Factors Affecting Reaction Rates Pre Lab

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I need answers for this pre laboratory from page 277 to 278, thanks

23 Experiment Jo A. Beran/Trey Hernandez Factors Affecting Reaction Rates The reaction rate of zinc metal decreases (left to right) with decreasing concentration of hydrochloric acid. • To study the various factors that affect the rates of chemical reactions Objective The following techniques are used in the Experimental Procedure: Techniques 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 (e.g., per second)—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 a species as a function of time in a chemical reaction are color (expressed as absorbance, Figure 23.1), temperature, pH, gas evolution (see opening photo), odor, and conductivity. The parameter chosen for following the rate of a particular reaction depends on the nature of the reaction and the species of the reaction. We will investigate four of five factors that can be controlled to affect the rate of a chemical reaction. The first four factors listed below are systematically studied in this experiment: Introduction • Nature of the reactants • Temperature of the chemical system • Presence of a catalyst • Concentration of the reactants • Surface area of the reactants Some substances are naturally more reactive than others and therefore undergo rapid chemical changes. For example, the reaction of sodium metal and water is a very rapid, exothermic reaction (see Experiment 11, Part F), whereas the corrosion of iron is much slower. Plastics, reinforced with fibers such as carbon or glass, are now being substituted for iron and steel in specialized applications where corrosion has historically been a problem. 10 –9 second merzavka/iStockphoto Nanosecond: 1 Figure 23.1 The higher concentration of light-absorbing species, the more intense is the color of the solution. Species: any atom, molecule, or ion that may be a reactant or product of a chemical reaction Nature of the Reactants Sodium metal and water: the reaction releases H2(g) which ignites with the oxygen in the air to produce a yellow/blue flame, the yellow resulting from the presence of Na+ in the flame Experiment 23 271 Temperature of the Chemical System Internal energy: the energy contained within the molecules/ions when they collide As a rule of thumb, a 10°C rise in temperature doubles (increases by a factor of 2) the rate of a chemical reaction. The added heat not only increases the number of collisions1 between reactant molecules but also, and more importantly, 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. Presence of a Catalyst Figure 23.2 Reaction profiles of an uncatalyzed and a catalyzed reaction A catalyst increases the rate of a chemical reaction without undergoing any net chemical change. Some catalysts increase the rate of only one specific chemical reaction without affecting similar reactions. Other catalysts are more general and affect an entire set of similar reactions. Catalysts generally reroute the pathway of a chemical reaction so that this “alternate” path, although perhaps more circuitous, has a lower activation energy for reaction than the uncatalyzed reaction (Figure 23.2). Concentration of the Reactants An increase in the concentration of a reactant generally increases the reaction rate. See the opening photo. The larger concentration of reactant molecules increases the probability of an “effective” collision between reacting molecules for the formation of product. On occasion, such an increase may have no effect or may even decrease the reaction rate. A quantitative investigation on the effect of concentration changes on reaction rate is undertaken in Experiment 24 . Surface Area of the Reactants Generally speaking, the greater the exposed surface area of the reactant, the greater the reaction rate. For example, a large piece of coal burns very slowly, but coal dust burns rapidly, a consequence of which can lead to a disastrous coal mine explosion; solid potassium iodide reacts very slowly with solid lead nitrate, but when both are dissolved in solution, the formation of lead iodide is instantaneous. Experimental Procedure Procedure Overview: A series of qualitative experiments are conducted to determine how various factors affect the rate of a chemical reaction. Caution: A number of strong acids are used in the experiment. Handle with care; do not allow them to touch the skin or clothing. Perform the experiment with a partner. At each circled superscript 1–19 in the procedure, stop and record your observation on the Report Sheet. Discuss your observations with your lab partner and your instructor. Ask your instructor which parts of the Experimental Procedure you are to complete. Use a 250-mL beaker to prepare an ice water bath for Part B.3 and the hot water baths for Parts B.4 and C.3, 4. A. Nature of the Reactants 1. Different acids affect reaction rates. Half-fill a set of four labeled small test tubes (Figure 23.3) with 3 M H2SO4, 6 M HCl, 6 M CH3COOH, and 6 M H3PO4, respectively in a test tube rack. (Caution: Avoid skin contact with the acids.) Submerge a 1-cm strip of magnesium ribbon into each test tube. Compare the reaction rates and record your observations. 1 2. Different metals affect reaction rates. Half-fill a set of three labeled small test tubes (Figure 23.4) with 6 M HCl. Submerge 1-cm strips of zinc, magnesium, and copper separately into the test tubes. Compare the reaction rates of each metal in HCl and record your observations. 