea of reaction stoichiometry, consider the burning of hydrogen gas in oxygen to form water vapor: 2 H 2 (g) + O 2 (g) $ 2 H 2 O(g) In the balanced equation, the H 2 :O 2 ratio is 2:1. When hydrogen gas is combined with oxygen in a 2:1 ratio, all of the H 2 will react with all of the O 2 to produce H 2 O(g) and there are no leftover reactants. 16
Development of an Equation What if different ratios of H 2 to O 2 are used? There will be leftover reactants: Ratio of H 2 :O 2 Before (Reactants) After (Products) 4:1 4H 2 (g) + O 2 (g) 2H 2 O(g) + 2H 2 (g) 3:1 3H 2 (g) + O 2 (g) 2H 2 O(g) + H 2 (g) 2:1 2H 2 (g) + O 2 (g) 2H 2 O(g) 1:1 2H 2 (g) + 2O 2 (g) 2H 2 O(g) + O 2 (g) 1:2 2H 2 (g) + 4O 2 (g) 2H 2 O(g) + 3O 2 (g) Question 3.3: In the table above, circle the leftover reactants. What is the limiting reactant for each ratio H 2 :O 2? H 2 :O 2 Ratio Limiting Reactant 4:1 3:1 2:1 1:1 1:2 17
Laboratory 3 How could a chemist determine the ratio of H 2 :O 2 that does not have leftover reactants? They need a test to determine whether H 2 (g) or O 2 (g) remains. This is actually a simple experiment; if a wooden splint is lit, extinguished, and then placed in a container of hydrogen gas while still glowing, it produces a loud pop. However, if the glowing splint is placed in a container of oxygen gas, it reignites. Question 3.4: Suppose H 2 (g) and O 2 (g) are combined in the ratios shown in the previous table to produce H 2 O. For which ratios of H 2 :O 2 would the splint test result in a pop? For which ratios of H 2 :O 2 would the splint reignite? How does this analysis lead the chemist to the correct chemical equation? In Part D you will apply quantitative reasoning to determine the ratio of potassium iodide reacting with lead nitrate in your chemical equation. This will be accomplished by mixing drops of each solution in different ratios and then testing the supernatant (the liquid above the precipitate) for excess ions. Instead of doing a test with a glowing splint, however, you will determine whether there are excess ions by checking whether a precipitate can be formed with the ions (that may) be leftover. The solution in which neither ion is in excess indicates that the ions were added in the correct ratio to exactly form the product. Note that having neither ion in excess requires careful measurement of amounts of the two solutions. A very small amount of precipitate will form while testing the supernatant liquid if there is a slight excess of either reactant. It is therefore important not only to measure carefully, but also to note the amount of precipitate when testing for the ion in excess. Can you apply quantitative reasoning from Part D to determine the chemical equation? Try question 3.5. 18
Development of an Equation Question 3.5: When solutions of barium chloride and sodium sulfate are mixed, a white precipitate of barium sulfate forms: Barium chloride + sodium sulfate $ barium sulfate (a white precipitate) + sodium chloride tested: Drops of Barium Chloride Drops of Sodium Sulfate Drops of Water Tube 1 6 2 22 Tube 2 6 3 21 Tube 3 6 6 18 Tube 4 6 12 12 Tube 5 6 18 6 Observation White precipitate formed White precipitate formed White precipitate formed White precipitate formed White precipitate formed Inference for the Barium Chloride to Sodium Sulfate Ratio A 3:1 ratio produces barium sulfate. A 2:1 ratio produces barium sulfate. A 1:1 ratio produces barium sulfate. A 1:2 ratio produces barium sulfate. A 1:3 ratio produces barium sulfate. These results are inconclusive and do not narrow down the barium chloride to sodium sulfate ratio in the balanced chemical equation because every reaction produces barium sulfate. However, by testing each supernatant (the liquid above each precipitate) it is possible to determine which ions are leftover and in excess. In the table below, identify the ions present in the supernatant (Ba 2+, Cl, Na + 4 ) prior to additional barium chloride or sodium sulfate. Tube 1 Tube 2 Tube 3 Tube 4 Tube 5 Barium Chloride Added to Supernatant Observation No additional precipitate formed No additional precipitate formed No additional precipitate formed White precipitate formed White precipitate formed Sodium Sulfate Added to Supernatant Observation White precipitate formed White precipitate formed No additional precipitate formed No additional precipitate formed No additional precipitate formed Ions Present in the Supernatant Prior to Adding Additional Reagents Question 3.6: Use the information in the table above to determine the empirical formula of barium sulfate. Support your reasoning. 19
