7# .Jr-;ppppppXpppqq.qJqJzqxpr< r\r*rprrtrrrrrrrrrExperiment IV Electrical Conductivity of Aqueous Solutions PRE-LAB ASSIGNMENT Reading: Before coming to your discussion section, read the following: 1. Sections 4-4 through 4-6 in Olmstead and Williams. 2. The remainder of this experiment in this manual. Pre-lab assignment to hand in: Prepare written responses to the following questions and submit your work to your instructor at the beginning of your discussion period: 1. Why dont we immerse the conductivity probe above the immersion mark? 2. Why does a 0.01 M HCl solution have a greater conductivity than a 0.01 M NaCl solution? 3. Calculate the conductivity of a 0.01 M NaOH solution. 4. How many ml of a 0.02 M HBr solution must be added to 150 ml of a 0.0066 M KOH solution to neutralize all the OH-? List the ions present in the solution at this point. INTRODUCTION The purpose of this lab is to help you discover the nature of aqueous solutions by investigating the relationship between conductivity and strong and weak electrolytes. In this lab you will add increasing amounts of either acid or base to several electrolyte solutions. After each addition you will measure the conductivity of the solution. From the conductivity data, you will work to deduce the nature of the chemical reaction that occurred in the experiment. This experiment is a "discovery" -type experiment similar to Experiments II and III. The procedure will be carefully described, but the analysis of the data is left purposely vague. You will work in small groups to decide how best to work up the data. In the process you will have the chance to discover some principles, to use what you have learned in lecture, and to practice thinking about manipulative details and theory at the same time. Plotting your data in an appropriate manner should verify the nature of the chemical reactions that occurred and the relationship between conductivity and electrolyte composition. What is conductivity? The electrical conductivity of a solution is the quantity of electricity which will flow through a solution per unit area, per unit potential gradient, per unit time usually expressed as mho/cm. Note: mho (pronounced "mo") was derived as a unit of conductivity by reversing the letters in "ohm", the unit of resistance. In other words, if we were to plug in the circuit shown in Figure 1, the conductivity of the solution in the beaker would be proportional to the current that flows from electrode 1 to electrode 2. In order for current to flow, ions must be present in solution to carry the charge from one electrode to another. Increasing the number of ions in solution will increase the amount of charge which can be carried between electrodes and will increase the conductivity. Therefore, conductivity measurements provide us with a tool to learn something about the number of ions in a solution. SuperPaint!Macintosh HD:CHEM 141:figure 1 cond lab!Paint(119,41:493,396)  Figure 1. Schematic of a simple conductivity measurement system. Another factor in conductivity measurements is that not all ions carry charge (conduct electricity) equally. Small ions such as H+ and OH- move through solution very rapidly and are very good charge carriers. Larger ions such as NH4+ and Cl- move through solution at a slower rate and therefore do not conduct electricity as well. Uncharged species in solution do not carry any charge. Fortunately, chemists over the years have measured the ability of individual ions to carry charge (ionic conductance) and these data are available in chemical reference books. Table 1 is an abbreviated table of ionic conductances. Table 1. Ionic Conductances of Selected Ions . Ion Ionic Conductance (mho l eq.-1 cm-1.) cations H+ 0.34982 Na+ 0.05011 NH4+ 0.0735 K+ 0.0735 anions OH- 0.1986 Cl- 0.07635 CH3COO- 0.0409 Br- 0.0781 To calculate the conductivity of a solution you simply multiply the concentration of each ion in solution by its ionic conductance and charge and add these values for all ions in solution. Units of conductivity are mho/cm (mho = 1/W). Units of ionic conductance are mho l eq.-1 cm-1 where l is liters and eq. is an equivalent of charge or the ion concentration times its charge. An example calculation for a 0.00500 M NaCl solution is shown below. Cond. due to Na+ = 0.00500 mol/l * 1 eq./mol * 0.05011 mho l/eq cm. = 2.51x10-4 mho/cm Cond. due to Cl- = 0.00500 mol/l * 1 eq./mol * 0.0764 mho l/eq.cm = 3.82x10-4 mho/cm Total Conductivity = 2.51x10-4 + 3.82x10-4 = 6.33x10-4 mho/cm = 633 mmho/cm Notice that the final result is reported in mmho/cm which is 10-6mho/cm and is also the unit used by our conductivity meters. How can we use conductivity to study chemical reactions? Since conductivity is a function of both the concentration and the composition of the solution being measured, we can use conductivity to follow chemical reactions. Consider the titration of KOH with HBr shown in Figure 2. KOH(aq) + HBr(aq) ---> KBr(aq) + H2O ionic equation: K+ + OH- + H+ + Br- ---> K+ + Br- + H2O net ionic: OH- + H+---> H2O Initially the solution conductivity is high, mostly because OH- has a large ionic conductance. However, as OH- is neutralized