Prior to doing this lab students may benefit from familiarizing themselves with the equipment and ideas required for its successful completion. To that end I have written a brief exploration activity.

Lab: Conductivity


In this lab students will explore the difference between electrolytes and non-electrolytes. First, they will determine whether a variety of substances are electrolytes by measuring the conductivity of low concentration solutions. Second, they will learn about the idea of solute concentration and how it affects conductivity. Finally, they will explore the differences in conductivity between solutions containing compounds with different numbers of ions per formula unit.


The Nature of Electrolytes

When atoms combine to make compounds they bind to each other by electrical forces of attraction. Positive charges attract negative charges and these strong forces create a bond between atoms. These bonds can take different forms: some are covalent bonds in which atoms share valence electrons. The negatively charged electrons occupy the space between the positively charged nuclei and both nuclei are attracted to the electrons. Other bonds are ionic bonds in which atoms gain electrons to become a negatively charged anion or lose electrons to become a positively charged cation. These ions then attract each other directly with no electrons between them.

Salt.Dissolving (9K)
An ionic compound in the act of dissolving in water. See a full explanation with animations at this page:

Chemical bonds are formed and broken during chemical reactions that change one substance into another. Chemical bonds may also be temporarily broken by dissolving a compound in water. Some compounds have bonds that can be broken to create two or more freely moving ions when the compound dissolves in water. This is due to the fact that water is a strongly polar molecule. It has two oppositely charged poles that strongly attract other atoms and molecules. In the illustration at right water molecules are shown dissolving ions from a solid crystal. The water molecules’ positively charged poles attract negative ions away from the crystal and then surround them as the ions then become mobile and move around in the solution. The negative poles of the water molecules do the same for the positive ions. The ions remain surrounded by water molecules just as a famous pop star is always surrounded by an entourage. Chemical compounds which break up into ions when dissolved in water are called electrolytes. All strong electrolytes are ionic compounds made up of positively charged cations and negatively charged anions. There are other compounds which dissolve well in water but which do not break into separate ions. Such compounds are called non-electrolytes. All non-electrolytes are molecular compounds and they do not make water conductive when they dissolve because they are not made of ions: their atoms are all covalently bonded.

Concentration of Solutions

Solutions are homogeneous mixtures of two or more substances. For liquid solutions the solvent is often water. Solvents are materials in which other substances dissolve. Solids, liquids and gases may all dissolve in water and any of them may be called a solute. A solute is any substance that dissolves in a solvent. Solutions may be made to have a variety of different strengths, or concentrations. The concentration

Dissolved.Ions.diff.concentrations (8K)
The image on the left shows dissolved ions at low concentration.
The image on the right shows dissolved ions at high concentration.
of a solution is a quantitative measure of how much solute is dissolved in the solvent. Often concentration is measured in moles per liter. A mole is a unit related to the number of molecules of a substance and the liter is a common unit of volume.

Concentration is higher for solutions containing more of the solute and lower when less material has been dissolved. In the illustration at right there are two boxes showing different concentrations. On the left there are very few particles of the solute and the image represents a solution with a low concentration. On the right there are many more solute particles in a higher-concentration solution.


One interesting consequence of the fact that some compounds produce separate, mobile cations and anions when dissolved in water is that the solution conducts electricity. Imagine a complete circuit containing a light bulb and a battery. The light bulb lights because of the power supplied by the battery. If a wire is cut then the bulb goes out. But if the two cut ends of the wire are placed in an electrolyte solution of high enough concentration the bulb will light again. Very often it is possible to demonstrate this in a classroom. There is also a high-quality simulation available from the PhET project at the University of Colorado. It is called Sugar and Salt Solutions and is available here:

Electrons in a metal are mobile and can move among the atoms of the metal like water in a pool full of pebbles. When a battery is used to make an electrical current it causes the electrons to move in an organized way, all in one direction, turning the pool into a stream. This motion of the electrons in a piece of metal wire is what we normally think of as an electrical current. The electrical current in an electrolyte solution has a slightly different mechanism. The ions themselves move and carry electrical charge from one wire to the other. Electrons do not leave the metal and travel through the water to the other wire. Instead, cations move toward the negatively charged wire and anions move toward the positively charged wire. Fundamentally, an electrical current is only the movement of charged particles and it doesn’t matter whether the particles are electrons or ions.