Laboratory 3 Question 3.7: Write a balanced chemical equation for the reaction that occurs when barium chloride and sodium sulfate are mixed. Materials Required Equipment 15 4-mL test tubes medicine dropper micropipet micro spatula water bath (100-mL beaker) Bunsen burner, wire gauze thin-walled rubber tubing ring stand, ring Common Equipment centrifuge Chemicals 0.1 M lead nitrate, Pb(NO 3 ) 2 0.1 M potassium nitrate, KNO 3 3 M nitric acid, HNO 3 5% thioacetamide, CH 3 CSNH 2 3% hydrogen peroxide, H 2 O 2 1,2-dichloroethane, C 2 H 4 Cl 2 Cautions be sure to wash your hands thoroughly. 1,2-dichloroethane (dichloroethane) is a suspect carcinogen and may be irritating to the respiratory tract. Heating thioacetamide solutions produces H 2 S, an inhalation poison. Do this procedure well within the hood. Fume hoods must be on. Procedure solution. Always take drops from these tubes. A. Initial observations. In this section of the experiment you will observe the reaction that occurs when solutions of potassium iodide and lead nitrate are mixed. 1. Add 5 drops of each solution to a 4-mL test tube and observe what happens. It may be helpful to centrifuge the test tube after the reaction. Record your observations in your notebook. Save this tube for use in Part C. B. Qualitative tests for ions. In Part C we will use chemical tests to determine which ions are present in the product of anions and cations present in this experiment react. That is the task for this section. 20
Development of an Equation Procedure for Testing Anions In this section of the experiment, you will see what the reactions of the anions (iodide and nitrate) look like when they react with hydrogen peroxide. Be sure to make careful observations in your notebook. Later, you will use this test to determine which anion is present in a solution by repeating the experiment and comparing your observations. 2. Add 5 drops of KI, 5 drops nitric acid, 10 drops dichloroethane, and 5 drops hydrogen peroxide to a clean 4-mL test tube. Stopper and shake the test tube. You may decide to centrifuge the solution in the test tube to produce two distinct layers. 3. Add 5 drops of KNO 3, 5 drops nitric acid, 10 drops dichloroethane, and 5 drops hydrogen peroxide to a clean 4-mL test tube. Stopper and shake the test tube. You may decide to centrifuge the solution in the test tube to produce two distinct layers. Procedure for Testing Cations In this section of the experiment you will see what the reactions of the cations (lead and potassium) look like when they react with H 2 S. Be sure to make careful observations in your notebook. Later, you will use this test to determine which cation is present in a solution by repeating the experiment and comparing your observations. 4. To each of two clean, dry 4-mL test tubes, add 20 drops of thioacetamide and 2 drops of nitric acid. 5. To one tube add 5 drops Pb(NO 3 ) 2, and to the other 5 drops KNO 3. Heat both tubes in a water bath for at least 5 minutes and record your observations. On heating, the thioacetamide will decompose to form H 2 S. C. Identification of ions in the precipitate. In this section of the experiment you will test the product that formed when potassium iodide and lead nitrate solutions were mixed. To accomplish this you will re-dissolve the solid and determine which ions it is comprised of. Formation of Precipitate 6. Once again, add 5 drops of potassium iodide and 5 drops lead nitrate to a 4-mL test tube. (You may simply use the sample from Part A if you still have it.) Centrifuge the test tube, and then carefully remove and discard the clear aqueous layer. Wash the precipitate: add 1 ml of distilled water, stir, centrifuge the test tube, and discard the clear aqueous layer. Repeat the washing. Test for Anions 7. Add 5 drops of nitric acid, 10 drops of dichloroethane, and 5 drops of hydrogen peroxide to the solid in the test tube. Stopper and shake the test tube. 8. If necessary, alternate holding the test tube in a stream of hot tap water for 20 seconds and shaking for 40 seconds until the yellow solid is completely dissolved. Centrifuge the test tube. Record your observation of the dichloroethane layer and compare it with 21