2 Match the relative reactivity of the metals with the photos in Figure 23.5. 3 A 10°C temperature rise only increases the collision frequency between reactant molecules by a factor of 1.02—nowhere near the factor of 2 that is normally experienced in a reaction rate. 1 272 Factors Affecting Reaction Rates M M M M Figure 23.4 Setup for the effect of metal type on reaction rate Test Tube Rack Jo A. Beran/Trey Hernandez Jo A. Beran/Trey Hernandez Figure 23.5 Zinc, copper, and magnesium react at different rates with 6 M HCl. Identify the metals in the photo according to their reactivity. 3 Disposal: Dispose of the reaction solutions in the Waste Inorganic Test Solutions container. Ask your instructor to determine if both Parts B and C are to be completed. You should perform the experiment with a partner; as one student combines the test solutions, the other notes the time. The oxidation–reduction reaction that occurs between hydrochloric acid and sodium thiosulfate, Na2S2O3, produces insoluble sulfur as a product. 2 HCl(aq) + Na2S2O3(aq) —› S(s) + SO2(g) + 2 NaCl(aq) + H2O(l) B. Temperature of the Reaction: Hydrochloric Acid–Sodium Thiosulfate Reaction System (23.1) The time required for the cloudiness of sulfur to appear is a measure of the reaction rate. Measure each volume of reactant with separate graduated pipets. 1. Prepare the solutions. Pipet 2 mL of 0.1 M Na2S2O3 into each of a set of three 150-mm, clean test tubes. Into a second set of three 150-mm test tubes, pipet 2 mL of 0.1 M HCl. 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 Na2S2O3–HCl pair test tubes in an ice water bath for Part B.3. and a third pair of Na2S2O3–HCl pair test tubes in a warm water bath (< 60°C) for Part B.4. Allow each pair of test tubes to establish thermal equilibrium (~5 minutes) before continuing to Parts B.3, and 4. Courtesy of VWR International, LLC Jo A. Beran/Trey Hernandez Figure 23.3 Setup for the effect of acid type on reaction rate Courtesy of Thermo Fisher Scientific M 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 and room temperature, using all certain digits plus one uncertain digit. Experiment 23 273 Data Analysis, F 3. Record the time for reaction at the lower temperature. From the ice bath, pour the HCl solution into the Na2S2O3 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 lapse for the reaction and the temperature of the bath, using all certain digits plus one uncertain digit. 4 4. Record the time for reaction at the higher temperature. From the warm water bath, pour the HCl solution into the Na2S2O3 solution and proceed as in Parts B.2 and B.3. Record the appropriate data. 6 Repeat any of the above reactions as deemed necessary. 5. Plot the data. Plot temperature (y-axis) versus time (x-axis) on one-half of a sheet of linear graph paper or by using appropriate software for the three data points. Have the instructor approve your graph. 7 Further interpret your data as suggested on the Report Sheet. Disposal: Dispose of the reaction solutions in the Waste Inorganic Test Solutions container. C. Temperature of the Reaction: Oxalic Acid–Potassium Permanganate Reaction System The reaction rate for the oxidation–reduction reaction between oxalic acid, H2C2O4, and potassium permanganate, KMnO4, is measured by recording the time elapsed for the (purple) color of the permanganate ion, MnO4 —, to disappear in the reaction: 5 H2C2O4(aq) + 2 KMnO4(aq) + 3 H2SO4(aq) —› 10 CO2(g) + 2 MnSO4(aq) + K2SO4(aq) + 8 H2O(l) (23.2) Measure the volume of each solution with separate clean graduated pipets. As one student pours the test solutions together, the other notes the time. Data Analysis, F 1. Prepare the solutions. Into a set of three, clean 150-mm test tubes, pipet 1 mL of 0.01 M KMnO4 (in 3 M H2SO4) and 4 mL of 3 M H2SO4. (Caution: KMnO4 is a strong oxidant and causes brown skin stains; H2SO4 is a severe skin irritant and is corrosive. Do not allow either chemical to make skin contact.) Into a second set of three clean 150-mm test tubes pipet 5 mL of 0.33 M H2C2O4. 2. Record the time for reaction at room temperature. Select a KMnO4—H2C2O4 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 certain digits plus one uncertain digit. 8 3. Record the time for reaction at the higher temperature. Place a second KMnO4–H2C2O4 pair of test tubes in a warm water (~40°C) bath until thermal equilibrium is established (~5 minutes). 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. Record the temperature of the bath. 9 4. Record the time for reaction at the highest temperature. Repeat Part C.3 but increase the temperature of the bath to ~60°C. Record the appropriate data.10 Repeat any of the preceding reactions as necessary. 5. Plot the data. Plot temperature (y-axis) versus time (x-axis) on one-half of a sheet of linear graph paper or by using apropriate software for the three data points. Have the instructor approve your graph.11 Disposal: Dispose of the reaction solutions in the Waste Inorganic Test Solutions container. 274 Factors Affecting Reaction Rates Hydrogen peroxide is relatively stable, but it readily decomposes in the presence of a catalyst. 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 1 or 2 crystals of MnO2 to the solution and observe. Note its instability.12 Ask your instructor for advice in completing both Parts E and F. 1. Prepare the reactants. Into a set of four clean, labeled test tubes, pipet 5 mL of 6 M HCl, 4 M HCl, 3 M HCl, and 1 M HCl, respectively (Figure 23.6).2 Determine the mass (± 0.001 g)—separately (for each solution)—of four 1-cm strips of polished (with steel wool or sand paper) magnesium. Calculate the number of moles of magnesium in each strip. 