The conductivity probes you will use in this lab activity measure how well a solution conducts electricity. It gives a low reading for a low concentration of ions and a higher reading for higher concentrations of ions. In effect, the reading produced by the probe is a measure of the number of ions per liter. The number of ions per liter depends on two things. First, it depends on how much solid is dissolved. The more material that is present in solution, the more conductive the solution will be. Second, it depends on the chemical formula of the dissolved material. Some compounds have more ions in their chemical formulas than others. Consider the following compounds and the number of ions each one makes when they dissolve:

sodium chloride       NaCl —> Na+ + Cl
calcium chloride      CaCl2 —> Ca2+ + Cl + Cl
sodium phosphate   Na3PO4 —> Na+ + Na+ + Na+ + PO43–

Each equation shows how many ions are produced from each formula when they dissolve in water. Sodium chloride produces two ions, calcium chloride produces three ions, and sodium phosphate produces four ions. Even when these salts are made into solutions with the same concentration (measured as whole formula units per liter) they produce different levels of conductivity. For example, a solution of sodium phosphate will be more conductive than a solution of sodium chloride with the same concentration because it has twice as many ions per formula unit.

Pre-Lab Questions

Answer the questions below in your lab notebook.
  1. Define the terms solute, solvent, and solution in your own words.
  2. In your own words, describe the process by which an ionic solute dissolves in water to make an electrolyte solution.
  3. What is an electrolyte and how is it different from a non-electrolyte?
  4. In your own words, describe how electrical current is moved through an electrolyte solution and how this is different from how electrical current moves through a metal wire.
  5. Do electrolytes carry electricity or produce it? Justify your answer by thinking about the example of the light bulb and the battery from the introduction.
  6. Write chemical equations showing how MgBr2 and AlCl3 separate into ions when they dissolve in water. Use the examples given in the text inside a box as a guide.
  7. How many ions does one formula unit of MgBr2 break up into when it dissolves? How many ions per formula unit does AlCl3 make?
  8. Based on your knowledge of ionic compounds, which of the following two compounds will produce a solution with higher conductivity if the same number of formula units are dissolved? AlCl3 or MgBr2? Justify your answer.
  9. How does increasing the concentration of a solution of the same electrolyte change the conductivity?

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The chemicals used in this activity are all of low toxicity and are used in dilute solutions. Even so, it is best to wear standard laboratory safety equipment such as splash goggles, long pants, and nitrile gloves, if desired. Tie back long hair and do not wear dangling jewelry in the lab.

Part I: Electrolytes vs. Non-electrolytes


In this part of the lab you will measure the conductivity of six different substances: distilled water, hydrogen peroxide (H2O2) solution, sucrose (C12H22O11) solution, sodium chloride (NaCl) solution, calcium chloride (CaCl2) solution, and sodium phosphate (Na3PO4) solution.

  1. Wear your safety goggles while working in the lab.
  2. On a separate piece of paper create a data table to record one number for each of the six substances (incl. water) you are going to test in this part of the lab.
  3. Set up the Conductivity Probe using the ring stand and clamp. Do not clamp the probe too tightly and do not submerge the probe completely into water at any time.
  4. Set the range selector on the Conductivity Probe to “0 - 200 μS/cm”.
  5. Connect the probe to the computer and start the Logger Lite software.
  6. The software will immediately recognize the probe and produce a data collection screen.
  7. Go to the “Experiment” menu, select Change Units, select the probe and choose mg/L. This is milligrams per liter of equivalence to NaCl.
  8. Obtain and label five 100-mL beakers with the formulas of the substances you will be testing: NaCl, CaCl2, Na3PO4, C12H22O11, and H2O2. You will use these beakers throughout this lab without emptying them until it is time to clean up.
  9. Fill each beaker with 50 mL of distilled water by measuring with a graduated cylinder.
  10. Test the conductivity of one of the beakers containing distilled water by submerging the electrode at the bottom completely. Water alone will be the first material you test. Hold the beaker up to the probe and submerge the end of the probe in the water. Write down the conductivity value you see on your screen. You will not record this information digitally.
  11. Now add just one drop of each solution your teacher has provided for you to each labeled beaker.
  12. Stir the NaCl solution with the glass rod. Rinse the rod and dry it before stirring the next solution. Stir each one with the stirring rod, rinsing and drying in between each one.
  13. Use the Conductivity Probe to measure the conductivity of the NaCl solution. Record the value in a data table you write down. Don’t forget to include distilled water as one of the substances you tested.
  14. Rinse the end of the probe carefully using the wash bottle full of distilled water. This prevents the probe from giving inaccurate readings when you test the next solution.
  15. Now use the probe to measure the conductivity of each of the other solutions. Always be sure to rinse the probe between tests!

