ELECTRICAL CONDUCTION IN SOLUTIONS

Transcription

1 ELECTRICAL CONDUCTION IN SOLUTIONS Partially adapted from "Chemistry with Computers" Vernier Software, Portland OR, 1997 INTRODUCTION Conductivity Electricity is typically thought of as the conduction or movement of electrons through a metallic solid (such as a copper wire) in response to the application of an electric field. However, electricity is the motion of any charged particle. In fact, electricity is observed when an electric field is applied to a liquid solution containing electrolytes (charged chemical species). Compounds that dissolve in water and form a large number of ions are known as strong electrolytes. Compounds producing a few ions are called weak electrolytes. Compounds that do not form any ions are called non-electrolytes. A compound s electrolytic strength is directly related to the amount of current that can be propagated through solution. This current can be measured with a conductivity probe. An electric field is created between two oppositely charged metal plates (electrodes) located on each side of the hole near the bottom of the probe body (Figure 1). Figure 1. Figure 2. Cations move toward the negative plate and anions move toward the positive plate. Like a copper wire in a circuit, a solution containing ions completes the electrical circuit between the electrodes (Figure 2). Eventually, all the ions would reach the oppositely charged plate, so the polarity of the plates must reverse (the negative plate becomes the positive plate and vice versa) at regular intervals to start the movement of the electrolytes all over again. 1

2 The solution conductance has the units of inverse ohms (Ω -1 ) which are formally defined as microsiemens (µs). The magnitude of the conductivity value is directly related to electrolytic strength and concentration. A stronger electrolyte has more ions in solution, and, therefore, a higher conductance value. Likewise, a more concentrated solution of electrolyte has more ions in solution, and, therefore, a higher conductivity value. Furthermore, the magnitude of the conductivity value is usually directly related the mass of the electrolyte. The conductivity of a solution is typically becomes greater for electrolytes going down a group on the periodic table. A notable exception to this trend is H +, the lightest of all ions, has the Figure highest 3 conductivity. The reason for this is that H + tunnels through aqueous solution by chemical reaction with H 2 O: H + reacts with H 2 O forming hydronium ions, H 3 O +, which then go on to react with another H 2 O molecule forming a new H 3 O + and regenerating the original H 2 O molecule. This reaction sequence repeats itself over and over until H + (probably not the original H + ) reaches the negative electrode. (Hydroxide ions, OH, also have a similar interaction with H 2 O resulting in the second highest conductivity of all ions.) All other electrolytes have to physically push through the H 2 O molecules (without chemical reaction) as they move toward the oppositely charged electrode, resulting in significantly lower conductivity values. Using Conductivity to Find an Equivalence Point The equation for the reaction in this experiment is: Ba 2+ (aq) + 2 OH (aq) + 2 H + (aq) + SO 4 2 (aq)! BaSO 4 (s) + 2 H 2 O(l) Before reacting, Ba(OH) 2 and H 2 SO 4 are almost completely dissociated into their respective ions. Neither of the reaction products, however, is significantly dissociated. Barium sulfate is a precipitate and water is predominantly molecular. As 0.2 M H 2 SO 4 is slowly added to Ba(OH) 2 of unknown concentration, changes in the conductivity of the solution will be monitored using a conductivity probe. When the probe is placed in a solution that contains ions, and thus has the ability to conduct electricity, an electrical circuit is completed across the electrodes that are located on either side of the hole near the bottom of the probe body (Figure 3). This results in a conductivity value that can be read by the interface. The unit of 2

