Properties of Solutions: Electrolytes and Non-Electrolytes

Transcription

1 Please type ALL data directly into your lab notebook Properties of Solutions: Electrolytes and Non-Electrolytes In this experiment, you will discover some properties of strong electrolytes, weak electrolytes, and non-electrolytes by observing the behavior of these substances in aqueous solutions. You will determine these properties 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 (see Figure 1). This results in a conductivity value that can be read by the computer. The unit of conductivity used in this experiment is the microsiemens per centimeter, or µs/cm. Figure 1 The size of the conductivity value depends on the ability of the aqueous solution to conduct electricity. Strong electrolytes produce large numbers of ions, which results in high conductivity values. Weak electrolytes result in low conductivity, and non-electrolytes should result in no conductivity. In this experiment, you will observe several factors that determine whether or not a solution conducts, and if so, the relative magnitude of the conductivity. Thus, this simple experiment allows you to learn a great deal about different compounds and their resulting solutions. In each part of the experiment, you will be observing a different property of electrolytes. Keep in mind that you will be encountering three types of compounds and aqueous solutions: Ionic Compounds These are usually strong electrolytes and can be expected to 100% dissociate in aqueous solution. Example: NaNO 3 (s) Na + (aq) + NO 3 (aq)

2 Molecular Compounds These are usually non-electrolytes. They do not dissociate to form ions. Resulting solutions do not conduct electricity. Example: CH 3 OH(l) CH 3 OH(aq) Molecular Acids These are molecules that can partially or wholly dissociate, depending on their strength. Example: Strong electrolyte H 2 SO 4 H + (aq) + HSO 4 (aq) (100% dissociation) Example: Weak electrolyte HF H + (aq) + F (aq) (<100% dissociation) OBJECTIVES In this experiment, you will Write equations for the dissociation of compounds in water. Use a Conductivity Probe to measure the conductivity of solutions. Determine which molecules or ions are responsible for conductivity of solutions. Investigate the conductivity of solutions resulting from compounds that dissociate to produce different numbers of ions. MATERIALS computer H 2 O (distilled) Vernier computer interface 0.05 M NaCl Logger Pro 0.05 M CaCl 2 Vernier Conductivity Probe 0.05 M AlCl ml beaker 0.05 M HC 2 H 3 O 2 wash bottle with distilled water 0.05 M H 3 PO 4 tissues 0.05 M H 3 BO 3 ring stand 0.05 M HCl utility clamp 0.05 M CH 3 OH (methanol) H 2 O (tap) Glacial acetic acid PROCEDURE 1. Obtain and wear goggles! CAUTION: Handle the solutions in this experiment with care. Do not allow them to contact your skin. Notify your teacher in the event of an accident. 2. The Conductivity Probe is already attached to the interface. It should be set on the µs/cm position. 3. Prepare the computer to monitor conductivity by opening the file 13 Electrolytes from the Chemistry with Vernier folder. 4. Obtain the Group A solution containers. The solutions are: 0.05 M NaCl, 0.05 M CaCl 2, and 0.05 M AlCl Measure the conductivity for each of the solutions.

3 a. Carefully raise each vial and its contents up around the Conductivity Probe until the hole near the probe end is completely submerged. Important: Since the two electrodes are positioned on either side of the hole, this part of the probe must be completely submerged. b. Briefly swirl the beaker contents. When the reading has stabilized, record the value. c. Before testing the next solution, clean the electrodes by surrounding them with a 250 ml beaker and rinsing them with distilled water. Blot the outside of the probe end dry using a tissue. It is not necessary to dry the inside of the hole near the probe end. 6. Obtain the four Group B solution containers. These include 0.05 M H 3 PO 4, 0.05 M HC 2 H 3 O 2, 0.05 M H 3 BO 3, and 0.05 M HCl. Repeat the Step 5 procedure. 7. Obtain the five Group C solutions or liquids. These include 0.05 M CH 3 OH, 0.05 M C 2 H 6 O 2, distilled H 2 O, tap H 2 O and glacial acetic acid. Repeat the Step 5 procedure. DATA TABLE Solution Conductivity (µs/cm) A - CaCl 2 A - AlCl 3 A - NaCl B - HC 2 H 3 O 2 B - HCl B - H 3 PO 4 B - H 3 BO 3 C - H 2 O distilled C - H 2 O tap C - CH 3 OH C- Glacial Acetic Acid Part 2 1. Prepare the computer for data collection by opening the file 14 Conductivity Solutions from the Chemistry with Vernier folder. 2. Add 70 ml of distilled water to a clean 100 ml beaker. Obtain a dropper bottle that contains 1.0 M NaCl solution. 3. Before adding any drops of solution: a. Click. b. Carefully raise the beaker and its contents up around the Conductivity Probe until the hole near the probe end is completely submerged in the solution being tested. Important: Since the two electrodes are positioned on either side of the hole, this part of the probe must be completely submerged. c. Monitor the conductivity of the distilled water until the conductivity reading stabilizes.

