Chemical Reactions 2

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This document outlines the purpose, procedure, and experiments for a chemistry lab activity on chemical reactions. The purpose is to observe different types of chemical reactions, identify reaction types, predict products, and write and balance chemical equations. The procedure involves setting up reactions in test tubes, recording observations, and cleaning up. Experiment 1 mixes zinc and copper nitrate, Experiment 2 mixes ammonium hydroxide and copper nitrate, and Experiment 6 dissolves iron chloride and sodium hydroxide in water.

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TOPIC 4 LAB ACTIVITY 1 Chemical Reactions By: Alberto Cánovas, Noelia Fajardo & Silvia Jiménez

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Thanks everybody! The lesson plans presented were actually outdated and can still be improved. I was also a college student when I did these. There were minor errors but the important thing is, the structure and flow of activities (for an hour-long class) are included here. I appreciate all of your comments! Please like my fan page on facebook search for JUNNIE SALUD. *The detailed LP for English is from Ms. Juliana Patricia Tenzasas. I just revised it a little. For questions about education-related matters, you can directly email me at mr_junniesalud@yahoo.com

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Chemistry 201 Laboratory- Harold Washington College Reactions of Copper and Percent Recovery Introduction In Chapter 4 of your textbook, we have covered the 3 main classes of reactions in aqueous solution: precipitation reactions, acid-base reactions, and oxidation-reduction reactions. These can be summarized as follows: • In Precipitation reactions: a solid product forms from aqueous reactants. The product can be predicted by using solubility rules. • In Acid-base neutralization reactions: an acid (proton donor) reacts with a base (proton acceptor) to form an ionic salt and sometimes water. You can recognize this type of reaction by the presence of an acid and base in the reactants. • In Oxidation-reduction reactions (or redox for short): electrons are transferred between reactants. So you will always see changes in oxidation number within this type of reaction. In this lab, we will be examining these reactions by carrying out a sequence of reactions starting from metallic copper in the form of a wire. According to the Law of Mass Conservation, covered in section 2.3 of your textbook, mass is neither created nor destroyed. So since no copper is added or removed during the course of these reactions, you should be able to quantitatively recover all of the copper metal you started with at the end of these reactions (if everything goes well). You need to carefully observe and record your observations. You should be able to carefully describe the reactions in terms of color change (color of solution or solid forming), gas formation (as bubbles), heat generation (change in temperature), or precipitate formation, throughout the experiment. At the end of the reactions series, you will collect the copper recovered and calculate the percent recovery (% yield) of copper using this formula: % yield = actual amount of Cu recovered x 100 % original amount of Cu Materials and Equipment • Piece of copper wire (about 0.5 g) • concentrated HNO3 (aq) • 3.0M NaOH • 6.0M H2SO4 • solid zinc • 2-250 mL beakers, 400 mL beaker • 50 mL graduated cylinder • boiling chips • stirring rod • iron ring and ring stand, wire gauze, • Bunsen burner • evaporating dish • electronic balance Chemistry 201 Laboratory- Harold Washington College SAFETY NOTES • Safety goggles should be worn at all times • Sulfuric acid, nitric acid, and sodium hydroxide are corrosive chemicals and should be handled with care. • Gloves are to be worn when handling these chemicals. • Waste should be put in the designated container in the hood. Never throw chemicals down the drain. PROCEDURE 1. Obtain a piece copper wire, measure its mass to the nearest 0.01 g, and place it in a 400 mL beaker. 2. IN THE HOOD, add 4-5 mL of concentrated nitric acid, HNO3, to the beaker. (Caution: Be careful not to get any of the nitric acid on yourself. If you do, wash it o ...

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Water Testing Lab Background Information: Chemists can detect and identify ions in water solution in several different ways. In this lab you will use some chemical tests to check for the presence of certain ions in aqueous solution. Positively charged ions are called cations; negatively charged ions are called anions. The tests you will perform are confirming tests. If the test is positive, it confirms that the ion in question is present. In each confirming test you will look for a change in solution color, or the appearance of an insoluble material called a precipitate. A negative test (no color change or precipitate) doesn’t necessarily mean the ion in question is not present. The ion may simply be present in such a small amount that the color or precipitate cannot be seen. Purpose: § To use chemical tests to determine the presence of ions in three different water samples. Objectives: § To test for the presence of cations iron (III) Fe3+, calcium Ca2+, copper Cu+, as well as the anions chloride Cl-, phosphate PO43-, and sulfate SO42-. § To perform each of the following confirming tests on three different samples: 1. A reference solution (known to contain the ion of interest) 2. Tap water (which may or may not contain the ion) 3. A control (distilled water, known not to contain the ion) § Test for water hardness using ion-testing strips. § Complete a confirmative ion-testing using a complete ion-testing strip. Materials: § 10 test tubes § Test tube rack § 10 mL graduated cylinder § Test reagent solutions in dropper bottles o Potassium thiocyanate (KSCN) o Potassium oxalate (K2C2O4) o Acetic acid (HC2H3O2) o Silver nitrate (AgNO3) o Barium Chloride (BaCl2) § Reference ion solutions o Iron (III) Nitrate - Fe(NO3)3 o Calcium Chloride – CaCl2 o Iron (II) Sulfate – FeSO4 Procedures: PART ONE 1. Safety goggles should be worn at ALL times during this lab exercise. If you spill any of the reagents on your hands, wash them immediately. 2. Clean five test tubes with water and dry all glassware before starting the experiment. 3. In this lab, you are testing water samples for the presence of four different ions: Ca2+, Fe3+, Cl-, and SO42-. For EACH of the test procedures you will test a reference solution (known to contain the ion), a control (distilled water-known not to contain the ion), and your water samples brought from home. Follow the diagram below and be sure to keep track of your water samples: 4. Measure 1 pipette full (about 2 mL) of reference solution, control, and the three water samples into five clean test tubes. Label the three test tubes Reference (R), Control (C), and Water Sample #1, 2, & 3. 5. Follow the table below to test for each ion: 6. As you perform each test, record your observations in the data table below. 7. When you have completed an ion test, wash your test tubes with water and dry them thoroughly before moving on to the next ion test. P.

