CONDUCTANCE OF STRONG AND WEAK ELECTROLYTES

In this experiment the measurement of the electrical conductivity for ionic solutions with decreasing concentration is extrapolated to zero concentrations to determine the limiting conductivity of each type of solution at zero concentration.   The conductivity of three strong electrolytes will be used to determine the limiting conductivity of a weak electrolyte, acetic acid.  Then the ratio of the experimental conductivity to the limiting conductivity will be used to determine the equilibrium constant of acetic acid.  The conductivity is to be measured by two means, a conductivity probe and a conductivity meter.

 

PROCEDURE

Apparatus

Conductance cells (S); Jenway conductivity meter (S); Vernier conductivity probe (S); 50 ml beakers(S); pipettes (D); volumetric flasks (D); 25 ^oC thermostat (L).

 

Chemicals

100 ml 0.0200 N KCl (P); 100 ml 1.00 N HCl (L); 100 ml each of 0.1, 0.05, 0.025, 0.0100 and 0.00500 N HCl (P); 100 ml 1 N acetic acid (L); 100 ml each of 0.05, 0.025, 0.0125, 0.00625, and 0.00312 N acetic acid solutions (P); 100 ml 1 N NaCl (P); 100 ml each of 0.1, 0.050, 0.025, 0.0100, and 0.0050 N NaCl (P); 100 ml 1 N NaC₂H₃O₂ (P); 100 ml each of 0.1, 0.050, 0.025, 0.0100, and 0.0050 N NaC₂H₃O₂ (P); distilled H₂O (L).

 

It is very important to prepare a solution of KCl that is exactly 0.0200 N. The other solution can be approximate, but their value must be known exactly.  Calculate and weigh out the amount of KCl required into a 100 ml volumetric flask and add deionized water to the 100 ml mark.  Prepare an approximately 100ml of 1.00 N HC₂H₃O₂ using glacial acetic acid.  Then make the solutions of HCl and acetic acid by diluting the 1.00 N HCl provided, and the 100 ml of 1.00N acetic acid just prepared. Use pipettes and volumetric flasks for all dilutions!

 

Example Dilutions

10 ml 1.00 N HCl diluted to 100 ml yields 0.10 N HCl

25 ml 1.00 N HCl diluted to 500 ml yields 0.05 N HCl

25 ml 0.10 N HCl diluted to 100 ml yields 0.025 N HCl

20 ml 0.05 N HCl diluted to 100 ml yields 0.0100 N HCl

 

Prepare the 100 ml of 1.00 N NaCl and 1.00 N NaC₂H₃O₂ solutions by weighing out the necessary amount of the predried salts. The other concentrations of the salt solutions are prepared by appropriate dilutions of these 1.00 N solutions.  The instruction manual for the Jenway conductance meter and the Vernier conductance probe is reproduced in the manual section.

 

Measure the conductance of the KCl solution in order to determine the cell constant using both the conductivity probe and bridge. For the instruments used in this laboratory it should be near 1, if it is not, contact the instructor. Save the KCl solution for check measurements on the cell constant! Measure the conductance of the distilled H₂O, then measure the conductance of each solution of HCl. Proceed from one solution to the next by rinsing the cell with a few ml of the solution to be tested.   Immerse the conductance cell in a 50 ml beaker 2/3 filled with the appropriate solution.  There must be enough solution to ensure that the electrodes are completely covered. Measurements are done on the most dilute solutions first in order to minimize errors due to imperfect rinsing and absorption of ions on the electrodes. Now repeat the series of conductance measurements with the acetic acid, NaCl and NaC₂H₃O₂ solutions. When you have finished for the day, rinse the cell with distilled water and store in distilled water (as it was received) and return bridge and cell to stockroom.

 

Calibration

The Conductivity Probe can be easily calibrated at two known levels, using any of the Vernier data-collection programs.

