Kinetics: Factors that Affect Rates of Chemical Reactions

Transcription

1 Objective- Study several factors that influence chemical reaction rates, including: 1. Concentration 2. The nature of the chemical reactants 3. Area in a heterogeneous reaction 4. The temperature of the reacting system 5. The presence of catalysts Background Effects of concentration In order to react, molecules or ions must encounter one another. These encounters (or collisions) are often violent enough to produce a rearrangement of the atoms and the formation of new products. The more molecules there are per unit volume, the more likely they are to collide with one another, so increasing the concentrations of the reactants generally increases the rate of reaction. However, the dependence of reaction rate on the concentrations of the reactants is intimately linked to the reaction pathway; in some instances, an increase in the concentration of a reactant has no effect (or even a slowing effect) on the rate of reaction. Experiment 4 we will determine quantitatively the dependence of reaction rate on the concentrations of the reactants. Influence of the nature of the reactants The alkali metals all react with water. The net reaction is the same for each metal: M+ H 2 O M + (aq) + OH - (aq) + 1/2 H 2 (g) (1) where M represents Li, Na, K, Rb, or Cs. The rate of reaction increases noticeably in the sequence lithium to cesium. Lithium reacts briskly, but cesium reacts violently, producing enough heat to ignite the hydrogen produced. The metals that react most rapidly have the lowest melting points and the smallest ionization energies. We will examine the reaction of alkali earth metal with water. Effect of surface area in a heterogeneous reaction In a reaction between a solid and a gas (or liquid), the reactants can get together only at the surface of the solid. Increasing the surface area of the solid, therefore, one should observe an increase in the number of atoms available to react. Effect of temperature Increasing the temperature of a reaction increases the kinetic energy of the reacting molecules. They move about faster, and thus collide with greater force, producing the bond-length and bond-angle distortions as well as bond-rupture characteristic of reacting molecules. So a molecule at a higher temperature has higher energy and a higher probability of reacting.

2 Effect of catalysts Factors that Influence Reaction Rates. A catalyst increases the rate of a chemical reaction and can be recovered from the reaction mixture unaltered. Catalysts generally function by interacting with one or more reactants to form an intermediate that is more reactive. In effect, catalysts change the reaction pathway by lowering the activation energy that is required for the reaction to take place. Pre-Lab and Experimental Procedure Special Supplies: Mortar and pestle; apparatus for collecting gases shown in Figure below. Chemicals: 3% (~0.9 M) H 2 O 2 ; solution A (containing in the same solution 0.2 g/l starch, 0.5 M CH 3 COOH, 0.05 M NaCH 3 COO, 0.30 M KI, and M Na 2 S 2 O 3 ); Mg and Ca metal turnings; 0.1% phenolphthalein indicator; marble chips {CaCO 3 }; 1 M HCl; steel wool; iron nails (approx. 4 cm long); methanol; 50:50 (v:v) 2-propanol water; 1 M CuCl 2 ; 3 M CuCl 2 ; 3 M FeCl 3 ; KMnO 4 (s); M KMnO 4 ; 0.5 M H 2 C 2 O 4 (oxalic acid); 6 M H 2 SO 4. Part 1. Effects of Concentration on Reaction Rates {Anwer all questions in your notebook and/or report form} (a) Place 5 ml of solution A in a 50-mL beaker. From a small graduate cylinder add 5 ml of 3% H 2 O 2. Mix the solution once, and time the number of seconds required for the solution to turn blue. Record the reaction time, t. Pour reaction mixture down the sink and rinse beaker and graduate cylinder with deionized water. (b) Repeat the procedure placing 5 ml of solution A in 50 ml beaker. In the graduated cylinder mix 4 ml of water, with 1 ml of 3% H 2 O 2. Add the contents of the graduated cylinder to the beaker, mix and record the reaction time. Is it longer? Which reactant concentrations have been changed? What is the relationship between the reaction time and the initial concentration of H 2 O 2? Part 2. The Influence of the Nature of the Reactants Obtain three or four small pieces of magnesium metal turnings and about the same amount of calcium metal turnings. Add enough water to each of two 15 mm x 150 mm test tubes to fill them about one-third full. Place the magnesium turnings in one test tube and the calcium turnings in the other. Both of these metals are in the same chemical family in the periodic table. Which metal is more reactive with water? To determine what gas is produced, invert a clean dry test tube over the mouth of the tube in which the metal is reacting more vigorously for about s. Remove but keep the tube inverted and bring a burning splint or match toward the mouth of the tube. Is the gas flammable? What is the chemical composition of the gas? Look closely at the slower-reacting metal as you agitate the test tube. Is any gas being produced (small gas bubbles may be observed)? Add a drop of phenolphthalein indicator solution to each test tube containing the metals. What other product is formed besides a gas (Hint: check equation 1)? Write chemical equations describing the reactions of the metals with water. Is there a correlation of the reaction rates with the ionization potentials of these metals (is it easier to remove electrons from magnesium or calcium)? Look up the first and second ionization energies of these metals and record these values in your notebook/report form. -2-

