Determination of the Conductance of Strong and Weak Electrolytes

Objective:

1. To measure the conductance of potassium chloride, hydrochloric acid, sodium chloride and sodium acetate

2. To determine the dissociation constant of acetic acid

Introduction:

An electrolyte is any substance containing free ions that make the solution to be electrically conductive. Strong electrolyte can dissociate completely in water to form ions while weak electrolyte dissociate partially in water to form its ions. The conductance of solution is depends on the degree of dissociation of electrolyte. The higher the degrees of dissociation of electrolyte, the more ions are produced in the solution, hence the better the conductivity of electrolyte. According to Ohms’ law,

E = I R

where E is the potential difference, I is the current measured and R is the resistance. The term conductance is generally used for dealing with electrolyte and this is defined as the reciprocal of the resistance of the solution. The relationship between resistance and the conductance of the solution is defined as below:

L (ohm^-1 or Ω^-1) = 1 / R

Once the resistance, R is known, the conductivity or specific conductance (X) may be obtained from

X (Ω^-1 cm^-1) = d / AR

where d is the distance separation between two electrodes of the measurement cell, A is refers to the area. The cell constant k of the conductivity cell is defined as

k = d/A

and hence

X = kL

The specific conductance, X is the reciprocal of the resistance in ohm of a 1cm of liquid at a specified temperature. The molar conductivity, Λ of an electrolyte solution is defined as:

Λ (Ω^-1 mol^-1 cm²) = X /C

For weak electrolyte, the increase of molar conductivity with increasing dilution is ascribed to increased dissociation of the electrolyte molecules to free ions. However, a thermodynamic equilibrium exists between the un-dissociated molecules and the ions formed from dissociation. The dissociation degree, α at the given concentration, C is given by

α = Λ / Λ_o

where Λ_o is the molar conductivity in the limit of zero concentration or limiting molar conductivity. For strong electrolytes, the molar conductivity is higher than those of weak electrolyte at high concentrations. As the solutions become dilute, the molar conductivities also increase in the case of weak electrolytes but the variation is less steep than for weak electrolytes.

Kohlrausch’s Law states that at time infinite dilution, the molar conductivity of an electrolyte can be expressed as the sum of the contribution from its individual ions.

Λ_o = Λ⁺_o + Λ^-_o

where Λ⁺_o and Λ^-_o are refers to the conductivities of cation and anion at infinite dilution. In this experiment, the molar conductivity of weak electrolyte is determined by using Kohlrausch’s Law.

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Given that the molar conductivities of strong electrolyte are expressed as the following:

Alternatively, Λ_o and the dissociation constant, k_a of weak electrolyte may be obtained from the Ostwald dilution law:

Materials:

0.2000M potassium chloride solution, 0.1000M acetic acid, 0.1000M, hydrochloric acid, 0.1000M sodium chloride solution, 0.1000M sodium acetate solution

Apparatus:

Conductivity meter, 100ml dilution flasks, pipette, burette

Procedure:

Part 1: Determination of cell constant

1. The conductance (L) of 0.2000M potassium chloride solution was measured.

2. The cell constant (k) was determined by using equation 6 with the given specific conductance (X) of this solution of 2.768x10^-3 ohm^-1 cm^-1.

Part 2: Measurement of conductance

1. From the solution acetic acid provided, the successive solution with concentration of 0.0500M, 0.0250M, 0.0125M, 0.00312M, 0.00156M and 0.00078M solution were prepared.

2. The conductance of these solutions was measured.

3. The procedure was repeated with hydrochloric acid, sodium chloride, and sodium acetate.

4. The conductance of water used was measured.

5.

Results and calculations:

Part I Determination of Cell Constant

Conductance, L of:

KCl = 30.07 mΩ^-1

Distilled water = 0.00 mΩ^-1

Specific conductance, X = 2.768 × 10^-3 Ω^-1 cm^-1

Cell constant, k = X / L

= 2.768 mΩ^-1 cm^-1 / 30.07 mΩ^-1

= 0.09205 cm^-1

Table 1 Conductance of solutions with different concentration

Table 2 Specific conductance of each solution with different concentration

Table 3 Molar conductivity of each solution with different concentration

By using Kohlrausch’s law, the

Table 4 Degree of dissociation and dissociation constant of CH₃COOH at different concentration