• Select the conductivity range setting on the probe box: low = 0 to 200 mS, medium = 0 to 2000, mS, and high - 0 to 20,000, mS. Note: if you are not sure of which setting to use, you may first want to load a stored Vernier calibration for one or more of the settings to determine an approximate value for the solution to be sampled.

 

• Zero Calibration Point: Simply perform this calibration point with the probe out of any liquid or solution (e.g., in the air). A very small voltage reading will be displayed on the computer.

• Standard Solution Calibration Point: Place the Conductivity Probe into a standard solution (solution of known concentration), such as the sodium chloride standard that is supplied with your probe. Be sure the entire elongated hole with the electrode surfaces is submerged in the solution. Wait for the displayed voltage to stabilize. Enter the value of the standard solution 2768 m ohm^-1 cm^-1 at 25 ^oC. For further information on preparing and interpreting standard solutions, see subsequent sections on calibration.

 

This method of calibrating is easy enough that we recommend that you perform a calibration whenever you use the probe. As an alternative, you can save a calibration performed using a conductivity range setting (range setting or standard value in the calibration name), and reload it at a later date.

 

For even better results, the two-point calibration can be performed using two standard solutions that bracket the expected range of conductivity or concentration values you will be testing. For examples if you expect to measure conductivity in the range of 600 mg/L to 1000 mg/L (TDS), you may want to use a standard solution that is 500 mg/L for one calibration point and another standard that is 1000 mg/L for the second calibration point.

 

Taking Measurements using the Conductivity Probe

Once the Conductivity Probe has been calibrated, you are ready to take readings

• Rinse the tip of the Conductivity Probe with distilled water. Optional: Blot the inside of the electrode cell dry only if you are concerned about water droplets diluting or contaminating the sample to be tested.

• Insert the tip of the probe into the sample to be tested. Important: Be sure the electrode surfaces in the elongated cell are completely submerged in the liquid.

• While gently swirling the probe, wait for the reading on your computer screen to stabilize. This should take no more than 5 to 10 seconds.

• Rinse the end of the probe with distilled water before taking another measurement.

• If you are taking readings at temperatures below 15°C or above 30°C, allow more time for the temperature compensation to adjust and provide a stable conductivity reading.

• Important: Do not place the electrode in viscous, organic liquids, such as heavy oils, glycerin (glycerol), or ethylene glycol. Do not place the probe in acetone or non-polar solvents such as pentane or hexane.

 

Instrumentation

Conductance cell: The conductance cell for the Jenway conductance meter consists of two parallel electrodes of area 1cm² separated by a distance of 1 cm as seen in Figure 1.  This configuration will have a cell constant of 1 making the conductivity and the conductance of the cell equal in numeric value.  This cell, when attached to the meter, can be used to measure conductivity as seen in Figure 2.

 

 

 

 

 

 

 

 

 

 

 

 

 

Figure 1.  Picture of Jenway conductance cell.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Figure 2.  Picture of Jenway conductance cell attached to the Jenway meter.

 

The electrodes of the conductance cell must have an adherent coating of platinum black and should be immersed in distilled water whenever the cell is not use. If the electrodes are allowed to dry out, it is difficult to rinse out electrolytes from them, and it is advisable to dissolve off the coating with aqua regia (under the hood) and plate out a fresh deposit. The electrodes and cell are rinsed out thoroughly first with distilled water and then with conductance water, which is especially pure water prepared by multiple distillations. The conductance cells must be handled with great care; the electrodes must not be touched, and they must not be moved respect to each other during the course of an experiment.

Controls

                Switches the instrument on and to Standby.

 

These keys are used to change a parameter.

               The keys are used to adjust the cell constant or temperature coefficient when displayed.

 

  These keys are used to move the cursor between the menu options.

 

                  This key is used to select the displayed menu option.