13 M M M E. Concentration of Reactants: Magnesium–Hydrochloric Acid System Data Analysis, A M Figure 23.6 Setup for the effect of acid concentration on reaction rate 2. Record the time for completion of the reaction. Add the first magnesium strip to the 6 M HCl solution. START TIME. Record the time for all traces of the magnesium strip to disappear. Repeat the experiment with the remaining three magnesium strips and the 4 M HCl, 3 M HCl, and 1 M HCl, solutions. 14 3. Plot the data. Plot mol HCl (y-axis) versus time in seconds (x-axis) for the four tests mol Mg on one-half of a sheet of linear graph paper or by using appropriate software. Have the instructor approve your graph. 15 Data Analysis, F Disposal: Dispose of the reaction solutions in the test tubes in the Waste Inorganic Test Solutions container. CLEANUP: Rinse the test tubes twice with tap water and twice with deionized water. Discard each rinse in the sink; flush the sink with water. A series of interrelated oxidation–reduction reactions occur between iodic acid, HIO3, and sulfurous acid, H2SO3, that ultimately lead to the formation of triiodide ion, I3–, and sulfuric acid, H2SO4, as the final products. 3 HIO3(aq) + 8 H2SO3(aq) —› H+(aq) + I3 –(aq) + 8 H2SO4(aq) + H2O(l) (23.3) F. Concentration of Reactants: Iodic Acid–Sulfurous Acid System The triodide ion, I3– ([I2•I]–), appears only after all of the sulfurous acid is consumed in the reaction. Once the I3– forms, its presence is detected by its reaction with starch, forming a deep-blue complex. I3 –(aq) + starch(aq) —› I3 – • starch(aq) (deep blue) (23.4) 1. Prepare the test solutions. Review the preparation of the test solutions in Table 23.1, page 276. Set up five, clean and labeled test tubes (Figure 23.7) in a test tube rack. Measure the volumes of the 0.01 M HIO3, starch, and water with dropping (or Beral) pipets.3 Calibrate the HIO3 dropping pipet to determine the volume (mL) 2 Remember to properly rinse the pipet with the appropriate solution before dispensing it into the test tube. 3 Be careful! Do not intermix the dropping pipets between solutions. This error in technique causes a significant error in the data. Figure 23.7 Setup for changes in HIO3 concentration on reaction rate Experiment 23 275 Table 23.1 Reactant Concentration and Reaction Rate Solution in Test Tube Add to Test Tube Test Tube 0.01 M HIO3 Starch H2O 0.01 M H2SO3 1 2 3 4 5 3 drops 6 drops 12 drops 15 drops 20 drops 1 drop 1 drop 1 drop 1 drop 1 drop 17 drops 14 drops 8 drops 5 drops 0 drops 1.0 mL 1.0 mL 1.0 mL 1.0 mL 1.0 mL per drop. 16 Calibrate a second dropping (or Beral) pipet with water to determine the number of milliliters per drop.17 Calibrate a third dropping (or Beral) pipet for the 0.01 M H2SO3 solution that delivers 1 mL; mark the level on the pipet so that quick delivery of 1 mL of the H2SO3 solution to each test tube can be made. Alternatively, use a calibrated 1-mL Beral pipet. 2. Record the time for the reaction. Place a sheet of white paper beside the test tube (Figure 23.8). 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 (seconds) for the deep-blue I3–•starch complex to appear.4 Figure 23.8 Viewing the reaction rate in a test tube 3. Complete remaining reactions. Repeat Part F.2. for the remaining reaction mixtures in Table 23.1. Repeat any of the trials as necessary. 18 4. Plot the data. On one-half of a sheet of linear graph paper or by using appropriate software, plot for each solution the initial concentration of iodic acid,5 [HIO3]0 (y-axis), versus the time in seconds (x-axis) for the reaction. 19 Data Analysis, F Disposal: Dispose of all test solutions in the Waste Inorganic Test Solutions container. CLEANUP: Rinse the test tubes twice with tap water and discard each into the Waste Inorganic Test Solutions container. Two final rinses with deionized water can be discarded in the sink. The Next Step (1) The dissolution of dissolved gases such as CO2(aq) in carbonated beverages, changes significantly with temperature changes. Study the kinetics of the dissolution of dissolved gases such as CO2(aq) or O2(g) using such things as Mentos candy, salt, rust, and so on. The study may be qualitative or quantitative. For the dissolution of O2(g), refer to Experiment 31 in this manual. (2) Corrosion of iron in deionized water, tap water, boiled deionized/tap water, salt water (varying concentrations), and so on all affect the economy. 4 Be ready! The appearance of the deep-blue solution is sudden. Remember that in calculating [HIO3]0, the total volume of the solution is the sum of the volumes of the two solutions expressed in liters. 5 276 Factors Affecting Reaction Rates Experiment 23 Prelaboratory Assignment Factors Affecting Reaction Rates Date __________ Lab Sec. ______ Name ____________________________________________ Desk No. __________ Adam Hart-Davis/Science Source 1. Identify the major factor affecting reaction rates that accounts for the following observations: a. Tadpoles grow more rapidly near the cooling water discharged from a power plant. b. Enzymes facilitate certain biochemical reactions but are not consumed. c. Hydrogen peroxide antiseptic rapidly decomposes when applied to an open wound. 2. Chlorofluorocarbons photodissociate to produce chlorine atoms, Cl•, which have been implicated in decreasing the concentration of ozone, O3, in the stratosphere. The decomposition of the ozone follows a reaction sequence of O3 + Cl• l ClO• + O2 ClO• + O l Cl• + O2 What role (factor affecting reaction rates) do chlorine atoms have in increasing the depletion rate of ozone? 