Concentration of CH3COOH (mol dm-3)	Degree of dissociation, α = Λ / Λo	Dissociation constant, ka of CH3COOH
0.10000	0.0189	3.64 x10-5
0.05000	0.0310	4.96 x10-5
0.02500	0.0388	3.92 x10-5
0.01250	0.0511	3.44 x10-5
0.00625	0.0622	2.58 x10-5
0.00312	0.1245	4.85 x10-5
0.00156	0.1950	7.37 x10-5
0.00078	0.2471	6.33 x10-5

*degree of dissociation is calculated by using the formula of α = Λ / Λ_o

*dissociation constant is calculated by using

Average value of dissociation constant, k_a

= [(3.64 + 4.96 + 3.92 + 3.44 + 2.58 + 4.85 + 7.37 + 6.33) × 10^-5] / 8

= 4.636 × 10^-5

 

Table 5 1/Λ of CH₃COOH with CΛ.

1 / Λ (Ω cm-2 mol)	CΛ (× 10-3 Ω-1 cm-1)
1.597	0.0626
0.971	0.0515
0.776	0.0322
0.590	0.0212
0.484	0.0129
0.242	0.0129
0.154	0.0101
0.122	0.0064

Graph of 1/Λ versus CΛ was plotted.

From Graph 4,

Compared to y = mx + c

m = 1/ (k_a x Λ_o²), c = 1 / Λ_o

1 / Λ_o = 0.013 Ω cm^-2 mol

Λ_o = 1 / 0.013 Ω cm^-2 mol

= 76.92 Ω^-1 cm² mol^-1

m = 23 = 1/ (k_a x Λ_o²)

k_a = 1/ (76.92)² (23)

= 7.35 x10^-6

 

Precaution steps:

1. Use distilled water to rinse the electrodes before use.

2. Shake the electrode briefly to ensure that no air bubbles trapped in electrode.

3. Ensure the electrode surface is completely submerged in solution.

Qualitative Determination of Organic Compounds by Infrared Spectroscopy

Objective:

1. To determine the structural information on cinnamic acid prepared by using KBr disc technique

2. To identify an unknown sample by comparing it with the known sample

Introduction:

Infrared spectrometer (also known as spectrophotometer) is the instrument that used to determine the absorption spectrum for a compound. Two types of spectrometers are commonly being used in the analysis or research laboratory which includes dispersive and Fourier Transform (FT) infrared spectrometer. Both of these types of spectrophotometer provide the spectra of compounds in the common range of 4000 to 400 cm^-1. Infrared region (IR) is divided into three regions which includes near IR, mid IR and far IR. The unit of IR measurement is wavenumber, the number of waves per centimeter, (cm^-1). Although two provide nearly identical spectra for a given compound. FT infrared spectrometers provide the infrared spectrum much more rapidly than the dispersive spectrometers.

Infrared spectroscopy is resulted from the molecule vibration in which it includes stretching and bending. Every type of bond has a different natural frequency of variation and two of the same type of bond in two different compounds are in two slightly environments, two different structures have exactly the same infrared absorption pattern, or infrared spectrum. Although some of the frequencies absorbed in the two cases might be the same, but no identical infrared spectrum will appear in the two different molecules. Hence, infrared spectrum can be used for molecules much as a fingerprint can be used for human beings.

The main function of the infrared spectrum is to determine the structural information about chemicals either organic or inorganic. The functional groups that exist in a compound can be determined based on the spectra obtained. For example, infrared spectroscopy is also can be used to identify the chemicals from spills, paints, polymers, coatings, drugs and contaminants. The absorption of each type of bond is regularly found only in certain small portions of the vibrational infrared region. A small range of absorption can be defined for each type of bond. However, the absorptions are normally due to some other type of bond if outside this range. For an organic compound, it has a very rich and detailed spectrum while for an inorganic compound is usually much simpler.

In order to prepare a Kbr (potassium bromide) pellet, mortar and pestle, die set, sample holder and hydraulic KBr press are very important in preparation.

