Forces and Bonding: Conductivity

• To explore the behavior of substances dissolved in water and other solvents
• To distinguish strong electrolytes, weak electrolytes, and nonelectrolytes
• To explore the nature of the bonding within substances based on their ability to conduct electricity
• To formulate generalizations relating structure to solution behavior

Pure substances can be classified according to physical state (solid, liquid, or gas at a particular temperature); physical properties (metal, semimetal, nonmetal); bonding (ionic, covalent, metallic); conductivity when dissolved in water (electrolyte, nonelectrolyte); and in numerous other ways. The conductivity (or lack thereof) of a substance or a solution of the substance provides a lot of information about the nanoscale nature of the substance when pure or dissolved. Conductivity by a pure substance or a solution of the substance implies the presence of either

By measuring conductivity of solutions, we can deduce whether a substance dissolves as uncharged molecules or as ions, and to what extent. In this experiment, you will build a conductivity tester and use it to explore a number of substances in several solvents. This will provide a phenomenological basis for describing and beginning to understand the forces that operate between fundamental structural units of substances at the molecular level.

In this experiment, it will be necessary for you to prepare approximately 0.1 molar (0.1 M) solutions of a number of substance in water or another solvent. A solution is a homogeneous mixture in which one or more substances, called solutes, are said to be dissolved in another substance, called the solvent. In most cases, the solvent is a liquid. Solutes may be solids, liquids, or gases. A solution is described by stating its concentration, the amount of solute contained in a unit amount of either solvent or solution. The most frequently used concentration unit is molarity, M. A solution that is 0.1 molar, or 0.1 M, contains 0.1 mole of substance per liter of solution, where 1 mole of a substance is a number of grams of the substance numerically equal to its mass on the atomic mass scale. Thus, for example, one mole of sodium chloride, NaCl, is 58.443 g of NaCl. To prepare 1 L of a 0.1 M aqueous (water) solution of some substance X, you would weigh a mass of substance X equivalent to 0.1 mole, and dissolve that mass in enough water to give a total volume of 1 L. Of course, 1 L is a large volume of solution. In this experiment, you will be testing small volumes on the order of 1 mL, so you would not want to prepare a liter of solution, only to discard 999 mL of it. Instead, you will prepare only the volume of each solution that you need--about 1 mL. To prepare 1 mL (0.001 L) of a 0.1 molar solution of a substance requires that you weigh 0.0001 mole (or 0.1 mmole) of the substance, and dissolve it in enough water to give 1 mL. That this produces a 0.1 M solution can be seen by taking the ratio of the moles dissolved to the total volume:

Thus a 0.1 M solution of a substance contains 0.1 mole of the substance per L of solution, or 0.1 mmole of substance per mL of solution.

This small volume of solution will allow you to carry out the necessary tests without waste of reagents.

In general, if you want to prepare "a" liters of a solution in which the solute has concentration, "x", you will weigh out "x*a" moles of solute and dissolve it in enough water to give a total volume of "a" liters. For example, to prepare 2.5 mL of a 0.10 M solution of NaCl in water, weigh 0.10*0.0025 moles NaCl and dissolve it in enough water to give a total volume of 0.0025 L, or 2.5 mL.

Finally, suppose you want to prepare a certain volume of a solution of a particular solute by diluting a concentrated stock solution of that solute. For example, suppose you need 6.0 mL of 4.0 M HNO₃, which you must prepare from concentrated HNO₃, which is 16.0 M. The reasoning goes like this. The desired solution is 4 x less concentrated than the stock solution, which tells us that the stock solution must be diluted by a factor of 4. To end up with 6.0 mL of stock solution, then, requires that we start with 6/4 = 1.5 mL of stock solution, and add it to enough water (approx. 4.5 mL) to give a total volume of 6.0 mL. Dilution is an often-used procedure in the lab, so please be sure you understand the thinking here.

For your reference, the molarities of some common laboratory reagents are given here:

In order for you to properly classify your substances you must know a few operational definitions. First, a substance acts as an electrolyte in a particular solvent if a solution of the substance in the solvent conducts electricity. Conductivity results because the electrolyte produces ions in solution. A substance producing a relatively high conductivity is called a strong electrolyte; one producing a low conductivity, a weak electrolyte; and one producing no conductivity, a nonelectrolyte. It is possible for a particular substance to be a strong electrolyte in one solvent, and a weak or nonelectrolyte in another.

