CONDUCTOMETRY

Theoretical background

Electrical conduction is the ordered movement of charged particles in an electric field. In

metals the electrons, while in the the melts or aqueous solutions of electrolytes the dissociated ions
carry the electric current. The conductivity of an electrolyte solution can be measured by the
conductometric method. Conductivity is the reciprocal of resistance (I/U), the unit of measurement is
siemens (S). In chemically pure water, the few ions from autoprotolysis conduct electricity only to a
small extent. The conductivity of aqueous electrolyte solutions depends on the intrinsic nature of the
anions and cations they contain, i.e. their mobility and charge, but also on their concentrations. The
actual concentrations of the ions, in turn, are dictated by the extent to which the electrolyte dissociates.
Therefore, the measurement of conductivity is suitable for determining the degree of dissociation.
The conductivity of an electolyte solution can be measured by immersing two metal plates of
the same size into it, and then measuring the current generated upon applying a voltage to them. Using
a given voltage, the current generated depends not only on the chemical nature and amount of ions, but
also on the surface and distance of the metal plates (geometric factors). It is wise to fix the geometric
factors: by convention, metal plates with surface areas of 1 cm2 are fixed at a distance of 1 cm from
each other. In this case, 1 cm 3 of solution is actually tested. The conductivity measured in such a
standardized configuration is the specific conductivity of the solution, denoted by κ (kappa), in S/cm.

Fig. 1. The standardized cell for measuring specific conductivity

When solutions of different concentrations and electrolytes are prepared and tested, the
measured specific conductivity of the more dilute solution is always lower, and that of the more
concentrated solution is always higher, because the ions carry the current, so less or more of them are
located between the metal plates of the measuring cell. However, the extent in the difference of
conductivities between the particular concentrations is distinct for each substance, as the amount of
ions also depends on the degree of dissociation. The divergence which is due to the difference in the
initial concentrations can be eliminated by introducing the concept of equivalent conductivity. The
symbol for the equivalent conductivity is  (capital lambda), its unit of measurement is Scm2/mol. For
substances that dissociate into monovalent ions, the equivalent conductivity is the total conductivity of
a portion of the solution that contains exactly 1 mole of solute. This can theoretically be measured by
connecting exactly as many conductivity measuring cells in parallel as many altogether contain
exactly 1 mole of the solute. The plates of the standardized conductivity cell hold 1 cm 3 of solution,

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so e.g. in the case of a 1 M solution 1000 pieces, with a 0.1 M solution 10,000 pieces of standardized
cells should be connected in parallel to contain 1 mole of solute. In practice, for substances that
dissociate into monovalent ions, the equivalent conductivity can be calculated from the specific
conductivity measured in one standardized cell and from the molarity of the solution:

 = (1000/c)

where c is the molarity of the solution (M). (The 1000/c formula is actually the volume of the portion
of the given solution that contains 1 mole of solute, expressed in cm 3)
Strong electrolytes always dissociate completely regardless of the concentration. In the case of
infinitely dilute solutions, the dissociation is always complete, regardless of the strength of the
electrolyte. The equivalent conductivity of such a solution (Λ ∞) is a constant, the value of which
depends on the chemical nature of the ions in it. In fact, there is no solute in the infinitely dilute
solution, so the conductivity of such a solution cannot be determined directly in practice. The
conductivity of a solution, on the other hand, is an additive property, so thus the conductivity of
multicomponent electrolyte solutions can be interpreted as the sum of the conductivity values of the
individual ions. The individual ions of a weak electrolyte can occasionally be found separately but in
the fully dissociated state in the solutions of other strong, i.e. fully dissociating electrolytes. The
conductivity of a solution containing a reasonable amount of a strong electrolyte is measurable. In this
way, the equivalent conductivity for the fully dissociated state of the weak electrolyte, i.e. theoretically
at the infinite dilution, can be determined by measurement and then calculation.
For example: ∞ acetic acid = (0.01M sodium acetate + 0.01M HCl) ─0,01M NaCl.
For weak electrolytes, the equivalent conductivity is directly proportional to the degree of
dissociation, which in turn increases with dilution. As a result of the dilution, the equivalent
conductivity in weak electrolytes increases up to a limit, so the degree of dissociation ( ) of the
electrolyte solution can be calculated from the equivalent conductivity of a given concentration c (Λc)
and the infinitely dilute solution (Λ∞):

(In the case of strong electrolytes, this correlation is not true. Interestingly, in concentrated
solutions, the equivalent conductivity of strong electrolytes is also lower than in dilute solutions.
However, this is due to the fact that because of the congestion, the distance between the oppositely
charged ions decreases, an attractive electrostatic interaction develops between them, they arrange in
pairs and are less willing to separate and move in opposite directions, so thus their mobility decreases.)

Determination of the degree of dissociation and dissociation

constant of acetic acid by conductometry

The measuring equipment

The conductivity is measureed with a Jenway 3540 conductivity (and pH) meter, to which a
bell-shaped conductivity cell is connected. The conductivity cell consists of a glass bell hiding two
platinum plates in the standardized arrengement. The plates form the conductivity cell with the
electrolyte solution to be measured between them. The small hole in the upper part of the glass bell

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must sink below the electrolyte solution during the measurement in order to ensure the unobstructed
flow of the electrolyte solution.
After switching on the conductometer (red button), use the left and right arrows at the bottom
of the device to select Cond from the menu. In the top left corner of the display, the specific
conductivity () is shown in µS or mS. During the further calculations, pay attention to the unit-to-unit
conversion.

Fig. 2. The Jenway 3540 conductivity meter (left), and the bell-shaped conductivity cell (right, in close-up)

Materials and equipment

 0.01 M CH3COOH solution
 0.05 M CH3COOH solution
 0.1 M CH3COOH solution
 4 beakers

The experiment, evaluation of the results

1. Fill separate beakers with three CH3COOH solutions of different concentrations.
2. Immerse the bell electrode in the first beaker (0.01 M acetic acid) so that the solution covers
the top hole of the glass bell.
3. Read the specific conductivity () of the solution measured from the conductometer.
4. Rinse the bell electrode by sinking-lifting it a couple of times into-from a beaker filled with
distilled water before placing it in a new solution.
5. Similarly measure the conductivity of the other two solutions.
6. Fill in the table below.
7. Determine how the specific and equivalent conductivity and the degree of dissociation
changed with dilution. (For acetic acid  = 344 Scm2/mol).
8. Calculate the dissociation constant of the acetic acid using the
ca2
K a¿
1−a
formula.

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 κ  = (1000/c)
(S/cm) (S cm2/mole)

0.01 M CH3COOH

0.05 M CH3COOH

0.1 M CH3COOH