Mortar and pestle

  

Die set

Hydraulic KBr press

sample holder

A KBr pellet is a dilute suspension of solid in a solid. It sis usually obtained by first grinding the sample in anhydrous KBr at a ratio of approximately 1 part sample to 10-100 parts KBr. The mixture is then being placed on a steel plate containing a paper card with a hole punched in it. The sample is placed into the centre so that it will lie in the infrared beam when placed on the spectrometer. A second steel plate is placed over the sample and the steel sandwich is placed in a hydraulic press and subjected to pressures of 15000 psi. Decompression usually will give a KBr pellet that is reasonably transparent both to visible light and infrared radiation. The only limitation of KBr is that it is hydroscopic. This is usually a good idea to obtain a spectrum run as a Nujol mull with sample. Since Nujol mull is a hydrocarbon and has no affinity for water when compared to KBr. Any absorption in Nujol between 3400- 3600 cm^-1 can be attributed to the sample and not to the absorption of water by KBr.

Materials: KBr powder, cinnamic acid, acetone, unknown sample, mortar and pestle

Apparatus: FTIR machine, hydraulic KBr press

Procedure:

Sample preparation by using KBr disc

1. The pure KBr powder was grinded by using mortar and pestle.

2. KBr was placed into the die set in which the sample is sandwiched by the steel plate with smooth surface.

3. The die set was compressed tightly by using a hydraulic KBr press.

4. The spectrum of pure KBr was obtained.

5. The cinnamic acid and KBr powder were mixed homogeneously and the spectrum of this mixture was obtained.

6. An unknown sample’s spectrum was obtained through KBr disc.

7. The spectrum of cinnamic acid and unknown were compared.

Results:

The spectrums of cinnamic acid and unknown compound, please refers to the spectrums obtained.

Discussion:

Potassium bromide (KBr) powder is always being used in the infrared spectroscopy analysis in which KBr does not absorb infrared region from 4000 cm^-1 to 400 cm^-1. KBr is transparent to IR radiation in the range of above 400 cm^-1 and has no absorption bands in the region traditionally used for IR spectroscopy. The absorption bands recorded on the KBr disc are come from the sample or impurities present in the KBr mixture. The disadvantage of using KBr pellet in IR spectroscopy is due to its hygroscopic property. The KBr pellet absorbs the water that appears in the atmosphere and hence causes the existence of a broad band in the IR spectrum at around 3400cm^-1 to 3200cm^-1. In order to prevent this problem, the KBr powder must always being kept in the oven before it is being used in the mixing of sample. Besides, KBr disc is prepared in the solid state because this can prevent the reaction between sample and atmospheric contaminants or solvent to occur easily.

Every molecule will have its own characteristic spectrum. The bands that appear depend on the types of bonds and the structure of the molecule. The functional groups that present in the cinnamic acid involves –COOH and C=C bonds. Based on the spectrum obtained, the C=O double bond stretch exists at the wavenumber of 1681cm^-1 due to the conjugation effect and hence it shifted to lower wavenumber. Besides, the –COOH functional group has a broad band O-H stretch and a C-O stretch with the wavenumber of 3100cm^-1 ~ 2800cm^-1 and wavenumber of 1200cm^-1 ~ 1320cm^-1 respectively. Since the cinnamic acid is polar molecule, the formation of dimer of carboxylic acid caused the O-H stretch shifted to from 3400cm^-1~3300cm^-1 to the current wavenumber of 3100cm^-1~ 3000cm^-1. Besides, the O-H stretch is actually overlapped the sp² C-H stretch in which sp² C-H should appears between 3100cm^-1 ~ 3000cm^-1. The two bands at 1418cm^-1 and 1628cm^-1 are represented by the aromatic C=C double bond in cinnamic acid. The aromatic C=C double bond is actually present between 1700cm^-1 ~ 1500cm^-1, but they shifted to lower wavenumber due to the conjugation effect of benzene ring. In addition, the band of aliphatic C=C double bond is overlapped by the aromatic C=C double bond stretch which actually present between 1680cm^-1 ~ 1620cm^-1 with a single peak (aromatic has two peaks). The figure 1 below shows the structural molecule of cinnamic acid:

Figure 1 Structure of cinnamic acid

Based on the spectrum of unknown compound, the very characteristic of the compound is the C=O double bond stretch near 1700cm^-1. Since the C=O stretch is present at around 1700cm^-1, no conjugation occurs in the C=O double bond. The band is very sharp because C=O is a strong absorber. The bond absorption takes place at 1604cm^-1 indicated that the compound contains C=C double bond. This is an aliphatic compound instead of aromatic compound since there is absence of C=C stretch at 1400cm^-1. Due to the presence of C=C double bond, the sp² C-H stretch is also present in the spectrum which is indicated by the absorption band at 3056cm^-1. This compound is not an amide although there is a peak present at 3417cm^-1. This is because amide has a strong band in this IR region. The particular band with 3417cm^-1 may be due to the contribution of other bonds. Based on the presence of C=O and C=C, the unknown compound is predicted as an aliphatic ketone with γ or higher degree of C=C double bond.