the conductivity falls. After all the OH- is neutralized the H+ concentration increases driving the conductivity back up. Figure 2. Titration of 150 ml of 0.0066 M KOH with 0.02 M HBr From the figure it is easy to see the point where all the OH- has been neutralized (ie. the equivalence point of the titration). Something we will leave you to consider is the role that spectator ions play in contributing to the observed conductivity over the course of the titration. Outline of this experiment. You will work in pairs and perform conductivity titrations on three different solutions. Every pair will perform a titration of NaOH with HCl. Half of the class will also titrate NH3 with HCl and NH4Cl with NaOH. The other half of the class will titrate acetic acid (CH3COOH)with NaOH and sodium acetate (CH3COO Na) with HCl. At the end of class you will work in groups of four to analyze your data so that each group has all of the titrations. The objective of the lab is to use the conductivity data to predict the ionic composition of your test solutions before, during, and after the titrations. PROCEDURE 1. Work alone or in pairs. 2. Figure 3 is a cartoon of the conductivity pen that you will use in this experiment. Please be careful with the pens, as they are expensive. Do not immerse the pen above the immersion limit (water will short out the electronics). Take the cover off of the pen and place the end of the pen into a 250 ml beaker which is filled with 75 ml of distilled water. Allow the pen to soak for 5 minutes while you get your other solutions ready.  Figure 3. Conductivity pen. 3. You will be performing three titrations. Table 2 details the starting solution and titrant for each titration. Notice that there are two concentrations of NaOH present in the lab; be careful to use the correct one. Table 2. Solutions to be titrated. Group Starting Solution Titrant . everyone 150 ml of 0.0066 M NaOH 0.02 M HCl group 1 150 ml of 0.0066 M NaC2H3O2 0.02 M HCl group 1 150 ml of 0.0066 M HC2H3O2 0.02 M NaOH group 2 150 ml of 0.0066 M NH3 0.02 M HCl group 2 150 ml of 0.0066 M NH4Cl 0.02 M NaOH Place 150 ml of the starting solution into a 250 ml beaker. You may use the graduations on your beaker to measure the solution volume. Add a stir bar to the beaker and stir the solution slowly using a magnetic stirrer. 4. Hold the conductivity pen in the stirring solution and read the conductivity on the display. The number on display multiplied by ten is the conductivity of the solution in units of mmho/cm. If the conductivity of the solution is too high the pen display will go off scale (the meter displays a 1 followed by three blanks). In this case, press the range switch while you are taking a measurement and multiply your result by 100 instead of 10 to get mmho/cm. 5. Titrate the starting solution by adding 10 ml increments of titrant to the solution (a graduated cylinder is OK for measuring titrant volume). Allow the solution to mix for a few seconds after each addition and measure the conductivity. Record your results in your lab book. Add a total of 90 ml of titrant to the beaker. 6. Rinse your beaker well with distilled water and repeat steps 3-5 for your other two solutions. DATA ANALYSIS Form a group of four and combine all of your data. Enter your data into Cricket Graph and try plotting it. What do your data tell you about the nature of the solutions with which you worked? Here are some hints on how you might proceed. 1. Consult with your lab instructor often, but don't be frustrated if your instructor doesn't answer all of your questions or asks a question in return. You will have a better sense of accomplishment if you work out the details on your own. You will hopefully gain a better grasp of the theory of this experiment if you try many ideas, some of which might be wrong. 2. Try to come up with a way of plotting the data that makes the underlying relationships as clear as possible. The spread sheet will help speed calculations. 3. Chemical theories are often judged on the basis of how well they unify and explain data from different experiments. See if your plots can do that. LABORATORY NOTEBOOK Keep track of all measurements in your lab notebook. Record all observations in your lab notebook. Remember to include units on all numbers that have them. All members of the group should keep track of the various calculations that you tried, in their notebooks. Copy and tape printouts of your calculations into your lab notebook. All members of the group should have a copy of the final graphs taped into their notebooks. REPORT Your report will contain only the answers to the questions below. 1. Give the names of the other three members of your group. All members should attach a copy of the final graphs that your group produced. 2. Which of the five starting solutions would you consider strong electrolytes and which would you consider weak electrolytes? How can you tell? Explain your conclusions. 3. Write the full ionic equation for each titration. 4. Reproduce an enlarged version of the table below in your lab book and turn in a photocopy of the table. 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