Part II: Electrolyte Conductivity vs. Concentration


In this part of the lab you will collect data to show the proportion between concentration and conductivity for three strong electrolytes. You will not be testing the conductivity vs. concentration for the non-electrolytes.

  1. On the “Experiment” menu, click “Data Collection…”. Choose “Events with Entry”.
  2. For “Name” enter “Number of Drops” and for “Units” enter “Drops”. Click OK
  3. Right-click in the middle of your graph and choose Graph Options. Go to the Axes tab. Set the maximum value for the x-axis to 10 so your data will fill your graph.
  4. Change the switch on the Conductivity Probe so that it points at “0 – 2,000”. The software will ask you to confirm the change. Click the button marked “Use Sensor Setting” in the dialog box that pops up.
  5. Click the “Collect” button to start data collection.
  6. Raise the beaker containing 50 mL distilled water and one drop of NaCl solution up over the end of the probe. Click the Keep button and enter 1 for the volume in drops. Do not rinse off the conductivity probe between tests.
  7. Add one drop of NaCl solution and stir it using your clean stirring rod.
  8. Test the conductivity of the solution by holding it so the end of the probe is submerged. Swirl it gently, without spilling and watch the conductivity value for about 5 seconds. Once the reading reaches a steady value, click “Keep” and enter 2. Do not rinse off the conductivity probe between tests.
  9. Repeat this process until you have reached a total of 8 drops.
  10. Stop data collection by clicking the “Stop” button.
  11. Next you will use the software to analyze your result by producing a linear regression. This will calculate the equation of the line and plot a best-fit line for your data.
    1. On the “Analyze” menu choose “Linear Fit”.
    2. When it plots the line and gives you the equation it will be in the form y = mx + b. The y variable stands for conductivity, the x value stands for volume in drops.
    3. Record the value of the slope in a data table you create yourself on a separate piece of paper.
    4. Click the File Cabinet icon to “Store” the data from the first run.
  12. Rinse the electrode with the wash bottle.
  13. Repeat the data collection procedure with the CaCl2 solution.
  14. Remember to store the data from Run 2 using the File Cabinet icon.
  15. Repeat the data collection process with the Na3PO4 solution.
  16. As instructed by your teacher, print this graph and/or export the data for further analysis and formatting in a spreadsheet program. In your report you must include a well-formatted graph with all three runs properly labeled and showing the slope of each line.

Post-lab Questions

Answer the following questions in a typed document on a separate paper. As part of your report you must include the answers to the Pre-lab questions, your formatted and labeled graph, and the answers to the questions below. This lab does not require a formal lab report.

  1. Create a data table to compare the conductivity of each different solution with only one drop of the stock solution. It must have 2 columns: “Name and Formula” and “Conductivity Measurement”. Use borders on the table. Include it as the answer to this question in your typed document.
  2. Why do some compounds show essentially zero conductivity? (That is, the same conductivity or nearly, as distilled water).
  3. Which of the substances you tested in part one were electrolytes? Which of them were non-electrolytes? Use your data to justify your answers and recall that distilled water should have essentially zero conductivity. If your distilled water had a reading slightly above zero then take this value as equivalent to zero conductivity.
  4. Why are some electrolytes more conductive than others despite the fact that you added the same amount of each one? Use the chemical formulas of each compound to explain why the conductivity values follow the pattern that you see.
  5. Paste the graph you created in Logger Lite into the answer for this question. The graph must have Drops on the x-axis and Conductivity on the y-axis. Each line plotted must have a clear label showing the equation of the line.
  6. Describe what happens to conductivity as concentration increases. What kind of proportion (direct or inverse) exists between the conductivity of a solution and the concentration of ions in the solution?
  7. Compare the slopes of the three lines and write them in order from smallest to largest, identifying each one with the chemical formula of the salt. Use the chemical formulas of the electrolyte compounds to explain the difference in the slopes even though all three were generated starting with solutions with the same concentration.
  8. Write equations for when the following salts dissolve in water: Na2SO4, KBr, and FeCl3. Do not look for an answer online. The point is whether you can recognize the ions that compose an ionic compound and write an equation showing how these separate when the compound dissolves in water.
  9. Which of the salts in the previous problem would have the steepest slope if it were used in Part II of this lab? Which would have the smallest slope? Justify your answers.

Prior to doing this lab students may benefit from familiarizing themselves with the equipment and ideas required for its successful completion. To that end I have written a brief exploration activity.
Last updated: Mar 31, 2022       Home