3. Assuming that the rate of a chemical reaction doubles for every 10°C temperature increase, by what factor would a chemical reaction increase if the temperature were increased from –5°C (a cold winter morning) to 25°C (room temperature)? 4. Experimental Procedure, Part B a. Identify the visual evidence used for timing the reaction. b. A data plot is used to predict reaction rates at other conditions. What are the coordinates of the data plot? Experiment 23 277 5. Experimental Procedure, Part E.3 a. An 18-mg strip of magnesium metal reacts in 5.0 mL of 3.0 M HCl over a given time period. Evaluate the mol HCl mol Mg ratio for the reaction. b. What are the correct labelings of the axes for the data plot? 6. Experimental Procedure, Part F. A 1.0-mL volume of 0.010 M H2SO3 is added to a mixture of 6 drops of 0.010 M HIO3, 14 drops of deionized water, and 1 drop of starch solution. A color change in the reaction mixture occurred after 56 seconds. a. Assuming 20 drops per milliliter for all solutions, determine the initial molar concentration of HIO3 after the mixing mol HlO3 but before any reaction occurs (at time = 0). Hint: Units are . total volume (L) b. The rate of the reaction is measured by the disappearance of HIO3. For the reaction mixture in this question, what is mol HlO3/L the reaction rate? Express the reaction rate in units of to the correct number of significant figures. sec 7. The reactions in the Experimental Procedure, Parts C, E, and F, are timed. Identify the visual signal to stop timing in each reaction. a. Part C. b. Part E. c. Part F. 278 Factors Affecting Reaction Rates Experiment 23 Report Sheet Factors Affecting Reaction Rates Date __________ Lab Sec. ______ Name ____________________________________________ Desk No. __________ A. Nature of the Reactants 1. 1 List the acids in order of decreasing reaction rate with magnesium: _________, _________, __________,_________ 2. 2 List the metals in order of decreasing reaction rate with 6 M HCl: _____________, _____________,_____________ 3. 3 Identify the metals reacting in Figure 23.5 (from left to right). ______________, ______________,______________ B. Temperature of the Reaction: Hydrochloric Acid–Sodium Thiosulfate Reaction System 1. Time for Sulfur to Appear 2. Temperature of the Reaction 4 __________ seconds __________ °C 5 __________ seconds __________ °C 6 __________ seconds __________ °C 7 Plot temperature (y-axis) versus time (x-axis) for the three trials. Instructor’s approval of graph: _______________ 3. From the plotted data, interpret the effect of temperature on reaction rate. 4. From your graph, estimate the temperature at which the appearance of sulfur should occur in 20 seconds. Assume no changes in concentration. C. Temperature of the Reaction: Oxalic Acid–Potassium Permanganate Reaction System 1. Time for Permanganate Ion to Disappear 2. Temperature of the Reaction 8 __________ seconds __________ °C 9 __________ seconds __________ °C 10 __________ seconds __________ °C 11 Plot temperature (y-axis) versus time (x-axis) for the three trials. Instructor’s approval of graph: _______________ 3. From your plotted data, interpret the affect of temperature on reaction rate. 4. From your graph, estimate the time for the disappearance of the purple permanganate ion at 55°C. Assume no changes in concentration. Experiment 23 279 D. Presence of a Catalyst 1. 12 What effect does the MnO2 catalyst have on the rate of evolution of O2 gas? 2. Write a balanced equation for the decomposition of H2O2. E. Concentration of Reactants: Magnesium–Hydrochloric Acid System Concentration of HCl Volume of HCl mol HCl mass of Mg 6M ____________ ____________ ____________ ____________ ____________ ____________ 4M ____________ ____________ ____________ ____________ ____________ ____________ 3M ____________ ____________ ____________ ____________ ____________ ____________ 1M ____________ ____________ ____________ ____________ ____________ ____________ 13 mol HCl mol Mg mol Mg 14 Time (sec) Plot mol HCl (y-axis) versus time (x-axis). Instructor’s approval of graph: __________________________________ mol Mg 2. From your graph, predict the time, in seconds, for 5 mg of Mg to react in 5 mL of 2.0 M HCl. 1. 15 F. Concentration of Reactants: Iodic Acid–Sulfurous Acid System Molar concentration of HIO3 __________ Molar concentration of H2SO3 __________ 16 What is the volume (mL) per drop of the HIO3 solution? _____ 17 What is the volume (mL) per drop of the water? _____ Test Tube mL H2SO3 [HIO3]06 (diluted) Time (sec) __________ 1.0 _______________ __________ __________ __________ 1.0 _______________ __________ __________ __________ __________ 1.0 _______________ __________ __________ __________ __________ __________ 1.0 _______________ __________ __________ __________ __________ __________ 1.0 _______________ __________ mL HIO3 Drops of H2O mL H2O __________ __________ __________ 2 __________ __________ 3 __________ 4 5 18 Drops of HIO3 1 1. 19 Plot [HIO3]0 (y-axis) versus time (x-axis). Instructor’s approval of graph: ___________________________________ 2. How does a change in the molar concentration of HIO3 affect the time required for the appearance of the deep-blue I3–•starch complex? 3. Estimate the time, in seconds, for the deep-blue I3– •starch complex to form when 10 drops of 0.01 M HIO3 are used for the reaction. Assume all other conditions remain constant. 6 See footnote 5. 280 Factors Affecting Reaction Rates