 

The cell is filled with conductance water and inspected to make sure there are no air bubbles at the electrodes, and its resistance is measured; it is then rinsed and refilled, and the resistance is measured again.  If an accurate temperature is desired, the cell should be placed in a thermostat bath at the desired temperature. This process is repeated until the resistance has become essentially constant, showing that contaminating electrolytes in the cell have been rinsed out. The cell resistance will not become absolutely constant because the conductance water is very pure and traces of electrolytes insignificant in the later measurements will produce noticeable fluctuations. The specific conductance of the water used should be about 5 ohm^-1 cm^-1 or less, corresponding to a resistance of 200,000 ohms or more in a cell of unit cell constant. At these high resistances an accuracy of more than two significant figures is difficult to obtain without special precautions. For the other solutions studied, a precision of the order of a few tenths of a percent should be obtained in the resistance measurements.

 

Method for general purpose Conductivity Standard

Accurately weigh out 0.746 grams of dried A.R. grade Potassium Chloride (KCI) and dissolve in 1 litre of good quality water. This produces a 0.01N solution with a conductivity of 1413mS at 25°C.

 

·        CAL - pressing ENT with the cursor beneath this option illuminates the CAL annunciator momentarily and then calibrates the instrument to the nearest standard solution (10mS, 1413mS, 12.88mS or 0mS). “Err” is displayed if a cell constant outside the allowed range is calculated and the calibration aborted.

·        Secondary display - provides direct readout of solution temperature in °C or °F (selectable by DIP switch setting). Also provides readout of cell constant (K) and temperature coefficient (%).

·        Endpoint detection symbol - this is displayed once a stable reading is detected, and is maintained until the input changes.

·        Selected mode indicator.

·        Cursor - used to select/indicate required mode.

·        Selected unit of temperature(°C/°F), cell constant(K) and temperature coefficients(%).

·        Measurement unit which is being used, mS or mS.

 

When the cell is clean, as shown by a reasonably constant high resistance with conductance water, it is rinsed two or three times with 0.020 N potassium chloride solution (2768  m ohm^-1 cm^-1 at 25 ^oC⁵) and the resistance is then determined with this solution filling the cell. Additional measurements are made on fresh samples of this solution until successive determinations agree closely. The purpose of these measurements is to provide data to adjust the cell constant.  The cell is emptied and rinsed with the next solution for which the conductance is to be measured. It is advisable to make check determinations on each solution to make sure that the cell was thoroughly rinsed.

 

THEORY[1-4] If a wire is placed between the two terminals of a battery, current will flow. The amount of current which will flow depends on the resistance of the wire according to Ohm’s law, V = I R, where V is the voltage, I the current, and R the resistance of the wire. This resistance will depend on the composition of the wire and on its geometry. For example, two copper wires of the same type and geometry will have the same resistance. However if a copper wire with double the cross sectional area and the same length is used, the resistance of that wire will decrease by one half. If the length, L, of a wire is doubled while its cross sectional area, A, remains the same, the resistance of that wire will double. It is possible to characterize a material, such as copper, with a measure of resistance that is an intrinsic property of the material called its resistivity, r. This resistivity can depend on other properties such as its purity and crystalline state, but is an intrinsic property of the material, as is the melting temperature. However, resistance to the passage of current depends on the geometry of the object formed from a material.  The relationship between the resistance of a material and the geometric independent resistivity is:

 

                          r = R A/L                                                                                  (1)

It is often convenient to discuss the conductivity, k, instead of the resistivity. Conductivity is the inverse of the resistivity.

 

                            k= 1/r                                                                                     (2)

 

The SI unit of resistance is the ohm (W ). The SI unit of conductance is the siemen (S) or W ^-1, though many people still refer to this unit as simply ohm written backwards, or mho. Conductivity is then measured in siemens m^-1 (or Scm^-1 is more convenient)

 

If electrodes are immersed in a solution of say sodium chloride in water, the conductivity can be measured using a conductivity meter.  This conductivity will depend on the type and concentration of ions present, the solvent, the area of the electrode, and the distance between the electrodes.  If the solvent is pure water, there are few charged species (ions) present, K_w =[OH^-][H⁺]=10^-14, the conductivity measured will be low. The addition of ions such as in a 0.10 N solution of NaCl (L = 106.8 W ^-1 cm² mol^-1) will increase the conductivity of the solution to106.8 W^-1cm²mol^-1x 0.01mol/1000 cm³ = 10.68 x 10^-6W^-1 cm^-1 or 1068 x 10^-6W^-1 m^-1.