Laboratory 3 Test for Cations 9. Transfer the less-colored aqueous layer to a clean 4-mL test tube using a micropipet. Discard the remaining dichloroethane layer in a waste beaker. 10. Heat the aqueous solution in a water bath to destroy any unreacted hydrogen peroxide. Continue heating until the solution is completely colorless. 11. Add 20 drops of thioacetamide to the test tube, stir, and heat in a water bath. While waiting for a reaction to occur, start on the next step. After 5 minutes, compare your results with the reaction observed in Part B. D. Determination of the ratio of ions in the precipitate. In Part C, you determined which ions are in the precipitate. In this section you will make observations and determine the ratio of these ions in the solid sample. Preparing Test Tubes with Varying Ratios of the Reactants according to the table below. Notice the table has several empty cells. Reproduce the before continuing. In each case make the amount of lead nitrate constant and the total number of drops constant. Table 3.1 Solutions to be mixed in Part D. Drops of Lead Nitrate Drops of Potassium Iodide Drops of Water Lead Nitrate: Potassium Iodide Ratio Tube 1 12 4 32 3:1 Tube 2 12 6 2:1 Tube 3 12 12 24 Tube 4 12 1:2 Tube 5 36 0 13. The same medicine dropper should be used to measure all drops. (Do not use a mi- tion, and then rinse the medicine dropper well with distilled water before changing solutions. and then centrifuge the test tubes. The centrifuge must be balanced using pairs of test in your notebook. 22
Development of an Equation Determining which Test Tube, 1 5, Does Not Have an Excess of Cation or Anion When the ions in the potassium iodide and lead nitrate solutions were mixed they may have produced different amounts precipitate, as noted in step 4. For a given ratio, were any of the ions in excess? If so, they will still be in the supernatant. This can be tested for by taking a portion of the supernatant and seeing if it is still possible to form additional precipitate. Obviously, if more precipitate is formed, excess ions must have been present in supernatant (step 2). 16. To test a given supernatant for excess ions, use a medicine dropper and remove 5 drops of the solution into a different test tube. Be careful not to transfer any of the precipitate. You may need to re-centrifuge the tube to avoid transferring any solid. If you are testing for the presence of potassium or iodide ions in the supernatant, add 5 drops of lead nitrate and record the relative amount of precipitate that forms (if any). If you are testing for the presences of lead or nitrate ions in the supernatant, add 5 drops of potassium iodide. and make comparisons. Waste Disposal All solutions containing lead ions and dichloroethane must be collected in a beaker at your Your lab instructor will neutralize the solution with sodium hydroxide and dispose of the total volume in the appropriate container. Dispose of the used test tubes in the uncleanable glass waste container. 23
24 Laboratory 3
Laboratory 3 Report Sheet Development of an Equation Name: Date: TA Name: A. Initial observations. What atoms, ions, or molecules do you think were in the potassium iodide and lead nitrate solutions before they were mixed? List your observations upon mixing: Is there evidence a chemical reaction occurred? Support your answer. Is there evidence that a gas was formed during the reaction? B. Observations of qualitative tests for ions in solution. Testing the Anions Observations for the iodide reaction with H 2 O 2 : Observations for the nitrate reaction with H 2 O 2 : Testing the Cations Observations for the lead reaction with H 2 S (thioacetamide): Observations for the potassium reaction with H 2 S: 25