4 d. Click, and then lower the beaker away from the probe. Type 0 in the edit box (for 0 drops added). Press the ENTER key to store this data pair. This gives the conductivity of the water before any salt solution is added. 4. You are now ready to begin adding salt solution. a. Add 1 drop of NaCl solution to the distilled water. Stir to ensure thorough mixing. b. Raise the beaker until the hole near the probe end is completely submerged in the solution. Swirl the solution briefly. c. Monitor the conductivity of the solution until the reading stabilizes. d. Click, and then lower the beaker away from the probe. Type 1 (the total drops added) in the edit box and press ENTER. 5. Repeat the Step 4 procedure, entering 2 this time. 6. Continue this procedure, adding 1-drop portions of NaCl solution, measuring conductivity, and entering the total number of drops added until a total of 8 drops have been added. 7. Click when you have finished collecting data. Dispose of the beaker contents as directed by your teacher. Rinse the probe tip with distilled water from a wash bottle. Carefully blot the probe dry with a tissue. 8. Prepare the computer for data collection. From the Experiment menu, choose Store Latest Run. This stores the data so it can be used later, but it will be still be displayed while you do your second and third trials. 9. Repeat Steps 4 8, this time using 1.0 M AlCl 3 solution in place of 1.0 M NaCl solution. 10. Repeat Steps 4 8, this time using 1.0 M CaCl 2 solution. 11. Click on the Linear Fit button,. Be sure all three data runs are selected, then click. A best-fit linear regression line will be shown for each of your three runs. In your data table, record the value of the slope, m, for each of the three solutions. (The linear regression statistics are displayed in a floating box for each of the data sets.) 12. To print a graph of concentration vs. volume showing all three data runs: a. Label all three curves by choosing Text Annotation from the Insert menu, and typing sodium chloride (or aluminum chloride, or calcium chloride ) in the edit box. Then drag each box to a position near its respective curve. b. Print a copy of the graph, with all three data sets and the regression lines displayed. Enter your name(s) and the number of copies of the graph you want. Part 3 1. Use your small graduated cylinder to measure out 5.0 ml of deionized water in a small beaker. Add 15 drops 0.05 M H 2 SO 4 (sulfuric acid) to the beaker and stir the solution. 2. Measure and record the conductivity of this dilute sulfuric acid solution. Is this an electrolyte?

5 3. Add 1 drop 0.05 M Ba(OH) 2 (barium hydroxide) solution to the beaker, stir, and again measure the conductivity. (The barium hydroxide solution might be somewhat cloudy due to a small amount of insoluble barium carbonate, BaCO 3. This will cause no harm.) 4. Add a second drop of the Ba(OH) 2 solution, stir, and again obtain a conductivity reading. Repeat this procedure until a total of 75 drops Ba(OH) 2 solution have been added. Measure and record the conductivity each time. Drops of Ba(OH) 2 may stopped being added when the conductivity of the resulting solution has a conductivity approximately equal to the initial solution. 5. Make a graph plotting solution conductivity on the y-axis and the volume (drops) barium hydroxide added on the x-axis. PROCESSING THE DATA PART A 1. Based on your conductivity values, do the Group A compounds appear to be molecular, ionic, or molecular acids? Would you expect them to partially dissociate, completely dissociate, or not dissociate at all? 2. Why do the Group A compounds, each with the same concentration (0.05 M), have such large differences in conductivity values? Hint: Write an equation for the dissociation of each. Explain. 3. In Group B, do all four compounds appear to be molecular, ionic, or molecular acids? Classify each as a strong or weak electrolyte, and arrange them from the strongest to the weakest, based on conductivity values. 4. Write an equation for the dissociation of each of the compounds in Group B. Use for strong; for weak. 5. For H 3 PO 4 and H 3 BO 3, does the subscript 3 of hydrogen in these two formulas seem to result in additional ions in solution as it did in Group A? Explain. 6. In Group C, do all four compounds appear to be molecular, ionic, or molecular acids? Based on this answer, would you expect them to dissociate? 7. How do you explain the relatively high conductivity of tap water compared to a low or zero conductivity for distilled water? PART B Explain the slopes of the lines in the graph produced in terms of ions in the materials studied PART C Explain the changes in the conductivity and the changes in the appearance of the solution. This acid-base reaction can be written as: Ba(OH) 2 (aq) + H 2 SO 4 (aq) BaSO 4 (s) + 2H 2 O(l)