23 Experiment Jo A. Beran/Trey Hernandez Factors Affecting Reaction Rates The reaction rate of zinc metal decreases (left to right) with decreasing concentration of hydrochloric acid. • To study the various factors that affect the rates of chemical reactions Objective The following techniques are used in the Experimental Procedure: Techniques 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 (e.g., per second)—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 a species as a function of time in a chemical reaction are color (expressed as absorbance, Figure 23.1), temperature, pH, gas evolution (see opening photo), odor, and conductivity. The parameter chosen for following the rate of a particular reaction depends on the nature of the reaction and the species of the reaction. We will investigate four of five factors that can be controlled to affect the rate of a chemical reaction. The first four factors listed below are systematically studied in this experiment: Introduction • Nature of the reactants • Temperature of the chemical system • Presence of a catalyst • Concentration of the reactants • Surface area of the reactants Some substances are naturally more reactive than others and therefore undergo rapid chemical changes. For example, the reaction of sodium metal and water is a very rapid, exothermic reaction (see Experiment 11, Part F), whereas the corrosion of iron is much slower. Plastics, reinforced with fibers such as carbon or glass, are now being substituted for iron and steel in specialized applications where corrosion has historically been a problem. 10 –9 second merzavka/iStockphoto Nanosecond: 1 Figure 23.1 The higher concentration of light-absorbing species, the more intense is the color of the solution. Species: any atom, molecule, or ion that may be a reactant or product of a chemical reaction Nature of the Reactants Sodium metal and water: the reaction releases H2(g) which ignites with the oxygen in the air to produce a yellow/blue flame, the yellow resulting from the presence of Na+ in the flame Experiment 23 271 Temperature of the Chemical System Internal energy: the energy contained within the molecules/ions when they collide As a rule of thumb, a 10°C rise in temperature doubles (increases by a factor of 2) the rate of a chemical reaction. The added heat not only increases the number of collisions1 between reactant molecules but also, and more importantly, 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. Presence of a Catalyst Figure 23.2 Reaction profiles of an uncatalyzed and a catalyzed reaction A catalyst increases the rate of a chemical reaction without undergoing any net chemical change. Some catalysts increase the rate of only one specific chemical reaction without affecting similar reactions. Other catalysts are more general and affect an entire set of similar reactions. Catalysts generally reroute the pathway of a chemical reaction so that this “alternate” path, although perhaps more circuitous, has a lower activation energy for reaction than the uncatalyzed reaction (Figure 23.2). Concentration of the Reactants An increase in the concentration of a reactant generally increases the reaction rate. See the opening photo. The larger concentration of reactant molecules increases the probability of an “effective” collision between reacting molecules for the formation of product. On occasion, such an increase may have no effect or may even decrease the reaction rate. A quantitative investigation on the effect of concentration changes on reaction rate is undertaken in Experiment 24 . Surface Area of the Reactants Generally speaking, the greater the exposed surface area of the reactant, the greater the reaction rate. For example, a large piece of coal burns very slowly, but coal dust burns rapidly, a consequence of which can lead to a disastrous coal mine explosion; solid potassium iodide reacts very slowly with solid lead nitrate, but when both are dissolved in solution, the formation of lead iodide is instantaneous. Experimental Procedure Procedure Overview: A series of qualitative experiments are conducted to determine how various factors affect the rate of a chemical reaction. Caution: A number of strong acids are used in the experiment. Handle with care; do not allow them to touch the skin or clothing. Perform the experiment with a partner. At each circled superscript 1–19 in the procedure, stop and record your observation on the Report Sheet. Discuss your observations with your lab partner and your instructor. Ask your instructor which parts of the Experimental Procedure you are to complete. Use a 250-mL beaker to prepare an ice water bath for Part B.3 and the hot water baths for Parts B.4 and C.3, 4. A. Nature of the Reactants 1. Different acids affect reaction rates. Half-fill a set of four labeled small test tubes (Figure 23.3) with 3 M H2SO4, 6 M HCl, 6 M CH3COOH, and 6 M H3PO4, respectively in a test tube rack. (Caution: Avoid skin contact with the acids.) Submerge a 1-cm strip of magnesium ribbon into each test tube. Compare the reaction rates and record your observations. 1 2. Different metals affect reaction rates. Half-fill a set of three labeled small test tubes (Figure 23.4) with 6 M HCl. Submerge 1-cm strips of zinc, magnesium, and copper separately into the test tubes. Compare the reaction rates of each metal in HCl and record your observations. 2 Match the relative reactivity of the metals with the photos in Figure 23.5. 