Precaution steps:

1. Make sure the mortar and pestle, die set and sample holder are being cleaned and washed with acetone to avoid contamination.

2. Make sure the KBr disc does not expose too long to atmosphere.

3. Make sure the mixture of KBr and sample are mixed homogeneously.

4. The die set must be placed in the centre of the hydraulic press in order to prevent it to be spoiled.

Beer-Lambert Law

Objective:

1. To verify the Beer-Lambert Law

2. To determine the composition of complexes by using Job Method

Introduction:

Job method is also known as method of continuous variation. This method is used to determine the composition of a complex which is formed by two reacting species. It is most effective to be applied when only a single complex is formed in the solution. Job’s method is based on the concept that equimolar solution of metal-ion and ligand are mixed gradually by using different volume ratio. As the concentration of metal ion increase, the concentrations of ligand will decerease. It maintains the total number of mole reactants to be constant in a series of mixture of reactants. In this experiment, the mixture is made up of different fraction of nickel sulfate and ligand ethylene diamine. A wavelength at which the complex absorbs the strongest, λ_max is selected. Then, the absorbance of each solution in the series at the wavelength of maximum absorbance is determined by using spectrometric method.

Job Method is often used to determine the soluble Ni²⁺ - en complexes in the solution. The “n” value of the complexes also can be calculated.

Z + nL à ZL_n

where Z represents Ni²⁺ ion, L refers to the ethylene diamine (en). Different complexes could be formed in a mixture of metal Ni²⁺ ion and ligand en, for example, Ni(en)²⁺, Ni(en)₂²⁺, Ni(en)₃²⁺ and many more. The equilibrium constant of each Ni²⁺ - en complexes are shown in the following:

where K_1, K₂ and K₃ are the equilibrium constant for each reaction respectively. Species that formed the most in the solution depends on the relative value of the equilibrium constant. If the value of K₂ is bigger than K₁, that means the concentration of Ni(en)₂²⁺ is higher than Ni(en)²⁺. In this experiment, every solution prepared contains the same concentration of Z and L. this is to ensure that the value of equilibrium constant of the formation of Z(L)_n is large and the absorption for the species is maximum when the concentration of ligands en is exactly “n” times the concentration of ion Z. The value of “n” can be calculated if the concentration rate of L/Z is known for the solution that gives the maximum absorption.

Measurement of optical density (absorbance) at the λ_max will show the maximum when the ratio of ethylene diamine to nickel sulfate is equally present in the particular mixture. This is because the solution contains the highest concentration of complex. So, in the graph of absorbance against mole fraction of ligand in the mixture will show a region starting with zero and increasing as the mole fraction of ligand increases as well as the concentration of complex. The absorbance of complex will also increase at the same time. At this region with positive slope in the graph, the ligand ethylene diamine is acting as limiting reagent. Further addition of mole fraction of ligand will decrease the mole fraction of nickel sulfate in the mixture. Since the nickel sulfate is insufficient to form complex with excess ethylene ligand, thus the absorbance due to less formation of complex then falls. The diagram 1 below shows the general trend of continuous variation of complex.

Diagram 1

According to Beer-Lambert law, the equation can be expressed as A = εcl where A is the absorbance, ε is refers to the molar absorptivity of liquid (L mol^-1 cm^-1), c is concentration of absorbing material (mol L^-1) and l is the optical path length (cm). The amount of attenuation is depends on the concentration of absorbing molecules and path length over which absorption occurs. The absorbance of the complexes formed is directly proportional to the concentration of the complexes formed in the solution. The value of “n” could be calculated by using the formula as below:

Apparatus:

10ml volumetric flask, dropper, pipette

Materials:

UV-Vis spectrometer, 0.4M ligand ethylene diamine, 0.4M nickel sulfate solution, distilled water

Procedure:

1. 100cm³ of nickel sulfate (NiSO₄.6H₂O) with the concentration of 0.4M and 100cm³ of ligand-en with concentration of 0.4M were prepared.

