EXPERIMENT 15 USING CONDUCTIVITY TO LOOK AT SOLUTIONS: DO WE AVE CARGED IONS OR NEUTRAL MOLECULES? rev 7/09 GOAL After you complete this experiment, you should have a better understanding of aqueous solutions and the forms that solutes may be in: charged ions or neutral intact molecules. In the Preliminary Activity, you will learn about measuring conductivity and what conductivity tells us about a solution. In Investigation 1, you ll choose a researchable question, design and conduct your experiment, and then share this with the class. In Investigation 2, you will apply conductivity to a titration. INTRODUCTION Many substances dissolve in water. Some substances dissolve so that we get individual neutral molecules surrounded by 2 O molecules. Other substances break apart into ions so that we get separate charged ions surrounded by 2 O molecules. We can easily distinguish between the two in lab. Solutions that contain many ions will conduct electricity well we call these electrolytes or strong electrolytes. Solutions that contain only neutral molecules will not conduct electricity we call these non-electrolytes. Solutions that contain only a few ions will conduct electricity weakly we call these weak electrolytes. Figure 1 below shows how O 2 dissolves in water. Notice that in the pure sample shown at left, we have neutral molecules composed of two oxygen atoms each. When O 2 dissolves in water, each molecule remains intact. Water molecules keep the O 2 molecules apart. Since everything in our beaker is neutral, this solution will not conduct electricity and will be a non-electrolyte. Most molecular substances, that is, substances that exist as neutral molecules when they are pure, will also exist as neutral molecules when they dissolve in water. Figure 2 below shows how NaCl dissolves in water. Notice that in the pure sample shown at left, we have charged ions: Na 1+ and Cl 1-. NaCl does not have molecules. When NaCl dissolves in water, each ion goes its separate way. In the solution, water molecules keep the ions apart from each other. This solution contains charged ions that will conduct electricity well, making this solution a strong electrolyte. Ionic salts are compounds like NaCl that are composed of charged ions. You can recognize the formula of a salt by looking for a metal paired with a non-metal or polyatomic ion. Na 1+ Cl 1- Cl 1- Na 1+ Na 1+ Cl 1- O O O Na 2 O 1+ O O Cl 1- Occasionally, a substance that exists as neutral molecules when pure will react with water when it dissolves. The most common examples are acids and bases. NO 3, nitric acid, is a good example. As shown in the figure 1
below, pure NO 3 exists as neutral molecules. [Note that while the formula for NO 3 contains a negatively charged polyatomic ion, it does not contain a positively charged metal ion or polyatomic ion. It is NOT a salt like NaCl.] When NO 3 dissolves in water, however, the water molecules pull 1+ off of the acid, leaving NO 3 1- behind. Ions form in solution and thus this solution will conduct electricity. NO 3 is an example of a strong acid, so named because 100% of its molecules fall apart to ions and result in a strong electrolyte solution. We can write this as NO 3 (l) 2 O 1+ (aq) + NO 1-3 (aq) Eqn 1 Unlike NO 3, most acids are considered weak acids because only a small percentage of their molecules will be broken apart into ions by the water. 2 S is a weak acid. Typically less than 1% of its molecules are broken into ions when placed in water. Since this results in a solution with only a very few charged ions and mostly neutral molecules, it will conduct electricity weakly and be known as a weak electrolyte. Since most of the 2 S molecules remain intact, not broken into ions, we most often represent its solution as intact 2 S(aq), not as 1+ (aq) and S 1- (aq). We usually want to show how most of the solute is present rather than emphasizing the few ions that form. Of course, not everything dissolves in water. These insoluble substances may be either molecular or ionic, but in either case we leave their formulas intact and include the appropriate state of matter. For example, we might have C 4 (g) or PbI 2 (s). Their presence will not cause a solution to conduct electricity. Beware the confusing O group! To predict conductivity, you need to decide if a substance is ionic or covalent. If you see O in a formula, think carefully. When O 1- is paired with a metal ion, as in NaO or Fe(O) 2, you have an ionic compound. When O is attached to non-metals, as in O or C 3 O, you have a covalent (molecular) compound. In Investigation 2, you will use conductivity to monitor a titration. The reaction you will study is between aqueous NaO, sodium hydroxide, and aqueous Cl, hydrochloric acid. NaO (aq) + Cl (aq) NaCl (aq) + 2 O (l) Eqn 2 At the start of your titration, your beaker will contain only Cl, so the conductivity you measure will allow you to decide if Cl exists as ions or molecules when dissolved in water. As you add NaO from the buret, it will be immediately consumed by reaction with the Cl to form our reaction products. At the equivalence point of the reaction, you will have only the products of Eqn 2 in your beaker. Thus, the conductivity at this point will allow you to decide if the products are intact or broken apart into ions. As you continue past the equivalence point, the excess NaO will also be present in the beaker and influence the conductivity we observe. AZARDS The solutions from Investigation 1 include acids. Be careful not to spill these on your skin. Be careful not to dilute or cross-contaminate solutions. Rinse your conductivity probe between samples and blot it dry with a tissue. In Investigation 2 you will use NaO, which is a strong base and Cl which is a strong acid. Both are corrosive and severe irritants, so be careful when handling. The solutions used in this experiment can go down the drain. 2