3 A 10°C temperature rise only increases the collision frequency between reactant molecules by a factor of 1.02—nowhere near the factor of 2 that is normally experienced in a reaction rate. 1 272 Factors Affecting Reaction Rates M M M M Figure 23.4 Setup for the effect of metal type on reaction rate Test Tube Rack Jo A. Beran/Trey Hernandez Jo A. Beran/Trey Hernandez Figure 23.5 Zinc, copper, and magnesium react at different rates with 6 M HCl. Identify the metals in the photo according to their reactivity. 3 Disposal: Dispose of the reaction solutions in the Waste Inorganic Test Solutions container. Ask your instructor to determine if both Parts B and C are to be completed. You should perform the experiment with a partner; as one student combines the test solutions, the other notes the time. The oxidation–reduction reaction that occurs between hydrochloric acid and sodium thiosulfate, Na2S2O3, produces insoluble sulfur as a product. 2 HCl(aq) + Na2S2O3(aq) —› S(s) + SO2(g) + 2 NaCl(aq) + H2O(l) B. Temperature of the Reaction: Hydrochloric Acid–Sodium Thiosulfate Reaction System (23.1) The time required for the cloudiness of sulfur to appear is a measure of the reaction rate. Measure each volume of reactant with separate graduated pipets. 1. Prepare the solutions. Pipet 2 mL of 0.1 M Na2S2O3 into each of a set of three 150-mm, clean test tubes. Into a second set of three 150-mm test tubes, pipet 2 mL of 0.1 M HCl. 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 Na2S2O3–HCl pair test tubes in an ice water bath for Part B.3. and a third pair of Na2S2O3–HCl pair test tubes in a warm water bath (< 60°C) for Part B.4. Allow each pair of test tubes to establish thermal equilibrium (~5 minutes) before continuing to Parts B.3, and 4. Courtesy of VWR International, LLC Jo A. Beran/Trey Hernandez Figure 23.3 Setup for the effect of acid type on reaction rate Courtesy of Thermo Fisher Scientific M 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 and room temperature, using all certain digits plus one uncertain digit. Experiment 23 273 Data Analysis, F 3. Record the time for reaction at the lower temperature. From the ice bath, pour the HCl solution into the Na2S2O3 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 lapse for the reaction and the temperature of the bath, using all certain digits plus one uncertain digit. 4 4. Record the time for reaction at the higher temperature. From the warm water bath, pour the HCl solution into the Na2S2O3 solution and proceed as in Parts B.2 and B.3. Record the appropriate data. 6 Repeat any of the above reactions as deemed necessary. 5. Plot the data. Plot temperature (y-axis) versus time (x-axis) on one-half of a sheet of linear graph paper or by using appropriate software for the three data points. Have the instructor approve your graph. 7 Further interpret your data as suggested on the Report Sheet. Disposal: Dispose of the reaction solutions in the Waste Inorganic Test Solutions container. C. Temperature of the Reaction: Oxalic Acid–Potassium Permanganate Reaction System The reaction rate for the oxidation–reduction reaction between oxalic acid, H2C2O4, and potassium permanganate, KMnO4, is measured by recording the time elapsed for the (purple) color of the permanganate ion, MnO4 —, to disappear in the reaction: 5 H2C2O4(aq) + 2 KMnO4(aq) + 3 H2SO4(aq) —› 10 CO2(g) + 2 MnSO4(aq) + K2SO4(aq) + 8 H2O(l) (23.2) Measure the volume of each solution with separate clean graduated pipets. As one student pours the test solutions together, the other notes the time. Data Analysis, F 1. Prepare the solutions. Into a set of three, clean 150-mm test tubes, pipet 1 mL of 0.01 M KMnO4 (in 3 M H2SO4) and 4 mL of 3 M H2SO4. (Caution: KMnO4 is a strong oxidant and causes brown skin stains; H2SO4 is a severe skin irritant and is corrosive. Do not allow either chemical to make skin contact.) Into a second set of three clean 150-mm test tubes pipet 5 mL of 0.33 M H2C2O4. 2. Record the time for reaction at room temperature. Select a KMnO4—H2C2O4 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 certain digits plus one uncertain digit. 8 3. Record the time for reaction at the higher temperature. Place a second KMnO4–H2C2O4 pair of test tubes in a warm water (~40°C) bath until thermal equilibrium is established (~5 minutes). 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. Record the temperature of the bath. 9 4. Record the time for reaction at the highest temperature. Repeat Part C.3 but increase the temperature of the bath to ~60°C. Record the appropriate data.10 Repeat any of the preceding reactions as necessary. 5. Plot the data. Plot temperature (y-axis) versus time (x-axis) on one-half of a sheet of linear graph paper or by using apropriate software for the three data points. Have the instructor approve your graph.11 Disposal: Dispose of the reaction solutions in the Waste Inorganic Test Solutions container. 274 Factors Affecting Reaction Rates Hydrogen peroxide is relatively stable, but it readily decomposes in the presence of a catalyst. 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 1 or 2 crystals of MnO2 to the solution and observe. Note its instability.12 Ask your instructor for advice in completing both Parts E and F. 1. Prepare the reactants. Into a set of four clean, labeled test tubes, pipet 5 mL of 6 M HCl, 4 M HCl, 3 M HCl, and 1 M HCl, respectively (Figure 23.6).2 Determine the mass (± 0.001 g)—separately (for each solution)—of four 1-cm strips of polished (with steel wool or sand paper) magnesium. Calculate the number of moles of magnesium in each strip. 