3 conductivity used in this experiment is also microsiemens per centimeter, or µs/cm. SAFETY Wear safety goggles and lab aprons at all times in the lab. Some acids are very corrosive and can cause burns; wash any affected areas thoroughly with cold water. Some compounds (AlCl 3, ethylene glycol, Ba(OH) 2 ) are toxic when ingested; wash hands thoroughly before leaving lab. Ethanol is flammable; extinguish any open flames or spark sources. Before starting the experiment, the TA will asks you to do a quick demonstration or talk-through one of the following: 1) The set up in part A of the experiment (what to clamp, where the buret, stir plate, beaker, and conductivity probe go)! two student demo 2) How to use a Büchner funnel (how do you disconnect the hose, what do you need to do with the filter paper first, what the Büchner funnel looks like)! two student demo 3) How to use a digital scale to correctly get the mass of BaSO 4 (how do you tare the watch glass, what happens if you weigh a hot object, how many digits do you record from the scale) PROCEDURE Note: During all parts of this experiment the conductivity probe should be rinsed and placed in the beaker of water when not in use. All data files should be attached to your ELN. Part A. Finding an Equivalence Point. Work in groups of three. Two people begin Part A. The third person starts making the solutions in Part B Step #3. Make sure your TA has turned on the oven at the beginning of lab to ~110 o C. Ba(OH) 2 reacts quickly with dissolved CO 2 to form unwanted BaCO 3 (aq). Once your TA has opened the sealed, degassed Ba(OH) 2 solution, you must perform the procedures in Part A relatively quickly. 1. Using a graduated cylinder, measure out approximately ml of 0.2 M H 2 SO 4 solution into a 250 ml beaker. Record the exact H 2 SO 4 concentration in your data table. CAUTION: 3

4 H 2 SO 4 is a strong acid, and should be handled with care. Obtain a 50 ml buret and prerinse the buret with a few ml of the H 2 SO 4 solution. Use a utility clamp to attach the buret to the ring stand. Fill the buret a little above the 0.00 ml level of the buret. Drain a small amount of H 2 SO 4 solution so it fills the buret tip (no bubbles) and leaves the H 2 SO 4 at or just below the 0.00 ml level of the buret. Dispose of the waste solution into a 1000 ml beaker to be neutralized later. 2. Arrange the buret, conductivity probe, and stir plate as shown in Fig. 3. Set the selection switch on the amplifier box of the probe to 0-20,000 µs/cm range. Connect the conductivity probe to Ch1 of the LabQuest2. If there are issues with LabQuest2 picking up the conductivity probe, select the sensors menu, click sensor setup, & pick Conductivity MICS 3. Using a squeeze bottle, rinse the probe with deionized water and then blot dry with a KimWipe. Prepare the LabQuest2 for data collection by clicking on mode, and changing the mode to events with entry, name: volume, units: ml. To check that LabQuest2 is registering the correct conductivity range, submerge the probe into a small beaker of tap water (not DI water). The range should be between µs. Clean the probe again using a squeeze bottle to rinse with deionized water and then blot dry with a KimWipe. 4. This step can only be accomplished once all students in the lab are ready as stated above, Ba(OH) 2 solutions react with CO 2 to form the insoluble BaCO 3. Carbon dioxide dissolves in the solution quickly from room air. Avoid exhaling near the solution. CAUTION: Ba(OH) 2 is toxic. Handle it with care and wear gloves! Your TA will dispense ml of Ba(OH) 2 solution of unknown concentration into your 100 ml graduated cylinder. Perform the next few steps quickly to avoid BaCO 3 formation. Transfer the solution to a clean, dry 400 or 500 ml beaker. Then add 120 ml of DI water to the beaker. 5. Place the beaker beneath the buret on the stir plate and clamp the conductivity probe in place as shown in Figure 3. The conductivity probe a should extend down into the Ba(OH) 2 solution to just above the stirring bar and should be clamped to the side of the beaker. (If the H 2 SO 4 drips directly on the probe, false readings will occur.) 4

