Electrical Conductivity of Aqueous Solutions

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1 Electrical Conductivity of Aqueous Solutions PRE-LAB ASSIGNMENT: Reading: Chapter in Brown, LeMay, Bursten & Murphy.. Using Table in this handout, determine which solution has a higher conductivity,. M HCl or. M NaCl. Explain what creates the higher conductivity. 2. Write the balanced complete ionic equation for the titration of aq. NaOH with aq. HCl. INTRODUCTION: In this lab you will explore the nature of aqueous solutions by investigating the relationship between conductivity and strong and weak electrolytes. To do this, 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 another "discovery" exercise to get you thinking about how the experimental data you have collected relates to the concepts you are learning in class. You will work in small groups to decide how best to work up the data. Plotting your data will allow you to visually follow the chemical reactions that are occurring and deduce what ions that are present. What is conductivity? Conductivity is a measure of the concentration of ions in solution. By completing the circuit shown in Figure, we can measure the conductivity of the solution in the beaker. The conductivity is proportional to the current that flows between the electrodes. 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 that can be carried between electrodes and will increase the conductivity. The units microsiemens/cm (us/cm) and millisiemens/cm (ms/cm) are most commonly used to describe the conductivity of aqueous solutions. Figure. Schematic of a simple conductivity measurement system. * The Siemen was formerly called mho (pronounced "mo"), which was derived as a unit of conductivity by reversing the letters in "ohm," the unit of resistance.

2 -2- CH 4 Lab Exp #4 Another factor in conductivity measurements is that not all ions carry charge (conduct electricity) equally. H + and OH are unique and move through solution very rapidly and are very good charge carriers. 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. Table is a table of molar conductivity for the ions in this exercise. The molar conductivity is the conductivity of a solution for the ion containing one mole of charge per liter. Note that the molar conductivity of H + ions is 5-7 times the conductivity of other small cations. The molar conductivity of OH - is 3-5 times the conductivity of other small anions. To calculate the conductivity of a solution you simply multiply the concentration of each ion in solution by its molar conductivity and charge then add these values for all ions in solution. Table. Molar Conductivity of Selected Ions Ion Molar Conductivity (S L mol - cm - ) Cations: H Na +.5 NH K Anions: OH.986 Cl.7635 CH3COO *.49 Br.78 *CH 3 COO is the acetate ion, and it is often written as C 2 H 3 O 2. How can we use conductivity to study solutions? Substances that completely dissociate into ions (strong electrolytes) produce solutions with high conductivity. Substances that partially dissociate (weak electrolytes) will produce solutions with low conductivity. Substances that dissolve but do not dissociate produce solutions that have the very low conductivity of pure water. 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 an acid base titration of a solution of KOH with HBr. The reaction is represented in the following: balanced chemical equation: KOH(aq) + HBr(aq) KBr(aq) + H2O complete ionic equation : or the net ionic equation: K + + OH + H + + Br - K + + Br + H2O OH + H + H2O During the course of the titration OH - is consumed, Figure 2. The equivalence point is reached when the number of moles of H + added is equal to the number of moles of OH - in the starting solution. After the equivalence point, the moles of H + increase, because they are no longer consumed by reaction with OH -. For the spectator ions, the number of moles of K + remains constant and the number of moles of Br - increases as the titrant is added. How will these ion concentrations be reflected in the conductivity of the solution? Initially the solution conductivity is high, because OH - has a large molar conductivity. However, as OH - is neutralized the conductivity falls. After all the OH - is neutralized the H + concentration increases and the conductivity again rises rapidly. However, remember that the conductivity includes contributions from all ions in solution, not just H + and OH -.

3 -3- CH 4 Lab Exp #4 mol OH - (x -3 ) mol HBr (x -3 ) mol H + (x -3 ) mol HBr (x -3 ) mol K + (x -3 ) mol HBr (x -3 ) 2 mol Br - (x -3 ).5.5 mol HBr (x -3 ) Figure 2. Titration of 5 ml of.667 M KOH with.2 M HBr. Note the expanded scale for Br -. PROCEDURE: For 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 (NaC 2 H3O 2 ) with HCl. At the end of class you will work in groups of four to analyze your data so that each group has all five of the titrations plotted. You will use this conductivity data to predict the ionic composition of your test solutions before, during, and after the titrations.

6 -6- CH 4 Lab Exp #4 Supplemental Report Sheet (In each box, fill in all the chemical species present) Starting Solution/ Titrant NaOH / HCl ml titrant added 5 ml titrant added 8 ml titrant added NaC 2 H 3 O 2 / HCl HC 2 H 3 O 2 / NaOH NH 3 / HCl NH 4 Cl / NaOH

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