Figure 1. Portable Conductivity Meter in Electrolyte Solution.
See Figure 2 for a diagram of the meter
circuit.
This is a guided-inquiry experiment. Working in a team of students, you will devise experimental procedures to answer the following questions. After discussion of experimental plans, you will carry out your plan in lab. Then you will reconvene in the classroom to discuss results.
It is very important to prepare thoroughly for guided-inquiry experiments. The Assignment, background material, and Preparing for Lab questions will expose you to concepts that will prove useful in planning and understanding your experiments.
If an aqueous solution conducts electricity, then it must contain ions. So measuring the conductance of solutions can tell you whether the solutes in the solution are dissociated into ions, and whether chemical reactions in solution are producing or consuming ions.
Any solution, even one containing ions, provides considerable resistance to the flow of current through it. Conductivity is, roughly speaking, the reciprocal of this resistance -- high resistance means low conductivity; low resistance means high conductivity.
Resistance is measured in ohms, so conductance is measured in ohms-1, more commonly called mhos or siemens (S, the official Systeme Internationale (SI) name of the unit). Chemists measure the conductivity of a solution by using the solution to complete an electrical circuit, usually by inserting a pair of electrodes into the circuit, and immersing the electrodes in the solution (Figure 2). The resistance that the solution adds to the circuit is converted to conductivity by a computer chip, and displayed on a meter.
Figure 2. Circuit diagram of conductivity measurement. The solution completes a circuit that includes a battery of known voltage (left) and an ammeter (right). An internal computer chip converts the meter output to conductivity. A portable conductivity meter incorporates all these components into a single device (Figure 1).
Conductivity is roughly proportional to the concentration of ions in solution, but all ions do not conduct equally. Ions that move through solution easily conduct better. For example, small, fast moving ions like hydrogen ion (H+ ) impart greater conductivity to solutions than do bulky ions like bromide ion (Br - ), or heavily hydrated ions like sulfate ion (SO42- ).
Electrolytes are compounds that dissolve in water and dissociate, at least partially, into ions. In solution of elctrolytes, several different species might be present, including intact molecules and dissociated ions. Strong electrolytes dissociate completely into ions. For example, a 1.0 M solution of the ionic strong electrolyte AZ contains 1.0 M A+ ions, 1.0 M Z- ions, and 0.0 M AZ molecules. In other words, the ions are the only species present in a solution of a strong electrolyte. Diluting a 1.0-M solution of AZ with water to 0.50 M reduces the concentration of ions by one half, and thus reduces the conductivity of the solution by one half.
Weak electrolytes dissociate incompletely. For example, a 1.0-M solution of the weak electrolyte BY might contain less than 0.10 M B+ ions, the same molarity of Y- ions (why the same?), and greater than 0.9 M BY molecules. In other words, in a solution of the weak electrolyte BY, the predominant species are BY molecules, while B+ and Y– ions are minor species (present, but not as numerous as the predominant species). The dissociation of BY is an equilibrium process, in which BY molecules are constantly dissociating (forward reaction) and reforming (reverse reaction), at identical rates:
In this reversible reaction, a B+ ion and a Y- ion form whenever a BY molecule happens to dissociate spontaneously (the forward reaction), while a BY molecule forms whenever a B+ ion and a Y- ion happen to collide (the reverse reaction). At equilibrium, the forward and reverse processes occur at the same rate, so BY dissociates and reforms at the same rate, and [BY], [B+], and [Y-] all remain constant.
The probability of spontaneous dissociation of a BY molecule is constant, but the probability of a B+/Y- collision depends on the concentration -- low concentration makes collisions more rare, and makes the reverse reaction slower. This means that diluting a solution of BY slows down the reverse reaction more than it slows down the forward reaction. After dilution, dissociation outpaces collision until the concentrations of the ions rise, their collision rate increases, and the rate of the reverse reaction rises to match that of the forward reaction. As a result, dilution leads to dissociation of additional BY molecules before equilibrium is reached again. So diluting a 1.0-M solution BY to 0.50 M with water reduces the conductivity, but by less than the 50% expected with strong electrolytes.
Think about these principles as you observe conductivity behavior during this experiment.
Download and print the Procedure for this experiment. In guided-inquiry labs, "preparing for lab" means studying the background information (this page), and then answering the Pre-Laboratory Questions of the Procedure. Your instructor will take up your Pre-Laboratory Questions for evaluation, treating them as a pre-lab quiz.
Download the Report Form for this experiment, and follow instructions on the form. Bring it with you to lab. You will write data and observations on the form during lab, as you carry out your work . After lab, you will complete the Report Form by carrying out calculations to give your final lab results. Your instructor might ask you to plot your some of your data with Excel. If you need to use USM computers for this purpose, be sure to allow time to visit the USM computer lab for this work.
If it is not convenient for you to view this web page online, you might want to print it out as well.