Laboratory 3 Report Sheet C. Identification of ions in the precipitate that forms when lead nitrate solution is added to potassium iodide solution. Testing for Anions Observations for the reaction with H 2 O 2 : Inferences: What anions are in the precipitate? How do you know? Testing for Cations Observations for the reaction with H 2 S: Inferences: What cations are in the precipitate? How do you know? Pattern of Reactivity What kind of reaction (combination, decomposition, combustion or exchange) occurred when potassium iodide and lead nitrate solutions are mixed? Support your reasoning. 26
Development of an Equation Report Sheet continued D. Determining the exact ratio of ions in the precipitate. Testing of Precipitate Observations on Initial Precipitate Formed (relative amount, etc.) Tube 1 Tube 2 Tube 3 Tube 4 Tube 5 Testing of Supernatant Observations Inferences: Which ions were in the supernatant? List both cations and anions. Tube 1 Tube 2 Tube 3 Tube 4 Tube 5 27
Laboratory 3 Report Sheet Using all of your observations, what is the correct ratio of ions? Support your conclusion. Formula of the Precipitate Write the balanced equation for the reaction that forms this precipitate. What is your reasoning? You must convince the reader by clearly summarizing your observations and how they lead to your inferences. This is the key point in this lab! 28
Development of an Equation Report Sheet continued Draw particle diagrams consistent with your balanced equation. Before the reaction After the reaction Metals like lead frequently have different oxidation states, e.g., being +2 or +3, and these oxidation states may change during a chemical reaction. What is the oxidation state of lead in the precipitate? Support your reasoning with your experimental observations. 29
Laboratory 3 Report Sheet 30
Laboratory 3 Answer Clinic Development of an Equation Question 3.1: Write a balanced chemical equation for N 2 reacting with O 2 to form laughing gas. Answer 2 N 2 (g) + O 2 (g) $ 2 N 2 O(g) In the boxes below, draw a particle diagram showing the reactants (before) and the products (after) the reaction. The drawing should be consistent with your balanced equation. Reactants (Before) Products (After) 2N 2 (g) + O 2 (g) 2N 2 O(g) Question 3.2: Is the chemical reaction N 2 (g) + O 2 (g) $ 2 NO(g) an example of a combination, decomposition, combustion, or exchange reaction? Answer Combination reactions, which describe the reaction of two or more reactants to form one product. Questions 3.3: What is the limiting reactant for each ratio H 2 :O 2? Answer H 2 :O 2 Ratio Limiting Reactant 4:1 O 2 (g) 3:1 O 2 (g) 2:1 1:1 H 2 (g) 1:2 H 2 (g) 31
Laboratory 3 Answer Clinic Question 3.4: Suppose H 2 (g) and O 2 (g) are combined in the ratios shown in the table on page 17 to produce H 2 O. For which ratios of H 2 :O 2 would the splint test result in a pop. For which ratios of H 2 :O 2 would the splint reignite? How does this analysis lead the chemist to the correct chemical equation? Answer When the H 2 :O 2 = 2 there will not be any H 2 or O 2 left over. If H 2 :O 2 > 2 there is excess H 2 and the O 2 is the limiting reactant, and the splint tests produce a pop. If H 2 :O 2 < 2 there is excess O 2 and the H 2 is the limiting reactant, and the splints will be reignited. Question 3.5: Identify the ions present in the supernatant (Ba 2+, Cl, Na + 4 ) prior to additional barium chloride or sodium sulfate. Answer Tube 1: Ba 2+, Na +, Cl Tube 2: Ba 2+, Na +, Cl Tube 3: Na +, Cl Tube 4: Na +, Cl, SO 4 2 Tube 5: Na +, Cl, SO 4 2 Question 3.6: Use the information in the table on page 19 to determine the empirical formula of barium sulfate. Support your reasoning. Answer By testing the supernatant it appears that tube 3 does not have excess barium ions or excess sulfate ions. This tube was prepared using equal drops of barium chloride and sodium sulfate. This suggests that the ratio of barium to sulfate is 1:1, consistent with an empirical formula of BaSO 4. Question 3.7: Write a balanced chemical equation for the reaction that occurs when barium chloride and sodium sulfate are mixed. Answer BaCl 2 (aq) + Na 2 SO 4 (aq) $ BaSO 4 (s) + 2 NaCl(aq) 32