5 6. Before adding H 2 SO 4 titrant, click and monitor the displayed conductivity value (in µs/cm). Once the conductivity has stabilized, click. 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 are now ready to begin the titration. This process goes faster if one person manipulates and reads the buret while another person operates the computer and enters volumes. a. Add 2.0 ml of 0.2 M H 2 SO 4 to the beaker. When the conductivity stabilizes, again click. In the edit box, type the current buret reading. Press ENTER. You have now saved the second data pair for the experiment. b. Continue adding 2.0 ml increments of H 2 SO 4, each time entering the buret reading, until the conductivity has dropped below 2000 µs/cm. c. Continue adding 1.0 ml increments of H 2 SO 4, each time entering the buret reading, until the conductivity has dropped below 800 µs/cm. d. After this, use 5-drop increments (~0.5 ml) until the minimum conductivity has been reached at the equivalence point. Enter the volume after each 5-drop addition. When you have passed the equivalence point, continue using 5-drop increments until the conductivity is greater than 800 µs/cm again. e. Now use 1.0 ml increments until the conductivity is 2000 µs/cm. Then use 2.0 ml increments until all of the H 2 SO 4 solution has been used up. 8. When you have finished collecting data, click. Save the data file to your ELN. 9. Place beaker into an ice bath for 10 min. This lowers the solubility of BaSO 4 and settles the solid out of solution. 10. Using a Büchner funnel, filter the solution slowly (avoid stirring or swirling), trying to keep the solid BaSO 4 at the bottom of the beaker. Use a spatula and cold DI water to rinse the final amount of BaSO 4 from the beaker. Let it stay under vacuum for 5 minutes. 11. Scrape the BaSO 4 (white powder) onto a tared watch glass. Place the watch glass into the oven for 10 min. Remove carefully from oven as directed by your TA. Let cool. Weigh. 12. Place about ¼ inch of solid NaHCO 3 in the bottom of another 1000 ml beaker. Add water until it barely covers the NaHCO 3 and stir to make a slurry. Slowly pour the acidic solution 5

6 into the NaHCO 3 slurry and stir. When the foaming stops, check the solution with ph paper. If ph = 7 pour the solution down the drain. If ph is less than seven, repeat NaHCO 3 neutralization. Part B. Preparing Solutions. 1. The conductivity probe should be attached to the LabQuest2 and resting in a beaker of deionized water. Click on the red box that reads CH1:Conductivity in the pull down menu choose zero. This calibrates the probe. 2. Measure and record the conductance of tap water. It should be between µs higher than the deionized water. Why? If it is not, inform your TA. 3. You and your partners will measure the conductance for the six solutions below. Two are already prepared for you. You each need to create one solution from each of the set listed below. Deionized (not regular tap) water should be used as the solvent for all. Create only 10 ml of each solution using volumetric glassware. Clearly outline the calculations and procedures used to create these solutions in your lab notebook. Set I: 0.10 M NaCl, 0.10 M LiCl, & 0.10 M KCl. Solids will be provided M LiCl is already prepared. Set II: 0.10 M H 3 PO 4, 0.10 M HC 2 H 3 O 2, and 0.10 M HCl. Stock solutions (1.0 M) will be provided M H 3 PO 4 is already prepared. Part C. Measuring Conductance. 4. Measure the conductivity of deionized water ( the blank ). A blank theoretically should not have any conductivity. However, it is highly likely that contaminants may be present that will result in a nonzero conductivity reading. The conductivity value of the blank taken before each solution should be reported with that solution s conductivity value. To clean the conductivity probe should be rinsed thoroughly with deionized water from a wash bottle. Blot the outside of the probe end dry using a tissue. Do NOT dry the inside of the hole near the probe end. (The rinsings may be discarded in the sink.) 6

7 Carefully place the conductivity probe into a beaker of deionized water. The hole near the probe end has two electrodes positioned on either side which must be submerged as shown in Figure 1 to insure correct conductivity readings. Once the conductivity reading in the Meter window has stabilized, record the value in a data table. Discard the deionized water. Repeat this process before each solution s conductivity is measured. 5. Each partner member should measure the conductivity for each of the solutions prepared (creating 3 conductivity measurements of each solution). Pour each solution into a clean, dry, and labeled test tube. Carefully raise each test tube and its contents up around the blotted dry conductivity probe until the hole near the probe end is completely submerged in the solution. Briefly swirl the test tube contents. Once the conductivity reading in the Meter window has stabilized, record the value in the data table. Before testing the next solution, repeat step Group members should share all data with each other and with one other group. Report the data in two tables: your group and the other group. 7. Rinse the conductivity probe and return it to the beaker of water. Keep the NaCl solution for Part D. Discard all other solutions in the sink (the salts are the components found in tap water). Part D. Changing Concentrations. 8. Open a new file on the LabQuest2. In, select Sensor. In Sensor Setup window, change the CH1 sensor to Conductivity 200 MICS. Change the mode to events with entry, name: volume, units: drops. 9. Add 70 ml of DI water to a clean 100 ml beaker. Then: 7