Focus Questions

While you are performing the experiment, we would like you to consider the following focus questions. When you have completed the experiment, provide answers for them in your notebook.

1. Are all ionic substances aqueous electrolytes?
2. Are all aqueous electrolytes ionic compounds?
3. What is the effect of the solvent on conductivity?
4. Use your results to classify each substance as strong electrolyte, weak electrolyte, or nonelectrolyte in each solvent tested. Classify each substance as metallic, ionic, or covalent. Classify each substance as acid, base, salt, or other (acid = H⁺ + anion; base often = cation + OH-; salt = cation other than H⁺ + anion other than OH-).
5. Is it possible for a covalent substance to be an electrolyte (strong or weak)? Explain, and give examples.
6. Is it possible for an ionic substance to be a nonelectrolyte? Explain, and give examples.

• assorted glassware (small beakers, graduates)
• Pasteur pipets
• well plate
• 9-volt battery
• 9-volt battery clip
• small LED
• 330 ohm, 0.25 watt resistor
• perfboard (4-hole by 10-hole)
• adhesive tape
• solvents
• assorted pure substances in solid, liquid, or concentrated solution form

You will work with solutions of acids, bases, and salts, some of which are toxic. Avoid skin contact. Clean up all spills immediately. In case of skin contact, flush with plenty of water.

Record your observations in your notebook.

Assembling the Conductivity Meter. The conductivity meter consists of a simple series circuit, assembled as follows:

1. Insert the LED leads into the perfboard. The long lead goes through a corner hole and the short one goes into the adjacent hole along the length of the board.
2. Push the black lead of the battery clip through the LED short-lead hole, fully insert the LED, and twist the short LED and clip leads together. Bend parallel to the board and tape down.
3. Bend the long LED lead 90 degrees so that it extends beyond the short edge of the board.
4. Pass the brown-stripe lead of the resistor through the third hole along the short edge from the long LED lead. Insert the orange-stripe resistor lead through the third hole along the board length. Bend the edge lead 90, parallel to the LED long lead.
5. Force the red lead of the battery clip through alongside the second resistor wire, and twist the two leads together. Bend parallel to the board and tape down
6. Attach the battery to the clip. The assembled conductivity meter should look as shown here.
7. Test the meter by shorting across the two protruding leads (long lead of LED and resistor wire) with a piece of conducting metal (a penny, or a strip of aluminum foil). The LED should light. This brightness represents "full" conductivity.

Your conductivity meter is somewhat crude, in that it will not allow you to measure conductivity numerically. Instead, you will classify the ability of a substance to conduct as follows:

Substance and Solution Testing. After you have assembled and tested your conductivity meter, the instructor will assign you several solvents (one of them water) and a number of pure substances to investigate. Consider the formulas of the substances, and predict for each whether or not an aqueous solution will conduct electricity. Substantiate your predictions with equations. Discuss your predictions with the instructor before carrying out the tests.

Make two small glass stirring rods by sealing off the ends of melting point capillary tubes in a burner flame. Get help from the instructor if necessary.

You should carry out conductivity tests for

You should plan to prepare only about 1 mL of each solution. You may do this directly in the wells of the well plate by weighing the appropriate amount of substance and transferring it to a well. Test the conductivity of the solid by probing the solid with the conductivity meter leads, then add about 1 mL of the specified solvent by counting drops from a Pasteur pipet (approx. 30-40 d/mL). Test the solution conductivities by submerging the conductivity meter leads. Dip the tester leads in distilled water and wipe dry between tests.

Use your meter to test each of the "household products" available in the lab, and record the results.

Clean-up. When you have finished all of your work:

• Clean all glassware by recommended procedures, shake off excess water, and place in drying oven.
• Return all borrowed equipment to the instructor before leaving lab.
• Clean up your work area before leaving lab.

Meet for post-lab discussion.

Using a Pasteur pipet, dispose of solutions of NaCl, KCl, and CaCl₂ down the sink. Transfer all solutions containing transition-metal, Pb, Sn, and Hg solutions to the container marked Metal Waste.

1. DA Katz, C Willis, J. Chem. Ed. 1994, 71, 330-332.