Experiment 23 laboratory report PDF

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Beran's Manual Laboratory Report- Experiment 23- Factors Affecting Reaction Rates...

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Mary Johnson 02/27/20 Experiment 23 Laboratory Report Factors Affecting Reaction Rates Lab Partner: Kevin Hernandes Objectives: The purpose of this experiment is to study the various factors that affect the rates of chemical reactions. Introduction: Chemical kinetics is the study of chemical reaction rates, how reaction rates are controlled, and the pathway or mechanism by which a reaction proceeds from its reactants to its products. Reaction rates vary from the very fast, in which the reaction, such as the explosion of a hydrogen oxygen mixture, is essentially complete in microseconds or even nanoseconds, to the very slow, in which the reaction, such as the setting of concrete, requires years to complete. The rate of a chemical reaction may be expressed as a change in the concentration of a reactant or product as a function of time- the greater the change in the concentration per unit of time, the faster the rate of the reaction. Other parameters that can follow the change in concentration of species as a function of time in a chemical reaction or color, temperature, pH, gas evolution, odor, and conductivity. We will be investigating 4 to 5 factors that can be controlled to affect the rate of a chemical reaction: - nature of the reactants - Temperature of the chemical system - Presence of a catalyst - Concentration of the reaction - Surface area of the reactants Content/Definitions: More reactive substances, like the reaction of sodium metal and water, is a very rapid, exothermic reaction, unlike the corrosion of iron, which is much slower. A 10°C rise in temperature doubles the rate of a chemical reaction. The added heat not only increases the number of collisions between reactant molecules but also, increases their kinetic energy. On collision of the reactant molecules, this kinetic energy is converted into an internal energy that is distributed throughout the collision system. This increased internal Energy increases the probability for the weaker bonds to be broken and the new bonds to be formed. A catalyst increases the rate of a chemical reaction without undergoing any net chemical change. Catalysts can affect only one specific reaction, or the entire general reactions. An increase in the concentration of a reactant generally increases the reaction rate.

Generally speaking, the greater the exposed surface area of the reactant, the greater the reaction rate. B. - The oxidation reduction reaction that occurs between hydrochloric acid and sodium thiosulfate, Na2S2O3, produces insoluble sulfur as a product:

 > 2HCl(aq) + NA2S2O3(aq) — S(s) +  SO2(g) + 2NaCl(aq) +  H2O(l) C. - The reaction rate for the oxidation reduction reaction between the oxalic acid, and potassium permanganate, is measured by recording the time elapsed for the purple color of the permanganate ion, to disappear and reaction.

 2KMnO4(aq) +  3H2SO4(aq) —  > 5H2C2O4(aq) +  2MnSO4(aq) +  K2SO4(aq) +  8H2O(l) 10CO2(g) +

Procedure: A. Nature of the Reactants 1. Different acids affect reaction rates. Half fill a set of four labeled small test tubes with 3 M H2SO4, 6 M HCl, 6 M CH3COOH, and 6 M H3PO4, in a test tube rack. Submerge a 1 cm strip of magnesium ribbon into each test tube. Compare the reaction rates and record your observations. 2. Different metals affect reaction rates. Ha fill a set of three labeled small test tubes with 6 M HCl. Submerge 1 cm strips of zinc, magnesium, and copper separately into test tubes. Compared with the reaction rates of each metal in HCl and record your observations. B. Temperature of the reaction: hydrochloric acid sodium thiosulfate reaction system. 1. Prepare the solutions play pet 2 mL of 0.1 molar sodium thiosulfate into each of us set of three 150 mm, clean test tubes. Into a second set of three 150 mm test tubes, but I put 2 mL of 0.1 molar hydrochloric acid. Label each set of test tubes. The first pair of Na2S2O3-HCl Pair test tubes is to be combined at room temperature in part B.2. Place a second pair of the test tubes in an ice water bath for part B.3. And a third pair of test tubes in a warm water bath for part B.4. Allow each pair of test tubes to establish thermal equilibrium, for about five minutes, before continuing to part B.3 and 4. 2. Record the time for reaction at room temperature. Be prepared to start time for monitoring the reaction rate. Combine the first pair of Na2S2O3-HCl pair test tubes and start time. Agitate the mixture for several seconds. Stop time when the cloudiness of the sulfur appears. Record the time lapse in room temperature, using all certain digits plus one uncertain. 3. Record the time for reaction at a lower temperature. From the ice bath for the hydrochloric acid solution into the sodium thiosulfate solution. Start time, agitate the mixture for several seconds, and return the reaction mixture to the ice bath. Stop time when the cloudiness of the sulfur appears. Record the time labs for the reaction and the temperature of the bath.