13 M M M E. Concentration of Reactants: Magnesium–Hydrochloric Acid System Data Analysis, A M Figure 23.6 Setup for the effect of acid concentration on reaction rate 2. Record the time for completion of the reaction. Add the first magnesium strip to the 6 M HCl solution. START TIME. Record the time for all traces of the magnesium strip to disappear. Repeat the experiment with the remaining three magnesium strips and the 4 M HCl, 3 M HCl, and 1 M HCl, solutions. 14 3. Plot the data. Plot mol HCl (y-axis) versus time in seconds (x-axis) for the four tests mol Mg on one-half of a sheet of linear graph paper or by using appropriate software. Have the instructor approve your graph. 15 Data Analysis, F Disposal: Dispose of the reaction solutions in the test tubes in the Waste Inorganic Test Solutions container. CLEANUP: Rinse the test tubes twice with tap water and twice with deionized water. Discard each rinse in the sink; flush the sink with water. A series of interrelated oxidation–reduction reactions occur between iodic acid, HIO3, and sulfurous acid, H2SO3, that ultimately lead to the formation of triiodide ion, I3–, and sulfuric acid, H2SO4, as the final products. 3 HIO3(aq) + 8 H2SO3(aq) —› H+(aq) + I3 –(aq) + 8 H2SO4(aq) + H2O(l) (23.3) F. Concentration of Reactants: Iodic Acid–Sulfurous Acid System The triodide ion, I3– ([I2•I]–), appears only after all of the sulfurous acid is consumed in the reaction. Once the I3– forms, its presence is detected by its reaction with starch, forming a deep-blue complex. I3 –(aq) + starch(aq) —› I3 – • starch(aq) (deep blue) (23.4) 1. Prepare the test solutions. Review the preparation of the test solutions in Table 23.1, page 276. Set up five, clean and labeled test tubes (Figure 23.7) in a test tube rack. Measure the volumes of the 0.01 M HIO3, starch, and water with dropping (or Beral) pipets.3 Calibrate the HIO3 dropping pipet to determine the volume (mL) 2 Remember to properly rinse the pipet with the appropriate solution before dispensing it into the test tube. 3 Be careful! Do not intermix the dropping pipets between solutions. This error in technique causes a significant error in the data. Figure 23.7 Setup for changes in HIO3 concentration on reaction rate Experiment 23 275 Table 23.1 Reactant Concentration and Reaction Rate Solution in Test Tube Add to Test Tube Test Tube 0.01 M HIO3 Starch H2O 0.01 M H2SO3 1 2 3 4 5 3 drops 6 drops 12 drops 15 drops 20 drops 1 drop 1 drop 1 drop 1 drop 1 drop 17 drops 14 drops 8 drops 5 drops 0 drops 1.0 mL 1.0 mL 1.0 mL 1.0 mL 1.0 mL per drop. 16 Calibrate a second dropping (or Beral) pipet with water to determine the number of milliliters per drop.17 Calibrate a third dropping (or Beral) pipet for the 0.01 M H2SO3 solution that delivers 1 mL; mark the level on the pipet so that quick delivery of 1 mL of the H2SO3 solution to each test tube can be made. Alternatively, use a calibrated 1-mL Beral pipet. 2. Record the time for the reaction. Place a sheet of white paper beside the test tube (Figure 23.8). 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 (seconds) for the deep-blue I3–•starch complex to appear.4 Figure 23.8 Viewing the reaction rate in a test tube 3. Complete remaining reactions. Repeat Part F.2. for the remaining reaction mixtures in Table 23.1. Repeat any of the trials as necessary. 18 4. Plot the data. On one-half of a sheet of linear graph paper or by using appropriate software, plot for each solution the initial concentration of iodic acid,5 [HIO3]0 (y-axis), versus the time in seconds (x-axis) for the reaction. 19 Data Analysis, F Disposal: Dispose of all test solutions in the Waste Inorganic Test Solutions container. CLEANUP: Rinse the test tubes twice with tap water and discard each into the Waste Inorganic Test Solutions container. Two final rinses with deionized water can be discarded in the sink. The Next Step (1) The dissolution of dissolved gases such as CO2(aq) in carbonated beverages, changes significantly with temperature changes. Study the kinetics of the dissolution of dissolved gases such as CO2(aq) or O2(g) using such things as Mentos candy, salt, rust, and so on. The study may be qualitative or quantitative. For the dissolution of O2(g), refer to Experiment 31 in this manual. (2) Corrosion of iron in deionized water, tap water, boiled deionized/tap water, salt water (varying concentrations), and so on all affect the economy. 4 Be ready! The appearance of the deep-blue solution is sudden. Remember that in calculating [HIO3]0, the total volume of the solution is the sum of the volumes of the two solutions expressed in liters. 5 276 Factors Affecting Reaction Rates Experiment 23 Prelaboratory Assignment Factors Affecting Reaction Rates Date __________ Lab Sec. ______ Name ____________________________________________ Desk No. __________ Adam Hart-Davis/Science Source 1. Identify the major factor affecting reaction rates that accounts for the following observations: a. Tadpoles grow more rapidly near the cooling water discharged from a power plant. b. Enzymes facilitate certain biochemical reactions but are not consumed. c. Hydrogen peroxide antiseptic rapidly decomposes when applied to an open wound. 2. Chlorofluorocarbons photodissociate to produce chlorine atoms, Cl•, which have been implicated in decreasing the concentration of ozone, O3, in the stratosphere. The decomposition of the ozone follows a reaction sequence of O3 + Cl• l ClO• + O2 ClO• + O l Cl• + O2 What role (factor affecting reaction rates) do chlorine atoms have in increasing the depletion rate of ozone? 