3 Part 3. The Influence of Surface Area in a Heterogeneous Reaction Factors that Influence Reaction Rates. (a) Set up the apparatus shown in the Figure below, using a 125-mL Erlenmeyer flask, a 13 x 100 mm test tube to collect the gas, and a large beaker or plastic basin. Fill the test tube with tap water and, holding your thumb or finger over the mouth, invert it in the beaker of water. When the mouth of the inverted test tube is below the surface of the water you can let go the test tube should remain filled. Insert the end of the bent gas delivery tube into the mouth of the inverted test tube as shown. Place 1-4 pieces of coarse marble chips (CaCO 3 ) in the 125-mL Erlenmeyer flask. Add 25 ml of water; then add 25 ml of 1 M HCl. Immediately insert the one-hole stopper in the flask and record the time required to fill 1/3 of the test tube with gas. Multiply this time by a factor of three and record it in your notebook/report form. What gas is produced? Rinse the Erlenmeyer flask with tap water and place the used marble chips in the waste receptacle provided. Do not discard the marble chips in the sink. Obtain another 3 g of marble chips and, using a mortar and pestle, pulverize them to the size of grains of sugar. Using a rolled-up sheet of paper as a funnel, transfer the granular material to the 125-mL Erlenmeyer flask. Fill a test tube with water and invert it over the delivery tube as before. Add 25 ml of water to the flask, followed by 25 ml of 1 M HCl, and stopper immediately. Record the time required to fill the test tube with gas. Save this apparatus for use in part 4. Is there a correlation between the surface area of the marble and the reaction rate? Write an equation describing the chemical reaction between CaCO 3 and HCl. Place a small iron nail (cleaned so that it is free of rust) in a 20-mm-diameter test tube. In a second 20-mm test tube, place a 0.6-g ball of steel wool, using a glass rod to push it to the bottom of the test tube. Add 10 ml of 1 M CuCl 2 to each test tube and watch them closely. Observe and record any color changes and note in which tube they occur more rapidly. Is a new solid sub-stance forming? What do you think it is? Agitate each test tube and feel the lower portion of the tubes. Is one tube warmer than the other? What happened to the steel wool? Would the same thing eventually happen to the nail? Write an equation describing the chemical reaction that takes place in your notebook/report form. Part 4. The Influence of Temperature Measure 5 ml of solution A (used in part 1) into a test tube. Measure 4 ml of water and 1 ml of 3% H 2 O 2 into a second test tube. Place both test tubes in a 250-mL beaker of warm water (50 C). After about 3 min, pour solution A and then the H 2 O 2 solution into a 50-mL beaker, stir, when the blue color appears, note and record the elapsed time. Compare the reaction time with that observed for the second reaction mixture you studied in part

4 Factors that Influence Reaction Rates. Part 5. Effect of Catalysts on the Rate of Decomposition of H 2 O 2 Using the same apparatus as in part 3a, place 20 ml of 3% H 2 O 2 in the flask. Fill a 13x100 mm test tube with water and invert it over the gas delivery tube. Then add 20 drops of 3 M CuCl 2 to the flask, replace the stopper, continuously swirl the contents, and note and record how long it takes to fill the test tube with the evolved gas. The gas coming over first will contain some air that was present in the flask. After 2 3 min, test the gas in the Erlenmeyer flask by inserting a glowing splint into the flask. Does the gas support combustion? What is the identity of the gas? Rinse the flask and repeat the experiment, adding only the H 2 O 2 solution to the flask. Continuously swirl the contents of the flask. Does the gas evolve more slowly? After a few minutes wait, you should be able to tell whether the CuCl 2 catalyzes the reaction. Now fill another test tube with water and place it over the gas delivery tube. Remove the stopper just long enough to add 2 drops of 3 M FeCl 3. Insert the stopper and continuously swirl the solution. Note and record the time required to fill the test tube with gas. Rinse the flask. Add to it 20 ml of 3% H 2 O 2, 1 drop of 3 M CuCl 2, and 1 drop of 3 M FeCl 3, in that order. Quickly replace the stopper, continuously swirl the solution in the flask, and record the time required to fill a 13! 100 mm test tube with gas. Repeat the procedure using 20 ml of 3% H 2 O 2, drops of 3 M CuCl 2, and 5 drops of 3 M FeCl 3. Quickly stopper the flask and continuously swirl the contents of the flask. Note and record the time required to fill a 13x100 mm test tube with gas. -4-