2. The spectrums of each stock solution that has been prepared were obtained in the range of 500-65-nm.

3. Solutions with a total volume of 10cm³ were prepared in which the mole fraction of en and X is 0.0, 0.3, 0.4, 0.5, 0.6, 0.7, 0.8, 0.9 by using the stock solutions of NiSO₄.6H₂O and ethylene diamine.

4. The values of Y with, at least five different wavelengths were calculated and graph of X vs. Y was plotted for each results.

5. The value of X was determined when Y is at the maximum for each graph. By using the value of X in equation (6), the n values for Ni(en)_n²⁺ complexes were calculated.

Results and calculations:

Table 1 Absorbance (A) at different wavelength, λ (nm)

In order to plot the graph of X versus Y, the equation of Y = [A – (1-X) A_z]b is used.

where A_z is the absorption of pure Ni²⁺ at the selected wavelength, b is optical path length

Table 2 Absorbance (Y) at Different Wavelength, λ (nm)

From Graph 1,

At λ= 530 nm, X = 0.8.

n₁ = X / (1 - X)

= 0.8/ (1 - 0.8)

= 4

From Graph 2,

At λ = 545 nm, X = 0.715

n₂ = X / (1 - X)

= 0.715 / (1 – 0.715)

= 2.51

From Graph 3,

At λ = 578 nm, X = 0.655

n₃= X / (1 - X)

= 0.655/ (1 – 0.655)

= 1.90

From Graph 4,

At λ = 622 nm, X = 0.5

n₄= X / (1 - X)

= 0.5/ (1 – 0.5)

= 1

From Graph 5,

At λ = 640 nm, X = 0.455

n₅= X / (1 - X)

= 0.455/ (1 – 0.455)

= 0.83

Average of value of n = (4 + 2.51 +1.90 + 1 + 0.83) / 5

= 2.048

Thus, the n value for Ni(en)_n²⁺ complexes = 2 since the n value must be an integer.

Discussion:

The reaction between nickel sulfate and ligand ethylene diamine produces nickel(II) bis(ethylenediamine) complex as product. This metal complex is a cationic complex. This is because the en anion carrying 0 charge since it is a neutral ligand while Ni²⁺ carrying +2 charge. The en ligand did not loss any proton during the formation of Ni-en complex so that it is a neutral ligand. Hence, the net charge of the Ni(II) complex is +2. The ethylene diamine ligand is a bidendate ligand which each en ligand coordinated to the Ni²⁺ metal via two dative bonds. The lone pair electron was donated from each nitrogen atom in ethylenediamine to form the two coordination bonding to Ni(II) metal. Since en ligand are weakly coordinated to the metal species, the en ligand is easily to be bonded to the Ni(II) metal so that the reaction can be took place at room temperature because the reaction is just required a small amount of energy.

In this experiment, the value of n was determined. Thus, the composition of the Ni(II) complex also was determined based on the value of n obtained. The ratio of Ni(II) cation to the ethylenediamine ligand in the metal complex is 1:2. For each metal complex in the solution, the Ni²⁺ metal is coordinated with two ethylenediamine ligand. The Ni(II) complex is predicted has the geometry of tetrahedral with the Ni(II) in the centre of complex. In the structure of complex, two ethylenediamine neutral ligands were coordinated surrounding to the Ni²⁺ metal centre. The ethylenediamone ligand is known as bidentate ligand (chelating ligand) which has can forms two coordinating point to the metal. The chelating effect allows the ethylenediamine bonded strongly to the Fe(III) complex.

Based on the data obtained, Beer-Lambert law was verified. This is because the data indicated that the absorbance is directly proportional to the concentration of complex formed and the path length over which absorption occurs. The path length is defined as 1cm in the experiment, because the curvette size is generally manufactured in 1cm. Since the path length is being held constant, thus the absorbance increases as the concentration of complex formed in the solution increases.

Precaution steps:

1. Gloves and goggles must be worn in order to prevent direct exposure to any chemicals.

2. The readings of apparatus must be taken parallel to the eyes in order to avoid parallax error which can cause the deviation in mole fraction.