 

The measurement of the conductivity of an electrolytic solution begins with the measurement of the resistance for that solution using a conductivity cell. The resistance measurements must be accomplished using alternating current (AC) to avoid changes in concentrations and product build up on the electrodes which in turn would effect the measurement. Usually, Platinum electrodes, coated with colloidal Pt to absorb any gas products, are used to in order to reverse changes occurring during phases with reverse polarity.  After the resistance is determined, the resistivity is determined using equation (1).  Since the geometry of the electrodes can change during their use through physically deforming, the ratio L/A could change, therefore it is necessary to determine the value of L/A by using a solution with known conductivity to calibrate the conductivity cell by determining a cell constant.

 

Specific conductance s is the reciprocal of the specific resistance. If a cell could be constructed with electrodes of exactly 1 cm² area exactly 1 cm apart, the reciprocal of the cell resistance in reciprocal ohms would be numerically equal to the specific conductance in reciprocal ohms per centimeter since L/A = 1cm^-1. Two sample conductance cells are given in Figure 3. These conductance cells do not satisfy these conditions, and so it is convenient to define a constant factor k determined by the cell geometry and called the cell constant, such that

 

                                  k = k/R                                                                            (3)

 

where R is the measured resistance of the actual cell. The numerical value of k for a particular cell is determined experimentally by use of a standard solution of known specific conductance.

 

 

     

 

 

 

 

 

 

 

 

 

 

Figure 3. Two conductance cell cells with different areas and different spacing.

 

The specific conductance will clearly depend on the concentration of the electrolyte.  In general, if the concentration doubles, the conductivity will double.  The measurement of conductance is put on an per ion basis by defining the equivalent conductance L:

 

                           L = k V                                                                                                (4)

 

where V is the volume of solution containing 1 g equiv. of solute. In the cgs system, V is in cubic centimeters and L is in square centimeters per ohm per equivalent.   For ions with a single charge, the equivalent conductivity is equal to the molar conductivity, L_m, obtained by dividing the conductivity by its molar concentration, m.

                              L_m = k /m                                                                                             (5)

 

KOHLRAUSCH’S LAW

Before proceeding, let’s stop for a moment and think about the conductivity of ions on a simple qualitative basis. Suppose the compound AB dissolves in water to produce the ions A⁺ and B^-. The conductance of ionic solutions is the result of the movement of ions through the solution to the electrodes. When two electrodes in the solution are made part of a complete electrical circuit, the cations (+) are attracted to the negative pole (cathode) and the anions (-) are attracted to the positive pole (anode). Changes in the equivalent conductance of an electrolyte solution with changes in concentration may result from changes in both the number and the mobility of the ions present. If both anion and cation are monovalent, the overall electrical neutrality of the soluble solution assures that equal numbers of the two ions are present. They will not in generally have the same mobility, however, and thus do not share equally in the conduction of current.

 

The solution conducts electricity through motion of the ions under the effect of an electric field. At high concentrations, each ion is surrounded by other ions, both positive and negative. The field affecting any particular ion changes slightly because of these surrounding ions. At infinite dilution, the distance between nearest neighbor ions is large, and only the effect of the applied electric field is felt by individual ions. This is the reason for extrapolating the data to infinite dilution.

 

The conductivity of any particular ion will also be affected by the ease with which the ion can more through the water. Hence different ions should contribute differently to the total measured conductivity. The ease with which any ion moves through the solution depends on considerations such as the total charge and the size of the ion; large ions offer greater resistance to motion through the water than small ions.