LABORATORY OBSERVATIONS AND DATA Your instructor will assign you a partner to work with in lab. Record your partner s name in your lab notebook. You will each write up your own lab report, however, so be sure that you both have a complete set of notebook entries and lab data before leaving lab. As always, include what you do and what you observe in your lab notes. PRELIMINARY INVESTIGATION Procedure: 1. Set the switch on the Conductivity Probe to the 0 20000 µs/cm conductivity range. 2. Connect the Conductivity Probe and the data-collection interface. 3. Rinse the electrode with distilled water and the blot it dry with a tissue. Repeat this cleaning before each new solution. Zero the probe by pressing Ctrl 0. You shouldn t need to re-zero again unless you have problems. 4. Determine the conductivity of a distilled water sample. a. Add 30 ml of distilled water to a clean 50 ml beaker. b. Place the tip of the electrode into the distilled water. The hole near the tip of the probe should be completely covered by the water. c. Start data collection. d. Stop data collection after about 15 seconds. e. Use the Statistics option to determine the mean conductivity value. 4. Determine the conductivity of a sodium chloride solution. a. Add about 0.25 g of NaCl to the distilled water. b. Stir to dissolve the NaCl. c. Repeat b e of Step 4 above to determine the conductivity of the NaCl solution. d. When done, flush the NaCl solution down the drain with excess water. Discussion: Return to the prelab side of the room with your lab partner. Working as a pair, answer these questions in your lab notebook. Consult the Introduction for help. 1. List 2 factors that affect the conductivity of a solution. 2. Define and give an example of strong electrolyte, weak electrolyte, and non-electrolyte. 3. Consult the List of Available Materials posted near the instructor s bench. List 2 Researchable Questions that you could explore using these materials. Participate in the discussion led by your instructor. As directed by the instructor, choose a Researchable Question from the list generated by the class. 3
INVESTIGATION 1 1. Prepare your Research Plan in your notebook. The sections should be a. Researchable Question b. ypothesis c. Variables (Manipulated, Responding, and Controlled variable(s)) d. Materials List e. Safety Issues f. Procedure g. Data Table 2. Get your plan approved by the instructor, make any needed changes, and then conduct your Investigation 1. 3. After you complete Investigation 1, return to the prelab side of the room and complete an Investigation Report form (available at the Instructor s bench). 4. Prepare to share the key points orally with the rest of the class. 5. During the class discussion of Investigation 1, you will be asked to briefly share your investigation with the rest of the class. You may modify or expand upon your conclusions based upon class discussion. Take notes on the investigations of others in your lab notebook since the lab report will include questions on the work of other groups. (All the Investigation Report forms will be available for viewing on Inquire the day after you complete this experiment.) INVESTIGATION 2: MONITORING A TITRATION BY CONDUCTIVITY 1. Obtain about 100 ml of 0.100 M NaO in a clean beaker 2. Obtain a 50-mL buret and rinse the buret with a few milliliters of the NaO solution. Repeat this two more times. Fill the buret to above the 0.00-mL level of the buret. Drain a small amount of NaO solution so it fills the buret tip and leaves the NaO at the 0.00-mL level of the buret. Although we don t usually worry about starting at exactly 0.00 ml, in this procedure doing so will make data recording simpler. 3. Use a clean 250-mL beaker to get 10.0 ml of Cl of unknown concentration from the dispenser. Add 90 ml of distilled water to the beaker. 4. Add a stir bar to your Cl beaker and position this on the stirring plate. Position the conductivity probe so that it is in the solution but away from the spinning stir bar. Position the buret above your beaker so that solution from the beaker won t hit the conductivity probe. Set the selection switch on the amplifier box of the probe to the 0-20000 μs range. 5. Your computer should already be in the LoggerPro program, with the file Exp 3 Electrolytes in the folder Roanoke Experiments open. The vertical axis of the graph has conductivity scaled from 0 to 2000 μs. The horizontal axis has volume scaled from 0 to 25 ml. Left click on the 1000 mark on the y-axis and drag down until 4000 comes into view. 6. Before adding NaO titrant, click Collect and monitor the displayed conductivity value (in μs). Once the conductivity has stabilized, click Keep. In the edit box, type 0, the current buret reading in ml. Press ENTER to store the first data pair for this experiment. 7. You will do 2 trials a more general one and a detailed one. You are now ready to begin the first titration. This process goes faster if one person manipulates and reads the buret while another person operates the computer and enters volumes. 4
