Lab.14.CONDUCTOMETRY Key words:

electric current, ion migration, resistivity, insulator, conductivity, molar conductivity,

Kohlorausch equation, water, solution of salts, acids, bases and their conductivity, ionization,

degree of ionization, titration, equivalence point, electrolytes, nonelectrolytes, constant of

ionization.

Literature:

J. Brady; General Chemistry, Chapters 5, 14, 17, 18, 21.

Theoretical background

Conductometry measures electrolytic conductivity or resistance of the solution located between

two electrodes. This property of the solutes is the result of the ion movement under the influence

of the electric field applied. Only ionizable molecules (electrolytes) conduct electric current. The

magnitude of the conductivity depends on the amount of ions presented in the solution. The

electrolyte kind, its concentration and temperature affects both conductivity and specific

conductance.

Conductance is the extrinsic property while conductivity is the intrinsic property. This means that

conductance is the property of an object dependent of its amount/mass or physical shape and size,

while conductivity is the inherent property of the material that makes up the object. No matter how

the object changes in terms of shape/size/mass, as long as it is made of the same material and the

temperature remains the same, its conductivity does not change.

The specific conductance (κ) of a solution containing one electrolyte depends on the

concentration of the electrolyte. Therefore, it is convenient to divide the specific conductance by

concentration. This quotient, termed molar conductivity is denoted by Λm

Λ𝑚 = 𝜅

𝑐

Strong electrolytes dissociate almost completely in solution. The conductivity of a solution of a

strong electrolyte at low concentration follows Kohlrausch’s law :

Λ𝑚 = Λ𝑚0 − 𝐾√𝑐

where Λ𝑚0 is known as the limiting molar conductivity, K is an empirical constant and c is the

electrolyte concentration. (Limiting here means "at the limit of the infinite dilution".) In effect,

the observed conductivity of a strong electrolyte becomes directly proportional to concentration,

at sufficiently low concentrations i.e. when

Λ𝑚0 ≫ 𝐾√𝑐

2 As the concentration is increased however, the conductivity no longer rises in proportion.

Moreover, Kohlrausch also found that the limiting conductivity of an electrolyte; 𝜆+0 and 𝜆−

0 are

the limiting molar conductivities of the individual ions.

The following table gives values for the limiting molar conductivities for selected ions.

Table of limiting ion conductivity in water at 298K (approx. 25oC)

Cations Cations Anions Anions

H+ 34.982 Ba2+ 12.728 −OH 19.8 SO42− 15.96

Li+ 3.869 Mg2+ 10.612 Cl− 7.634 C2O42− 7.4

Na+ 5.011 La3+ 20.88 Br− 7.84 HC2O41− 40.2

K+ 7.352 Rb+ 7.64 I− 7.68 HCOO− 5.6

NH4+ 7.34 Cs+ 7.68 NO3

− 7.144 CO32− 7.2

Ag+ 6.192 Be2+ 4.50 CH3COO− 4.09 HSO32− 5.0

Ca2+ 11.90 ClO4− 6.80 SO3

2− 7.2

Co(NH3)63+ 10.2 F− 5.50

Contrary to strong electrolytes weak electrolytes are not fully dissociated in water solution. Thus

their molar conductance is lower than strong electrolytes (see Figure below).

Fig. 1. Comparison of the molar conductance of various electrolyte type.

In this case there is no limit of dilution below which the relationship between conductivity and

concentration becomes linear. Instead, the solution becomes ever more fully dissociated at weaker

concentrations, and for low concentrations of "well behaved" weak electrolytes, the degree of

dissociation of the weak electrolyte becomes proportional to the inverse square root of the

concentration.

𝚲𝒎

3 For a monoprotic acid, HA, obeying the inverse square root law, with a dissociation constant Ka,

an explicit expression for the conductivity as a function of concentration, c, known as Ostwald's

dilution law, can be obtained.

1

Λ𝑚=

1

Λ𝑚0 +

Λ𝑚 𝑐

𝐾𝑎 (Λ𝑚0 )2

The molar conductivity at a given concentration and limiting molar conductivity for weak

electrolytes may be used for determination of its degree of dissociation according to the formula

below:

𝛼 = Λ𝑚

Λ𝑚0

Conductometry is a measurement of electrolytic conductivity to monitor a progress of chemical

reaction. Conductometry has notable application in analytical chemistry, where conductometric

titration is a standard technique. In usual analytical chemistry practice, the term conductometry is

used as a synonym of conductometric titration, while the term conductimetry is used to describe

non-titrative applications.

Conductometric titration is a type of titration in which the electrical conductivity of the reaction

mixture is continuously monitored as one reactant is added. The equivalence point is the point at

which the conductivity undergoes a sudden change (see Fig.2a, b).

Fig.2. Conductometric titration of weak acid with strong base (2a) and precipitation titration of

barium chloride with sulfuric acid (2 b).