 

Suppose we now consider the compound CB, which dissociates on solution to produce C⁺ and B, where the ion B^- is the same as the B^- produced by the compound AB discussed above. One expects the contribution of the anion B^- to the total conductivity of the solution to be independent of the nature of the cation at infinite dilution.

 

Friedrich Kohlrausch(1840-1910) found that the molar conductivity varied as the square root of the concentration for many solutions.

 

L _m = L_m^o - k c^1/2                                                                                                              (6)

 

This basic relationship is one of Kohlrausch’s law where the limiting molar conductivity at infinite dilution, L_m^o , is a constant which depends on the electrolyte. The constant, k, depends more on stoichiometry of the electrolyte than its nature. The molar conductivities of several electrolytes are plotted as a function of the square root of concentration in Figure 5 to illustrate that concept.

 

 

 

 

 

 

 

 

 

 

 

 

Figure 4. The molar conductivity of NaCl, NaC₂H₃O₂, and KCl as a function of concentration.

 

 

 

 

 

 

 

 

 

 

 

 

Figure 5. The molar conductivity of NaCl and NaC₂H₃O₂, as a function of the square root of concentration.

 

 

The solid line in Figure 5 is a linear least squares fit of the data using equation (6) on the four most dilute solutions. The value of L^o for NaCl was found to be 126.0 W^-1 cm² mol^-1, while L^o for NaC₂H₃O₂ was 90.6 ± 0.1 W^-1 cm² mol^-1.

 

These molar conductivities at infinite dilution can be used to determine a conductivity at infinite dilution for individual ions to yield the law of the independent migration of ions. That is, the limiting conductivity of NaCl could be determined by the relationship:

 

L^o_NaCl = L ^o_Na+ + L ^o_Cl- (7)

 

When a current flows in an electrolyte solution, charge is carried by the motion of both anions and cations.

 

Figures 6 show plots of the molar conductances of several substances vs the square root of the concentration. The curves in Figures 6 indicate that this extrapolation can be carried out quite simply for strong electrolytes, whereas the problems engendered by weak electrolytes such as HAc (acetic acid) are severe. 

 

The increase of the equivalent conductance of solutions of strong electrolytes in the low-concentration range is not due to an increase in dissociation, because the dissociation is already complete, but to an increased mobility of the ions. In a concentrated solution of a highly ionized strong electrolyte, the ions are close enough to one another so that any one of them in moving is influenced not only by the electrical field impressed across the electrodes but also by the field of the surrounding ions. The ionic velocities are, then, dependent upon both forces. Arrhenius attempted to treat the electrolytic-conductance behavior of the strong electrolytes in the way in which he had successfully treated the weak electrolytes; such a treatment is, however, inconsistent with the experimental fact, discovered by Kohlrausch, that a plot of the equivalent

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Figure 3.  Molar conductivities of several electrolytes in aqueous solution at 298.15 K vs square root of the concentrations.

 

 conductance of a strong electrolyte against the square root of the concentration is very nearly linear. Debye and Hückel and Onsager have been able to calculate the effect of the surrounding ions on the mobility of any given ion and, for dilute solutions, have obtained results entirely consistent with the experimental facts. Complete dissociation is here assumed.

 

The conductivity of a solution is then given by the conductivities of the anion and the cation:                

                                L^o = L₊^o + L_-^o                                                                             (8)

Equation (8) describes the true state of affairs and states Kohlrausch’s law of

independent mobilities of ions in infinitely dilute solutions. It is amenable to easy

verification, since the difference in L^o for pairs of salts with a common ion should be the same regardless of the common ion.   Samples of such calculations are given at the end of this write-up in Table 2.