a. Add 1.0 ml of 0.100 M NaO to the beaker. When the conductivity stabilizes, again click Keep. In the edit box, type the current buret reading read to the nearest 0.00 ml. Press ENTER. You have now saved the second data pair for the experiment. b. Continue adding 1.0 ml increments of NaO. With each increment, let the conductivity stabilize, click Keep, and enter the buret reading. Continue until you have added 25-mL of NaO. c. When you have finished collecting data, click Stop. Note approximately where your equivalence point is (to the nearest 0.0-mL). This will be important in the next trial. d. Print a copy of the Graph window using the popup menu File, Print Graph. Enter your name(s). Print a copy for each team member. Print a copy of the Table window using the popup menu File, Print Data Table. Enter your name(s). Print a copy for each team member 8. After printing your graphs and data tables from the first titration, you are now ready to begin your second titration. 9. Proceed as you did before, using a clean 250-mL beaker to get 10.0 ml of Cl of unknown concentration from the dispenser. Add 90 ml of distilled water to the beaker. 10. Refill the buret to the 0.00-mL with the NaO solution you prepared in Step 1. 11. Add a stir bar to your Cl beaker and position this on the stirring plate. Position the conductivity probe so that it is in the solution but away from the spinning stir bar. Position the buret above your beaker so that solution from the beaker won t hit the conductivity probe. Set the selection switch on the amplifier box of the probe to the 0-20000 μs range. a. Begin as you did in the first titration, clicking Collect. A box will come up and there will be several saving options. Click Erase and Continue to begin a new trial. Monitor the displayed conductivity value. Once the conductivity has stabilized, click Keep. In the edit box, type 0, the current buret reading in ml. Press enter to store the first data pair for this experiment. b. Add 1.0-mL of 0.100 M NaO to the beaker. When the conductivity stabilizes, again click Keep. In the edit box, type the current buret reading read to the nearest 0.00 ml. Press ENTER. You have now saved the second data pair for the experiment. c. Continue adding NaO in 1.00-mL increments until you are within 1.0-mL of the equivalence point you located in your first titration. e. After this, use 2-drop increments (~0.1 ml) until the minimum conductivity has been reached at the equivalence point. Enter the buret reading after each 2-drop addition. When you have passed the equivalence point, continue using 2-drop increments until you are 1.0- ml past your equivalence point. f. Now use 1.0-mL increments until 25 ml of NaO solution have been added. g. When you have finished collecting data, click Stop. Note approximately where your endpoint is. Print a copy of the Graph window using the popup menu File, Print Graph. Enter your name(s). Print a copy for each team member. Print a copy of the Table window using the popup menu File, Print Data Table. Enter your name(s). Print a copy for each team member 12 Flush all of your used and unused solutions down the drain with excess water. Rinse your buret with distilled water and return it to the supply bench. 5
RESULTS AND DISCUSSION For Investigation 1, write several paragraphs that describe your investigation, including a summary of each of the points in your research plan. Include a clean table that clearly shows your data, including units. Now describe what you can conclude from your investigation. Make sure that you refer to your data as the basis for your conclusions. If your data is inconclusive or if you had experimental difficulties, describe these. For Investigation 2, examine your computer generated graph for the second titration. The equivalence point of the titration is at the minimum conductivity. You can find the approximate volume from the graph. Now examine the data table closely at this approximate volume. Find the lowest value for conductivity. The volume at this data point is the volume of NaO needed to reach equivalence. Now prepare a data table based on your second titration with the following 6 rows: Volume of Cl; Molarity of NaO; Volume of NaO; Moles NaO; Moles Cl; and Molarity Cl. Note that the first two rows, the volume of Cl and molarity of NaO are given in the procedure. The volume of NaO is the equivalence volume that you just determined from your graph and data table. Calculate the moles of NaO from the volume and molarity of NaO. Since our titration is a 1:1 reaction, the moles of Cl at equivalence is equal to the moles of NaO. Finally, calculate the molarity of the Cl solution from its moles and volume. Be sure to include sample calculations in your report. Attach your computer printouts. QUESTIONS 1. Write several paragraphs, summarizing things your class collectively learned about conductivity and types of electrolytes from Investigation 1. Use examples from this experiment to illustrate your points. Your writing must be clear, grammatically correct, and chemically correct. Consult the Investigation Report Forms on Inquire to find the information you need. (a) Find an investigation that compares several different compounds. Describe the experiment, what the data showed, and what you can conclude from it. (b) Find an investigation that compares solutions with different concentrations. Describe the experiment, what the data showed, and what you can conclude from it. (c) Find one more investigation that you found interesting. Describe the experiment, what the data showed, and what you can conclude from it. 2. In Investigation 2, you studied the reaction NaO (aq) + Cl (aq) NaCl (aq) + 2 O (l) (a) Which of these four substances is actually in the beaker at the equivalence point? [See the Introduction for help.] Does the conductivity at equivalence suggest that you have many or few dissolved ions at equivalence? If you have ions, give their formulas with charges. (b) Prior to reaching the equivalence point, which of the reactants was in excess (not the limiting reagent)? Does the conductivity prior to equivalence suggest that this exists as ions or intact molecules? If you have ions, give their formulas with charges. (c) After equivalence, the other reactant is in excess. Does the conductivity after the equivalence point suggest that this reactant exists as ions or intact molecules? 6
Investigation Report Form Experiment Number: Team Members: Researchable Question: Variables Procedure Method Summary Data What was learned/concluded 7