5 Report Form Name: Data and Observations Partner s Name: (if any)_lab Section: MW/TTH/M-TH/F (circle) 1. Effects of Concentration on Reaction Rates Record the elapsed time (number of seconds) for the solution A to turn blue after the addition of (a) 5 ml 3% H 2 O 2 s (b) 4 ml H 2 O + 1 ml 3% H 2 O 2 s The concentration of which reactant was changed from (a) to (b)?. (b) What is the relationship between initial concentration of H 2 O 2 and the reaction time? Briefly explain. 2. The Influence of the Nature of the Reactant (a) Record your observations below when Mg and Ca are added to water. Mg: Ca: (b) Based on your observation which metal is more reactive (c) What happened when the gas collected in the inverted test tube is brought near a burning splint or match? (d) What are the colors of each solution when phenolphthalein (acid/base indicator) is added?. Are the solutions acidic or basic? (e)write balanced chemical equations for the reactions(s) of Mg and Ca with water (f) Record the 1 st and 2 nd ionization energies of each metal and their sums. {Data may be obtained from your text or the CRC Handbook of Chemistry and Physics} Mg Ca First ionization Energy (units) Second ionization Energy (units) Sum (1 st and 2 nd IEs) (g) Which metal has the smaller Sum of ionization energies for the removal of two electrons? (h) From your observations and the data above, briefly explain how the reaction rates and ionization energies related? 3. The Influence of Surface Area in a Heterogeneous Reaction (a) Reaction of CaCO 3 + HCl (i) What gas is produced in the reaction? (ii) Record the times required to fill a test tube with gas for coarse CaCO 3 s pulverized CaCO 3 s -5-

6 Report Form (iii) Briefly explain why the reaction rate is faster for the pulverized CaCO 3? (iv) Write a balanced chemical equation for the reaction of CaCO 3 with HCl. (b) Reaction of Fe + CuCl 2 (i) Describe your observations (color change, etc.) and the rate of reactions of CuCl 2 solution with iron in the form of a nail: steel wool: (ii) What new substances are formed? Does the reaction evolve heat? YES NO {circle your answer} Which reaction is faster? nail steel wool {circle your answer} (iii) What evidence is there that the iron nail is reacting? (iv) Write a balanced net ionic equation for the reaction of Fe with Cu The Influence of Temperature (a) Solution A with H 2 O 2 Record the reaction time for the addition to solution A of 4 ml H 2 O + 1 ml 3% H 2 O 2 s From part 1, enter the reaction time at room temperature for (b) that had the same initial concentrations of reactants (4 ml H 2 O + 1 ml 3% H 2 O 2 ) s (i) From your observations what was the effect of increasing the temperature on the reaction time? Shorter Longer Slower Faster {circle your answer} the reaction rate? Shorter Longer Slower Faster {circle your answer} -6-

7 Report Form 5. Effect of Catalysts on the Rate of Decomposition of H 2 O 2 (a) Record the times required to fill a test tube with gas for each of the H 2 O 2 catalyst reaction mixtures. 20 drops of 3 M CuCl 2 : s No added catalyst: s 2 drops of 3 M FeCl 3 : s 1 drop 3 M CuCl drop 3 M FeCl 3 : s 5 drops 3 M CuCl drops 3 M FeCl 3 : s (b) Does the gas evolved support combustion?. How did you test for this? (c) What is the identity of the gas produced? Write a balanced chemical equation for the decomposition of hydrogen peroxide into water and oxygen. What evidence did you obtain that the combination of Cu 2+ and Fe 3+ is a more effective catalyst than either one alone? -7-

8 Report Form Questions/Problems 1. Imagine that you have a cube of CaCO 3 whose edge length is 1.0 cm. (a) What is the surface area of the cube (in square centimeters)? Now imagine that you pulverize it to form small cubes of edge length cm. (b) How many of the little cubes form? (c) What is the surface area of each little cube? (d) What is the new surface area (in square centimeters) of the pulverized CaCO 3? (e) If the rate of reaction of HCl with CaCO 3 (s) is proportional to the surface area of the CaCO 3, how much faster would the HCl react with the pulverized CaCO 3 than it did with the original large cube? 2. Why does the reaction rate of virtually all reactions increase with an increase in temperature? 3. If you were to make a glass of sweetened iced tea the old-fashioned way, by adding sugar and ice cubes to a glass of hot tea, which would you add first? -8-