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Lab 9: Chemical Reactions II Pre-lab Questions 1. What is a limiting reagent? 2. A student used 7.15 g of CaCl2 and 9.25 g of K2CO3 to make CaCO3. The actual yield was 6.15 g of CaCO3. Calculate the limiting reagent and the percent yield. Experiment: Synthesis of Garden Lime Procedure **Take photographs of your experiment set up and your results. Submit them with your laboratory report.** 1. Table 1 provides an example set of data for 1.0 g CaCl2. 2. For Trial 1, weigh into a 250 mL beaker the amount of calcium chloride (CaCl2) shown in Table 1. Record the exact mass you weigh out in the Trial 1 column of the Data section. 3. Measure 50.0 mL of distilled water into a 100 mL graduated cylinder. Pour the water into the 250 mL beaker with the calcium chloride. 4. Stir the solution with a stirring rod until all of the calcium chloride is dissolved. 5. Weigh out 2.5 g of potassium carbonate (K2CO3) in a 50 mL beaker. Record the exact mass in the Data section. 6. Measure 25.0 mL of distilled water into a 100 mL graduated cylinder. Add the water into the 50 mL beaker containing the potassium carbonate. 7. Stir the potassium carbonate in the distilled water with a stirring rod until it is all dissolved. 8. Pour the K2CO3 solution into the 250 mL beaker that has the CaCl2 solution. Rinse the beaker that contained the K2CO3 with a few mL of water and add this to the CaCl2 solution. Stir the mixture. 9. As soon as the reaction begins, record your observations in the Data section. Continue stirring until you see no more precipitate forming. 10. Set up the funnel in the Erlenmeyer flask as shown in Figure 2. HINT: Do NOT begin filtering yet! 11. Zero the scale and weigh a piece of filter paper and a watch glass. Record the masses of both items in the Data section. 12. Prepare a filtering funnel as shown in Figure 2: fold a piece of filter paper in half twice to make quarters, and open the paper to make a small cone (three quarters are open on one side and one quarter is on the opposite side). Place the paper cone into the funnel and hold it in place with your fingers. Pour a small amount of distilled water through the paper to secure it inside the funnel. 13. Filter the mixture by pouring it into the filter paper in the funnel. Use the stirring rod and distilled water in a wash bottle to transfer the entire solid into the filter paper. HINT: For best results, be sure to transfer all of the precipitate into the filter paper. Use a rubber policeman if it is available to help with the transfer. 14. Rinse the remaining solid in the filter paper twice with distilled water from a wash bottle to rinse off excess sodium chloride (NaCl). After all the liquid has filtered through, rinse the product with approximately 5 mL of ethanol to aid in its drying. Allow the ethanol to completely finish filtering through the paper. 15. Remove the filter paper carefully so as to not lose any product. Gently unfold the filter paper and lay it flat on the pre-weighed wat ...