 

WEAK ELECTROLYTES

THEORY[6,7]:  A factor which affects the equivalent conductance of a solution that was not considered above is the possible limited dissociation of the electrolyte.  Some electrolytes, known as weak electrolytes, do not dissociate completely in solution. Instead, there is an equilibrium between ions and associated electrolyte.  Acetic acid is a typical weak electrolyte. The apparent equilibrium constant for dissociation may be calculated as

    

where  a = degree of dissociation

c = concentration of the solute

 

According to the Arrhenius theory, the equivalent conductance at any concentration is related to the degree of dissociation by

 a = L/L^o

where L = equivalent conductance at concentration c

L^o = equivalent conductance at infinite dilution

 

In the case of a weak electrolyte the value of L^o cannot be obtained by the extrapolation to infinite dilution of results obtained at finite concentration, because L is a rapidly varying and nonlinear function of L_c. Instead, L^o is obtained by the of the law of Kohlrausch:

 

L^o_HR = L^o_HCl + L^o_NaR - L^o_NaCl                                                                                 (11)

 

This is equivalent to saying that for a simple, binary electrolyte like HR,  L^o_HR= I^o_H+ + I^o_R^-.

The equivalent conductance at infinite dilution is a sum of ionic contribution, but l^o_H+, for example, is independent of the electrolyte from which the hydrogen ions are obtained. Since HCI and the salts NaR and NaCI are all strong electrolytes, the values on the right-hand side of Eq. (11) can all be obtained by extrapolation, as discussed in previous section.

 

The apparent equilibrium constant K_a, is equal to the true equilibrium constant, which is expressed in terms of activities, only for ideal solutes.

                                                                    (12)

where g _i is the activity coefficient of species i. Since g _i? 1 for infinitely dilute solutions,

 

Lim K_a = K                                                                                                         (13)

c? 0

Arrhenius noted that the molar conductivities of electrolytes decreased with increasing concentration. He considered a solution of the salt AB to consist partly of unionized AB molecules and partly of cations and anions. He attributed the decrease in conductivity to the decrease in the degree of ionization of the electrolyte. Suppose that a solution of AB is prepared with n molecules of AB per cubic centimeter of solution.  The degree of ionization, a, is determined by the number of positive ions per cubic centimeter, n₊ = a n, or the number of negative ions, n_- = a n. Suppose that the velocities with which the ions move through the water are u₊ and u_-. The total current carried across a unit area is then

 

I = n₊eu₊ + n_-(-e)(-u_-) = ne(u₊ + u_-)a                                                                            (14)

 

where the negative sign for u_ arises because the two different ions move in opposite directions; the electron charge is denoted by e.

 

The mobility m of an ion is defined as the velocity per unit field strength, or

m₊ = u₊/E and is m _-= u_-/E. From Ohm’s law and the definition of conductivity, then k = I/E, where I is the current across a unit area and E is the electric field. The specific conductivity of the solution is

       k = ne(m₊ + m_-)                                                                                                 (15)

 

The equivalent conductivity is then

 

    L = N_oe(u₊ + u_-)a                                                                                              (16)

 

where N_o is the Avogadro number. Since N_oe is the charge on a mole of electrons, or the Faraday constant F, (15) can be written as 

 

                L^o = F(m₊ + m_-)                                                                                     (17)

 

It is here that Arrhenius made the critical assumption that at infinite dilution

ionization is complete, that is, as m ®0, a ®1. In the limit,

 

               L = F(m₊ + m_-)a                                                                                    (18)

 

and by combining Equations (17) and (18) we get the expression for the degrees of ionization,

 

                         a = L /L^o                                                                                  (19)

 

This model accomplishes two things. First, it enables the calculation of the degree of ionization of electrolytes from conductivity data, and secondly it provides an explanation of Kohlrausch’s law of independent migrations. The mobilities of the ions are independent of the chemical constitution and the terms in eq. 18 can be interpreted by taking

                            L _o⁺ = Fm ₊ and L _o^- = Fm -                                                          (20)

 

Suppose the dissociation of acetic acid at a concentration of c moles per dm³ is take as an example.