4. Record the time for reaction at the higher temperature. From the warm water bath for the hydrochloric acid solution into the sodium Thiosulfate solution and proceed as in parts B.2 and .3, record the appropriate data. 5. plot the data. Plot temperature versus time on 1/2 of a sheet of linear graph paper or by using appropriate software. C. Temperature of the reaction: oxalic acid- potassium permanganate reaction system. 1. Prepare the solutions. Into a set of three, clean 150 mm test tubes, pipette 1 mL of .01M KMnO4 (in 3 M H2SO4) and 4mL of 3 M H2SO4. Into a second set of three clean 150 mm test tubes pipette 5 mL of 0.33 M H2C2O4. 2. Record the time for reaction at room temperature. Select a KMnO4-H2SO4 pair of test tubes. Pour the H2C2O4 solution into the KMnO4 solution. Start time. Agitate the mixture. Record the time for the purple color of the permanganate ion to disappear. Record room temperature using all digits plus one uncertain digit. 3. Record the time for reaction at a higher temperature. Place a second pair of test tubes in the warm water bath (40C) until thermal equilibrium is established. Pour the H2C2O4 solution into the KMnO4 solution. Start time. Agitate the mixture for several seconds and return the reaction mixture to the warm water bath. Record the time for the disappearance of the purple color and record the temperature of the bath. 4. Record the time for reaction at the highest temperature. Repeat part C.3 but increase the temperature of the bath to about 60°C Record the appropriate data. 5. Plot the data. Plot temperature versus time on 1/2 of a sheet of linear graph paper or by using appropriate software for the three data points. D. Presence of a Catalyst 1. Add a catalyst. Place approximately 2 mL of a 3% H2O2 solution in a clean, small test tube. Add one or two crystals of MnO4 to the solution and observe. Note its instability. E. Concentration of reactants: magnesium- hydrochloric acid system. 1. Prepare the reactants. Into a set of four clean, label test tubes, pipette 5 mL of 6M HCl, 4M HCl, 3M HCl, and 1M HCl. Determine the mass separately for each solution of 41 cm strips of polished magnesium. Calculate the number of moles of magnesium in each strip. 2. Record the time for completion of the reaction. Add the first magnesium strip to the 6M HCl. Start time. Record the time for all traces of the magnesium strip to disappear repeat with the experiment with the remaining three magnesium strips for each of the different molarities of the hydrochloric acids. 3. Plot the data. Plot moles of hydrochloric acid over moles magnesium, versus time in seconds for the four tests on 1/2 of a sheet of linear graph paper by using appropriate software. F. Concentration of reactants: Iodic acid- sulfurous acid system 1. Prepare the test solutions. Review the preparation of the test solutions in table 23.1. Set up five, clean and labeled test tubes in a test tube rack. Measure the volume of the 0.01M HIO3, starch, and water with dropping pipettes. Calibrate the HIO3 dropping pipette to determine the volume per drop. Calibrate a second dropping pipette with water to determine the number of milliliters per drop. Calibrate a third dropping pipette for the 0.01 M H2SO3 solution that delivers 1 mL;

