LAB REPORT 3

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University of Texas, Rio Grande Valley *

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1112

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Chemistry

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Dec 6, 2023

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LAB REPORT Experiment 3: Reaction rates Chem 1112-03 The University of Texas Rio Grande Valley Professor Bhupendra Srivastava Name: Daniel Lara Lab partners: America perez Ally diaz Annette Obregon

Objectives  Describe and determine the factors that affect the rate of reaction, such as temperature, concentration and chemical structure.  Define the term catalyst Introduction In experiment 3 reaction rate of certain reactions will be analyzed behind the study of “kinetics”. Kinetic studies include not only the determination of the rate of a reaction, but also the mechanism by which the reaction progresses. Many factors affect the rate of a reaction such as temperature, concentration, and chemical structure. These factors are studied in and manipulated in order to optimize the speed and productivity of the reaction. Procedure: In preparation for Part III, gently heat (~55-60°C) a beaker of water on a hot plate. Add a thermometer to check the temperature. Part I: Effect of Changing Acid 1 . Arrange four small test tubes in a test tube rack. Fill each one with 2 mL (40 drops) of 3 M H 2 SO 4 sulfuric acid, 6 M HCl hydrochloric acid, 6 M CH 3 COOH acetic acid, and 6 M H 3 PO 4 phosphoric acid,

2. At the same time, place a 1-cm strip of magnesium ribbon, Mg into each test tube. Notice the reaction rates. Record your observations in Table 1 of the Data Sheet. 3. Remove any remaining magnesium strips before disposing of the acid in the acid waste container in the hood. The magnesium strips should be rinsed with water and placed in the solid waste container in the fume hood. Part II: Effect of Changing Metal 1. Arrange three small test tubes in a test tube rack. Fill each one with 2 mL (40 drops) of 6M HCl , as shown below.

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2. At the same time, place a 1-cm strip of zinc , magnesium , and copper separately into the test tubes. Notice the reaction rates. Record your observations in Table 2 of the Data Sheet. 3. Remove the metal strips before disposing of the acid in the acid waste container in the hood. The metal strips should be rinsed with water and placed in the solid waste container in the fume hood. Part III: Temperature of Reaction 1. Prepare a hot water bath (~55-60°C) and an ice bath large enough to hold two test tubes each. 2. Arrange six medium size test tubes in a test tube rack. Add 2 mL of 0.1 M Na 2 S 2 O 3 sodium thiosulfate, into each of the three labeled test tubes. Add 2 mL of 0.1 M HCl into each of the remaining three labeled test tubes. Note : When using dropper bottles, 20 drops = 1 mL. 3. Place one tube containing Na 2 S 2 O 3 and one tube containing HCl into the warm water bath, place another pair of tubes containing Na 2 S 2 O 3 and HCl into an ice bath. These should remain in the bath for about 5 minutes to reach thermal equilibrium. The third set of tubes remain at room temperature. 4. Set the timer to 0:00. Pour the room temperature HCl solution into the room temperature Na 2 S 2 O 3 solution and START TIME. Mix the solution with a stirring rod for several seconds. STOP TIME when the solution appears cloudy (the formation of sulfur). 5. Record the room temperature and the amount of time required for the reaction to occur in Table 3.

6. Reset the timer to 0:00. From the ice bath pour the HCl solution into the Na 2 S 2 O 3 solution and START TIME. Mix the solution with a stirring rod for several seconds and return the reaction mixture to the ice bath. STOP TIME when the solution appears cloudy. 7. Record the temperature of the ice bath and the amount of time required for the reaction to occur in Table 3. 8. Reset the timer to 0:00. From the warm water bath, pour the HCl solution into the Na 2 S 2 O 3 solution and START TIME. Mix the solution with a stirring rod for several seconds and return the reaction mixture to the warm water bath. STOP TIME when the solution appears cloudy. 9. Record the temperature of the warm water bath and the amount of time required for the reaction to occur in Table 3. 10. Dispose of all the reaction solutions in the waste container in the hood.

Part IV: Presence of a Catalyst The Reaction: Decomposition of hydrogen peroxide, H 2 O 2 , is catalyzed by manganese (IV) oxide, MnO 2 . Place approximately 2 mL of 3% H 2 O 2 in a small test tube. Add one small crystal or a pinch (use the tip of spatula) of MnO 2 to the solution. Record your observations in Table 4 of the Data Sheet. Part V: Concentration of Reactants Triiodide forms after all the acid is consumed and forms a dark blue complex with starch. 1. Calibrate the dropper bottles by determining the number of drops in 1 mL using a graduated cylinder. If instructor approves, 20 drops = 1 mL will simplify the calculations. Record your results in Table 5 of the Data Section. 2. Arrange five medium size test tubes in a test tube rack and label them 1 through 5. Add the HIO 3 and water as outlined in the table below. Note: The total volume in all of the test tubes should be the same. 3. Add 1 drop of starch to each test tube.

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4. Set the timer to 0:00. Measure 1.0 mL of 0.01 M H 2 SO 3 solution in a separate test tube and quickly add it to test tube #5 and START TIMER. Immediately swirl the solution in the test tube. Stop the timer when the deep blue starch complex appears. Be prepared; the color change is sudden for test tube #5. Record the elapsed time in Table 5. Test tube #5 is the fastest one, so if there is a problem with the reaction, students can notify instructor. 5. Repeat steps 3 and 4 for the test tubes 1-4. 6. Dispose of the solutions in the waste container in the hood. Clean the test tubes with two tap water rinses, followed by two DI water rinses. The two tap water rinses should be disposed of in the waste container in the hood also.

Results and Discussion Summary:  Our investigation revealed a hierarchy of acid reactivity, with HCI displaying the highest reactivity and the quickest appearance of hydrogen gas bubbles. H2SO4 ranked second in reactivity, consistent with its status as a strong and highly reactive acid, leading to rapid magnesium consumption. In contrast, CH3COOH exhibited the slowest reaction rate due to its weaker acidity compared to H3PO4.  In terms of metal reactivity, we ordered them as Mg (most reactive), Zn, and Cu, aligning with the expected activity series where Mg surpasses Zn, and Cu falls below both.  Furthermore, our findings illustrated that elevated temperatures prompt faster reactions and reduced reaction times, demonstrated by the swift nine-second reaction in the warm water bath experiment.  The addition of a catalyst expedited hydrogen peroxide decomposition, evidenced by the immediate appearance of oxygen bubbles when MnO2 was introduced. In contrast, without a catalyst, hydrogen peroxide would have decomposed slowly at room temperature.  Lastly, we observed a direct correlation between reactant concentration and reaction rate. As the molar concentration of HIO3 increased, the reaction rate accelerated, resulting in a shorter reaction time, as the rate increase facilitated quicker reaction completion. Conclusion: The experimental data gathered in this laboratory investigation clearly demonstrates that reaction rates exhibit an increase with higher temperatures, elevated concentrations, and the introduction of catalysts. Furthermore, the nature of the reactants plays a pivotal role, with more reactive species exhibiting swifter reactions. Our observations during this experiment substantiate the influence of each factor, as varying these conditions yielded distinct reaction rates. Catalysts were observed to initiate reactions without undergoing consumption or production in the process. Additionally, the reactivity of acids was seen to correlate with their strength, while the reactivity of metals adhered to the established activity series. Our investigation enabled us to deduce the concentration of a solution by considering its initial concentration and the final volume of the products.

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Concentration Temp (C°) Time (sec) Time (sec) Temp Vs Time

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