                                     HAc ? H⁺ + Ac^-                                                              (21)

                                    c(1-a)    ca    ca

The concentrations of the various species at equilibrium have been written in terms of c and the degree of dissociation.  Shortly after Arrhenius first proposed his ionic dissociation theory, Ostwald applied the law of mass action to a partially ionized substance. For HAc :

                                                                                    (22)

 

where the brackets indicate molar concentrations.

 

Measurements of K_a over a range of concentrations thus provide a way of estimating K and the activity coefficient factor g_H+g_A- /g_HA at the various concentrations. To a good approximation g _HA may be considered to be unity. The value of the product g_H+ g_A- is, by definition, g_± ² the square of the mean ionic activity coefficient for the weak electrolyte HA.

 

That is the limiting conductivity of a solution of ions is the sum of the limiting conductivities of the individual ions. This can be extended to determine the conductivity of the weak acid, acetic acid, at infinite dilution from the conductivities of hydrochloric acid and sodium acetate minus that of sodium chloride.

 

L _HAc^o = L _HCl^o + L _NaAc^o - L _NaCl^o

H⁺ + Ac^- = H⁺ + Cl^- + Na⁺ + Ac^- - Na⁺ -Cl^-

 

As the concentration of a weak acid increases, a new species, the undissociated acid (HAc) is present in the solution. This can be contrasted with the solubility of a salt. When a soluble salt such as sodium chloride dissolves in water, Na⁺ ions and Cl^- ions are present in solution. As the solution becomes saturated the solid sodium chloride precipitates out of solution. However, for acetic acid, the HC₂H₃O₂ species, will be in solution and in equilibrium with the ions H₃O⁺ and C₂H₃O₂^-. The conductivity of the solution of a concentration, c, can be used to determine the degree of dissociation, a =L /L^o and then the equilibrium constant:

 

[H₃O⁺] = ac, [C₂H₃O₂^-] = ac, [HC₂H₃O₂] = (1 - a)c

 

If the normalities of the acids are not known accurately, they are determined by titration with standard base. The conductance of a 0.05 N solution of acetic acid

is determined, and a 0.025 N solution is then prepared by quantitative dilution with conductance water. In this fashion, conductance measurements are made on 0.05, 0.025, 0.0125, 0.00625, 0.00312, and 0.00156 N acetic acid solutions.

 

Calculations

For the HCl, NaCl, and NaC₂H₃O₂ solutions, conductance measurements are then made on solutions of hydrochloric acid at concentrations of 0.1, 0.05, 0.025, 0.0125, 0.00625, 0.00312, and 0.00156 N. Successive dilutions must be made with great care; otherwise propagation of error will become excessive. They may be made with calibrated volumetric flasks and pipettes. For each concentration of each system measured, the equivalent conductance is calculated.  Plot equivalent conductance vs. Öc and c for HCl, NaCl, NaC₂H₃O₂, and HC₂H₃O₂ and discuss the effect of concentration on conductance for strong and weak electrolytes. From your plots of equivalent conductance vs. for the electrolytes, discuss the effect of concentration on equivalent conductance for strong and weak electrolytes.

 

Determine values of L^o for the strong electrolytes by an appropriate extrapolation of your data and compare your results to literature[3].  Using these values, obtain a value of L^o for HC₂H₃O₂ using Kohlrausch's Law. For the acetic acid the apparent equilibrium constant for dissociation should be calculated for each concentration and the values of the true equilibrium constants estimated. Now, calculate the apparent equilibrium constant for dissociation and estimate by means of an appropriate plot a value for the true thermodynamic equilibrium constant, K_a, for acetic acid. Compare your values with literature values.  At larger concentrations, however, the value of K deviates from its small concentration value as a result of the nonideality of the solution.  Comment on how the data obtained in this experiment demonstrates or does not demonstrate this principle.