mark the level on the pipettes so that quick delivery of the 1 mL of the H2SO3 solution to each test tube can be made. Alternatively, usually calibrated 1 mL Beral pipette. 2. Record the time for the reaction. Place a sheet of white paper beside the test tube. As one student quickly transfers 1.0 mL of the 0.01 M H2SO3 to the respective test tube, the other notes the time. Immediately agitate the test tube; record the time lapse for the deep blue I-3 starch complete to appear. 3. Complete remaining reactions. Repeat part F.2 for the remaining reaction mixtures in table 23.1. Repeat any of the trials as necessary. 4. Plot the data. On 1/2 of a sheet of linear graph paper or by using appropriate software, plot for each solution the initial concentration of iodic acid, versus the time in seconds for the reaction. Correlation of Data: A. Nature of the Reactants: 1. List the acids in order of decreasing reaction rate with magnesium: 6M HCl; 3M H2SO4; 6M CH3COOH; 6M H3PO4 2. List the metals in order of decreasing reaction rate with 6M HCl: magnesium; zinc; copper 3. Identify the metals reacting in Figure 23.5 (from left to right): zinc; copper; magnesium B. Temperature of the Reaction: Hydrochloric Acid-Sodium Thiosulfate Reaction System 1. Time for Sulfur to Appear (s)

Temperature of the Reaction (C)

46.14

~22

1:32.00

~10

23.11

~40

2.

3. As temperature increases, the rate of reaction increases and appearance of sulfur quickens. 4. Estimated temperature of sulfur to occur in 20 seconds: about 35 degrees C C. Temperature of the Reaction: Oxalic Acid- Potassium Permanganate Reaction System 1.

2.

Time for Permanganate Ion to Disappear (s)

Temperature of the Reaction

20.36

~22

36.79

40

7.49

60

3. Estimation of time of disappearance of the purple permanganate ion at 55 degrees C: about 10 seconds. D. Presence of a Catalyst 1. The MnO2 catalyst speeds up the rate of evolution… causing it to fizz up and bubble, releasing the O2 into the air. 2. Balanced equation for the decomposition of H2O2: H2O2(aq) + MnO2(aq) → MnO(s) + O2(g) + H2O(l) E. Concentration of Reactants: Magnesium-Hydrochloric Acid System Concentrati on HCl

Vol. HCl (mL)

Mol HCl

Mass Mg (g)

Mol Mg

Mol HCl/mol Mg

Time (s)

6M

~5

0.137

0.020

0.000823

166.46

8.56

4M

~5

0.137

0.021

0.000864

158.56

21.82

3M

~5

0.137

0.020

0.000823

166.46

27.80

1M

~5

0.137

0.019

0.000782

175.19

2:27.89

1. 2. Time for 5 mg of Mg to react in 5 mL of 2.0 M HCl: .005 g/ 24.305= 0.000206 mol Mg 0.137 mol HCl / 0.000206 mol Mg= 665.05

Summary of Results/ Discussion: We completed multiple mini experiments within this one lab experiment to study the various factors that affect the rates of chemical reactions. In part A, we used at first four different acids to react with the same substance (Mg).. then we used the acid that happened to react the fastest, HCl and placed three different metals in three different test tubes to see the reaction times. Overall, magnesium and HCl is the fastest reacting combination. The magnesium reacts with the a cid, producing visible bubbles of hydrogen gas. In Parts B and C we used temperature combination to determine which combination of reactants reacted the fastest. Using the Hydrochloric Acid- Sodium Thiosulfate Reaction, as the temperature rises, the reaction speeded up. This was similar to the Oxalic Acid- Potassium Permanganate Reaction System, except the reaction occurred more slowly in the 40 degree bath than the room temperature bath. In Part D, the catalyst lowered the activation energy, thus speeding up the reaction, but yet not using any of it up within the reaction. Lastly, in Part E, we now used different concentrations of Hydrochloric Acid and reacted them with Magnesium, finding the rate of reaction. The higher concentration reacted the fastest, the 6 M, while the other concentrations as they decreased, decreased the rate of the reaction.  The increase of the concentration of the reactant means more of a chemical present... More reactant particles will be moving together to allow more collisions to happen and so the reaction rate is i ncreased.

Conclusion: This experiment, through the individual procedures, allowed us to witness and time the reactions’ rates of different chemical combinations via different factors. We graphed some of these values and found a rate based on temperatures of how quickly these reactions may occur… and the other set of experiments we watched the sequence of reactions for different metal combinations… based on different chemicals or the different concentrations of one chemical (ex: HCl). From calculations, we now have a better understanding of how rates of reactions vary and lead to importance in many different chemical situations....