8 Click. Place the conductivity probe in the beaker. Make sure the conductivity is set to 200 MICS from the drop down menu. Click once the conductivity reading stabilizes. Type 0 in the edit box (for 0 drops added). Press the OK key to store this data pair. (This reading is the conductivity of the water before any salt solution is added.) 10. Create a plot of conductance vs. volume for the addition of NaCl solution. Add 1 drop of the 0.1 M NaCl to the water in the beaker. Gently stir the solution with the conductivity probe. Click once the conductivity reading stabilizes. Type 1 (for the total drops added) in the edit box and press OK. Add another drop, entering 2 this time. 11. Continue this procedure, adding 1-drop portions of the salt solution, measuring conductivity and entering the total number of drops added until a total of 8 drops have been added. Click when data collection is finished. The beaker contents can be discarded in the sink. Rinse the probe tip with deionized water and carefully blot it dry with a tissue. 12. Transfer the data to your computer and use Excel to find a best-fit linear regression line. Record the value of the slope, m, for each solution. (The linear regression statistics are displayed in a floating box for each of the data sets.) Attach the graph into your ELN with labeled axis and an appropriate title. Make sure to clear your address and password of the LabQuest2 so others can t access your account. Shutdown the LabQuest2 and not simply put it to sleep. To shutdown the LabQuest2: press the home key, select System! Shut Down! OK. 8

9 DISCUSSION Part A. (1) From the data table and graph that you created, determine the volume of H 2 SO 4 added at the equivalence point. The graph should give you the approximate volume at this point. The precise volume of H 2 SO 4 added can be determined by further examination of the data table for the minimum conductivity. Record the volume of H 2 SO 4. (2) Calculate moles of H 2 SO 4 added at the equivalence point. Use the molarity (M) of the H 2 SO 4 and its volume (L). (3) Calculate the moles of BaSO 4 present at the equivalence point. What is the theoretical yield of BaSO 4 in grams? (4) Calculate the percent yield of BaSO 4 recovered. Part C. (5) Based on conductivity values, classify the compounds in each set of solutions as molecular, ionic, or acids. What is the level of dissociation (complete, varying, or none) of each set? Set Classification Level of Dissociation I II (6) List the Set I solutions in order of highest to lowest conductivity. Provide a simple explanation of the trend. (7) Why do conductivity values of tap water and deionized water differ? What conductivity would a solution of EtOH most likely have? Hint: Does EtOH dissociate in solution? 9

10 (8) All Set II solutions have the same concentration (0.1 M), however, the conductivity values differ by such large amounts. a. Arrange compounds by increasing electrolytic strength using conductivity values. b. Write an equation for the dissociation of each compound. For complete dissociation use a single arrow,!, and if partial dissociation occurs use equilibrium arrows,!. c. For H 3 PO 4, does the subscript "3" of hydrogen result in additional ions in solution as it did for the solutions in Part A? Explain. (9) Using both groups data, calculate the mean and standard deviation for each solution in Sets I and II. Perform the Q-test on any value that may be an outlier. (At least 3 Q-tests should be done.) Part D. (10) Describe the change in conductivity as the salt concentration was increased by the addition of more drops. What kind of mathematical relationship exists between conductivity and concentration? QUALITATIVE ERROR ANALYSIS 1. What modifications could be made to the procedure to better account for random (indeterminate) errors? 2. List three potential systematic (instrumental, methodological, or personal) errors that could be made in this experiment. (Note: Be specific, systematic errors are in the details. For example, losing your solution because you knocked over the cuvette is not a systematic error it s a gross one.) 3. Did any gross errors occur? Did you mess up? Did the equipment or instrumentation fail? 10