Marked increases or decrease in conductance are associated with the changing concentrations of

the two most highly conducting ions—the hydrogen and hydroxyl ions. The method can be used

for titrating coloured solutions or homogeneous suspension (e.g.: wood pulp suspension), which

cannot be used with normal indicators. Acid-base titrations and redox titrations are often

performed in which commonly indicators are used to locate the end point e.g., methyl orange,

a b

4 phenolphthalein for acid base titrations and starch solutions for iodometric type redox process.

However, electrical conductance measurements can also be used as a tool to locate the end point

e.g. in a titration of a HCl solution with the strong base NaOH. As the titration progresses, the

protons are neutralized to form water by the addition of NaOH. For each amount of NaOH added

equivalent amount of hydrogen ions is removed. Effectively, the faster moving H+ cation is

replaced by the slower moving Na+ ion, and the conductivity of the titrated solution as well as the

measured conductance of the cell fall. This continues until the equivalence point is reached, at

which we have a solution of sodium chloride, NaCl. If more base is added an increase in

conductivity or conductance is observed, since more ions (Na+ and OH- are being added and the

neutralization reaction no longer removes an appreciable number any of them. Consequently, in

the titration of a strong acid with a strong base, the conductance has a minimum at the equivalence

point. This minimum can be used instead of an indicator dye to determine the endpoint of the

titration. Conductometric titration curve, that is a plot of the measured conductance or conductivity

values against the number of milliliters of NaOH solution.

5

Objectives:

Measurement of purity degree of water.

Determination of HCl and acetic acid in mixture by conductivity titrations.

Determination of end-point in precipitation titrations.

Part 1 Measurement of the electrical conductivity for:

a) tap water

b) tap water after boiling

c) distilled water

d) deionized water

e) mineral water

Procedure:

Step 1.

Prepare 40 mL of the tap water in beaker of 50 mL.

Step 2.

Measure the conductance of water and present result in Table 1.

Steps 3-9.

Repeat steps 1-2 for various water types (from points b-e)

Table 1. Conductance of the different type of water

Type of water Conductance (mS or μS)

Tap water

Tap water after boiling

Distilled water

Deionized water

Mineral water

Answer the question:

What components of the water cause different conductance?

6 Part 2

Titration of mixture HCl and CH3COOH.

Step 10.

To the 100 mL of volumetric flask add 6 mL HCl and 8 mL CH3COOH, distilled water to the mark

and mix.

Step 11.

Transfer 25 mL solution of acids mixture into a150 mL beaker, add approximately 75 mL water,

immerse the conductance cell. There must be enough solution to ensure that the electrodes are

completely immersed.

Step 12.

Read the initial value of conductance.

Step 13.

Carefully titrate the solution of acids with 0.5 mL 0.1 M · L-1 NaOH.

Step 14.

After addition each portion of alkali solution, switch out stirrer and read the value of conductance.

The titration will be finish when the conductance is equal to the initial value.

Step15.

Record the results in Table 2.

7 Table 2. Changes of solution conductivity during titration

V mL NaOH mS V mL NaoH mS

0.0 16.5

0.5 17.0

1.0 17.5

1.5 18.0

2.0 18.5

2.5 19.0

3.0 19.5

3.5 20.0

4.0 20.5

4.5 21.0

5.0 21.5

5.5 22.0

6.0 22.5

6.5 23.0

7.0 23.5

7.5 24.0

8.0 24.5

8.5 25.0

9.0 25.5

9.5 26.0

10.0 26.5

10.5 27.0

11.0 27.5

11.5 28.0

12.0 28.5

12.5 29.0

13.0 29.5

13.5 30.0

14.0 30.5

14.5 31.0

15.0 31.5

15.5 32.0

16.0

8 Step 16.

From the results recorded in Table 2 plot the graph conductance vs. volume of titrant alkali.

Determine the end-points from the graph.

The contents of acids in mixture calculate from:

1000

4CVM)g(m M

Plot of………

Part 3. Conductometric determination of salicylic acid KA values.

Step 17.

Prepare the 5 aliquots of salicylic acid using the dilution equation. Use graduated cylinders and

pipettes for accuracy. The concentrations of solution should be in accordance with salicylic acid

in Table 3.

Step 18.

Pour 40 mL of the 0.004M of salicylic acid into beaker of 50 mL.

9 Step 19.

Measure the conductance of salicylic acid solution and present result in Table 3.

Step 20.

Clean the electrode with distilled water and dry it with tissue.

Steps 20-28.

Repeat steps 18,19 for diluted salicylic acid solutions.

Table 3. Conductance of the salicylic acid solution

C mol/L

Λ

[mS]

κ = λk* cm

m

m

o

1

cK

2

pKa

0.004

0.002

0.001

0.0005

0.00025

k* = conductometric cell constant (0.3286 cm-1)

Base on results presented in table 3, graph Λ = f (√C) and determine value of Λ for C = 0.

10

Plot of…..

Average of pKa =

The values of pKa of salicylic acid let us classified it as strong or weak acid?