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phychem1phychem2 Lab 9: Chemical Reactions II Pre-lab Questions 1. What is a limiting reagent? 2. A student used 7.15 g of CaCl2 and 9.25 g of K2CO3 to make CaCO3. The actual yield was 6.15 g of CaCO3. Calculate the limiting reagent and the percent yield. Experiment: Synthesis of Garden Lime Procedure **Take photographs of your experiment set up and your results. Submit them with your laboratory report.** 1. Table 1 provides an example set of data for 1.0 g CaCl2. 2. For Trial 1, weigh into a 250 mL beaker the amount of calcium chloride (CaCl2) shown in Table 1. Record the exact mass you weigh out in the Trial 1 column of the Data section. 3. Measure 50.0 mL of distilled water into a 100 mL graduated cylinder. Pour the water into the 250 mL beaker with the calcium chloride. 4. Stir the solution with a stirring rod until all of the calcium chloride is dissolved. 5. Weigh out 2.5 g of potassium carbonate (K2CO3) in a 50 mL beaker. Record the exact mass in the Data section. 6. Measure 25.0 mL of distilled water into a 100 mL graduated cylinder. Add the water into the 50 mL beaker containing the potassium carbonate. 7. Stir the potassium carbonate in the distilled water with a stirring rod until it is all dissolved. 8. Pour the K2CO3 solution into the 250 mL beaker that has the CaCl2 solution. Rinse the beaker that contained the K2CO3 with a few mL of water and add this to the CaCl2 solution. Stir the mixture. 9. As soon as the reaction begins, record your observations in the Data section. Continue stirring until you see no more precipitate forming. 10. Set up the funnel in the Erlenmeyer flask as shown in Figure 2. HINT: Do NOT begin filtering yet! 11. Zero the scale and weigh a piece of filter paper and a watch glass. Record the masses of both items in the Data section. 12. Prepare a filtering funnel as shown in Figure 2: fold a piece of filter paper in half twice to make quarters, and open the paper to make a small cone (three quarters are open on one side and one quarter is on the opposite side). Place the paper cone into the funnel and hold it in place with your fingers. Pour a small amount of distilled water through the paper to secure it inside the funnel. 13. Filter the mixture by pouring it into the filter paper in the funnel. Use the stirring rod and distilled water in a wash bottle to transfer the entire solid into the filter paper. HINT: For best results, be sure to transfer all of the precipitate into the filter paper. Use a rubber policeman if it is available to help with the transfer. 14. Rinse the remaining solid in the filter paper twice with distilled water from a wash bottle to rinse off excess sodium chloride (NaCl). After all the liquid has filtered through, rinse the product with approximately 5 mL of ethanol to aid in its drying. Allow the ethanol to completely finish filtering through the paper. 15. Remove the filter paper carefully so as to not lose any product. Gently unfold the filter paper and lay it flat on t.

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Unit 4 Lab: StoichiometryCHEM 1806 Purpose: To practice the concept of Stoichiometry Materials: Student Supplied – Distilled water, paper towels, small paper cup, mug LabPaq materials: 100 mL glass beaker, Funnel, 25 mL cylinder, Goggles, Digital Scale, Weighing boat, 2.5 g Calcium chloride, dehydrate, Filter paper, 2 g Sodium carbonate. Procedure: 1. While wearing goggles, weigh out 1.0 g CaCl2 and put it into the 100 mL beaker. Record the actual mass of CaCl2 you are using in Table 1. 2. Add 25 mL of distilled water to the beaker and stir. This is now the “Calcium chloride solution” 3. Calculate how much sodium carbonate you will need for a full reaction. Write that calculation in the Results and Discussion section. 4. Weigh out the calculated amount of Na2CO3 and put it in a small paper cup. (Record actual mass in Data Table 1) 5. Add 25 mL distilled water to the paper cup and stir to make “sodium carbonate solution.” 6. Pour the sodium carbonate solution into the beaker with the calcium chloride solution. 7. Record observations. Write out the reaction of what occurred in the results section. (The solid is calcium carbonate.) 8. Fold the filter paper in half, then in quarters. Record the mass of the filter paper. 9. Place the filter paper in the funnel and open one side. 10. Place the funnel in a mug or other cup to balance it. 11. Carefully swirl and pour the contents of the beaker into the filter paper. (Don’t let the solution overflow.) 12. Rinse the beaker with 1-3mL of distilled water and pour that water into the filter paper as well. (If there is a little solid left in the beaker, leave it. It is experimental error.) 13. When the liquid has drained, lay the filter paper and precipitate on a plate or paper towels to allow it to dry. (mine took overnight, some people take 3 days) 14. When dry, weigh the precipitate + filter paper, record it, and calculate the mass of the precipitate. 15. In the discussion, calculate the theoretical yield and percent yield. (4 pts) Data: Table 1: Data Mass of CaCl2 Mass of Na2CO3 needed Actual mass of Na2CO3 used Observation of reaction Mass of filter paper Mass of filter paper + precip Actual yield of Calcium Carbonate 1. (3 pts) Calculate how much Na2CO3 you need to completely react your mass of CaCl2. (Show that calculation here) 2. (2 pts) Write out the reaction that occurred to form calcium carbonate. Remember to identify the phases and balance the equation. 3. (2 pts) Calculate the actual yield of Calcium Carbonate. (Show your work) 4. (2 pts) Calculate the theoretical yield of Calcium Carbonate. (Show your work) 5. (2 pts) Calculate your percent yield of Calcium Carbonate. (Show your work) Insert a picture of your filter paper with yield for 2 pts extra credit. ;-) When you are finished, save the file as LASTNAME_Stoich and upload it to the Dropbox for Unit 4 Labs. Unit 4 Lab: Ionic ReactionCHEM 1806 Purpose: To observe several ionic reactions and practice net ionic equations ...

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