3. Assuming that the rate of a chemical reaction doubles for every 10°C temperature increase, by what factor would a chemical reaction increase if the temperature were increased from –5°C (a cold winter morning) to 25°C (room temperature)? 4. Experimental Procedure, Part B a. Identify the visual evidence used for timing the reaction. b. A data plot is used to predict reaction rates at other conditions. What are the coordinates of the data plot? Experiment 23 277 5. Experimental Procedure, Part E.3 a. An 18-mg strip of magnesium metal reacts in 5.0 mL of 3.0 M HCl over a given time period. Evaluate the mol HCl mol Mg ratio for the reaction. b. What are the correct labelings of the axes for the data plot? 6. Experimental Procedure, Part F. A 1.0-mL volume of 0.010 M H2SO3 is added to a mixture of 6 drops of 0.010 M HIO3, 14 drops of deionized water, and 1 drop of starch solution. A color change in the reaction mixture occurred after 56 seconds. a. Assuming 20 drops per milliliter for all solutions, determine the initial molar concentration of HIO3 after the mixing mol HlO3 but before any reaction occurs (at time = 0). Hint: Units are . total volume (L) b. The rate of the reaction is measured by the disappearance of HIO3. For the reaction mixture in this question, what is mol HlO3/L the reaction rate? Express the reaction rate in units of to the correct number of significant figures. sec 7. The reactions in the Experimental Procedure, Parts C, E, and F, are timed. Identify the visual signal to stop timing in each reaction. a. Part C. b. Part E. c. Part F. 278 Factors Affecting Reaction Rates Experiment 23 Report Sheet Factors Affecting Reaction Rates Date __________ Lab Sec. ______ Name ____________________________________________ Desk No. __________ A. Nature of the Reactants 1. 1 List the acids in order of decreasing reaction rate with magnesium: _________, _________, __________,_________ 2. 2 List the metals in order of decreasing reaction rate with 6 M HCl: _____________, _____________,_____________ 3. 3 Identify the metals reacting in Figure 23.5 (from left to right). ______________, ______________,______________ B. Temperature of the Reaction: Hydrochloric Acid–Sodium Thiosulfate Reaction System 1. Time for Sulfur to Appear 2. Temperature of the Reaction 4 __________ seconds __________ °C 5 __________ seconds __________ °C 6 __________ seconds __________ °C 7 Plot temperature (y-axis) versus time (x-axis) for the three trials. Instructor’s approval of graph: _______________ 3. From the plotted data, interpret the effect of temperature on reaction rate. 4. From your graph, estimate the temperature at which the appearance of sulfur should occur in 20 seconds. Assume no changes in concentration. C. Temperature of the Reaction: Oxalic Acid–Potassium Permanganate Reaction System 1. Time for Permanganate Ion to Disappear 2. Temperature of the Reaction 8 __________ seconds __________ °C 9 __________ seconds __________ °C 10 __________ seconds __________ °C 11 Plot temperature (y-axis) versus time (x-axis) for the three trials. Instructor’s approval of graph: _______________ 3. From your plotted data, interpret the affect of temperature on reaction rate. 4. From your graph, estimate the time for the disappearance of the purple permanganate ion at 55°C. Assume no changes in concentration. Experiment 23 279 D. Presence of a Catalyst 1. 12 What effect does the MnO2 catalyst have on the rate of evolution of O2 gas? 2. Write a balanced equation for the decomposition of H2O2. E. Concentration of Reactants: Magnesium–Hydrochloric Acid System Concentration of HCl Volume of HCl mol HCl mass of Mg 6M ____________ ____________ ____________ ____________ ____________ ____________ 4M ____________ ____________ ____________ ____________ ____________ ____________ 3M ____________ ____________ ____________ ____________ ____________ ____________ 1M ____________ ____________ ____________ ____________ ____________ ____________ 13 mol HCl mol Mg mol Mg 14 Time (sec) Plot mol HCl (y-axis) versus time (x-axis). Instructor’s approval of graph: __________________________________ mol Mg 2. From your graph, predict the time, in seconds, for 5 mg of Mg to react in 5 mL of 2.0 M HCl. 1. 15 F. Concentration of Reactants: Iodic Acid–Sulfurous Acid System Molar concentration of HIO3 __________ Molar concentration of H2SO3 __________ 16 What is the volume (mL) per drop of the HIO3 solution? _____ 17 What is the volume (mL) per drop of the water? _____ Test Tube mL H2SO3 [HIO3]06 (diluted) Time (sec) __________ 1.0 _______________ __________ __________ __________ 1.0 _______________ __________ __________ __________ __________ 1.0 _______________ __________ __________ __________ __________ __________ 1.0 _______________ __________ __________ __________ __________ __________ 1.0 _______________ __________ mL HIO3 Drops of H2O mL H2O __________ __________ __________ 2 __________ __________ 3 __________ 4 5 18 Drops of HIO3 1 1. 19 Plot [HIO3]0 (y-axis) versus time (x-axis). Instructor’s approval of graph: ___________________________________ 2. How does a change in the molar concentration of HIO3 affect the time required for the appearance of the deep-blue I3–•starch complex? 3. Estimate the time, in seconds, for the deep-blue I3– •starch complex to form when 10 drops of 0.01 M HIO3 are used for the reaction. Assume all other conditions remain constant. 6 See footnote 5. 280 Factors Affecting Reaction Rates

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