Determination Of Stability Constant Of A Molecular Complex Biology Essay

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The aim of the experiment is to examine the donor-acceptor complex formed between iodine (ion-acceptor) and alcohol (ion-donator), and determine the stability constant of the complex from the spectroscopic data.

Data Treatment and Analysis

Using the 0.005M of iodine given in the laboratory, the following sample calculation was conducted to acquire 0.0005M of iodine for solution 2:

Mole of Iodine required in 25ml = 0.025L X 0.0005M = 1.25 X 10-5 moles

Amount of stock solution of 0.005M of Iodine needed = 1.25 X 10-5 mol / 0.005M

= 2.5 X 10-3 L

= 2.5ml

Using the 2M of n-butyl alcohol given in the laboratory, the following sample calculation was conducted to acquire 0.2M of n-butyl alcohol for solution 2:

Mole of n-butyl alcohol required in 25ml = 0.025L X 0.2M = 5 X 10-3 moles

Amount of stock solution of 2M of n-butyl alcohol needed = 5 X 10-3 / 2M

= 2.5 X 10-3 L

= 2.5ml

Amount of cyclohexane was calculated using following sample calculation from solution 2:

Amount of cyclohexane needed= 25ml - 2.5ml of Iodine - 2.5ml of n-butyl alcohol = 20.0ml

Table 1 show the calculated data from the above mentioned calculation

Solution

1

2

3

4

5

6

Amount of Iodine used (ml)

2.5

2.5

2.5

2.5

2.5

2.5

Initial concentration of Iodine (M)

0.0005

0.0005

0.0005

0.0005

0.0005

0.0005

Amount of n-butyl alcohol used (ml)

0.0

2.5

5.0

10.0

15.0

20.0

Initial concentration of n-butyl alcohol (M)

0.0

0.2

0.4

0.8

1.2

1.6

Amount of cyclohexane used (ml)

22.5

20.0

17.5

12.5

7.5

2.5

Total volume (25ml)

25

25

25

25

25

25

Using the UV-VIS spectrophotometer, the absorbance of each solution was recorded as follow in Table 2:

Table 2: Absorbance value for each solution at 520nm, 460nm, 440nm wavelength

Wavelength

Solution 1

Solution 2

Solution 3

Solution 4

Solution 5

Solution 6

520 nm

0.471

0.417

0.373

0.318

0.290

0.263

460 nm

0.101

0.167

0.208

0.265

0.314

0.341

440 nm

0.034

0.095

0.140

0.202

0.254

0.286

The isosbestic point is measured at 492.10nm with the absorbance of 0.323Å

Each wavelength gives different εa and εc unless at isosbestic point, hence, the εa and εc not at isosbestic point have to be calculated separately for each wavelength.

Using the data collected, Y was calculated using the sample calculation shown for solution 1 at wavelength 520nm:

Ysolution 1 = Absorbance (a) for 520nm solution 1 /Initial concentration of ion-acceptor [A]o solution 1

= 0.471/0.0005 M

= 942 M-1

Εa have to be calculated using solution 1. The calculation of for εa at 520nm wavelength is: Absorbance (a) =εa [A]+εc [C] where [C] is the concentration of complexes formed and concentration of donor alcohol does not absorb visible light.

In solution 1, concentration of donor alcohol is absent. No complex formed, [C] = 0

Hence, a = εa[A] where [A] = [A]o - [C]

When [C] = 0, a = εa[A]o ɭ

εa = a/ [A]o ɭ

= (0.471/0.0005 M ) x 1cm

= 942 M-1 cm-1

Using Y and εa, the X value can be calculated using

Where X=

Calculated value for Y, εa and X at 440nm, 460nm and 520nm wavelength can be found in Appendix 3.

Using the Regression function in Excel 2010 with the Add-ins of data analysis function, the coefficient 1/K and the three εc at different wavelengths of 440nm, 460nm, and 520nm was calculated with their standard of error. Summary output of Regression Statistic can be found in Appendix 4.

Table 4: Coefficient for 1/K, εc, 440, εc, 460, εc, 520 calculated

Coefficients

Standard Error

Intercept

0

#N/A

X (εa-Y/[D0])

1.0906

0.0596

εc,440

879.7589

29.2589

εc,460

1010.9579

29.4586

εc,520

233.7889

26.3779

Uncertainties are calculated and can be found in Appendix 5:

K = Ќ ± tSK

= 0.917 ± (2.2 x 0.05) M-1

= 0.917 ± (0.11) M-1

εc,440 = 879.7589 (±64.370)

εc,460 = 1010.9579 (±64.809)

εc,460 = 233.7889 (±58.031)

Discussion

Charge Transfer Complex

Charge transfer (CT) complex is the association product of two or more molecules and fraction of electron charge transfer from a donor to an acceptor. This phenomenon occur when the acceptor of electron have high electron affinity while the donor of electron have low ionization potential. In this experiment the iodine (I2), is the acceptor of electron and n-butyl alcohol (C4H9OH), is the donor of electron. The iodine-alcohol complex formed due to the donation of electrons from the Highest Occupied Molecular Orbital (HOMO) of n-butyl alcohol to the Lowest Unoccupied Molecular Orbital (LUMO) of iodine. For iodine, the LUMO is the anti-bonding (σ*) orbital and for alcohol, the HOMO is the non-bonding orbital. Hence when the electron donating species of n-butyl alcohol approaches the empty orbital of iodine, the HOMO of alcohol will interact to form complex with the LUMO of iodine where the non-bonding electrons in alcohol are filled into a new formed bonding orbital between the two species which is nearer toward the iodine. This can also be explained as a Lewis's acid-base theory of an electron transfer between the Lewis base (electron donator) and Lewis acid (electron acceptor) to form a Lewis adduct.

Colors of Iodine and Iodine-Alcohol Complex

During the experiment, it can be observed that iodine with no complex solution is in violet color and becoming progressively reddish-pink to then to brown when more alcohol donor are added and more complexes are formed. This changing of color is due to the complexes absorb the complentary green light from the visible light spectrum as compare to iodine absorbing yellow complentary light. This implies that the complex has much higher transition energy than iodine and is more stable which shows that the difference in energy levels between the bonding and anti-bonding orbitals of the complex is larger than between the s and s* orbitals of iodine. These tell that there is higher transition energy of the complex. Therefore, the iodine-complex will be of different color as compared to iodine as the complex has higher transition energy and absorbs shorter wavelength of light,

Shifting of Maximum Absorption wavelength (λmax)

In the spectrum attached in the appendix, there is an observation of the shifting of the maximum absorbance wavelength (λmax) to lower wavelength with the increase of concentration of alcohol. For Solution 1 containing only iodine and no donor present, the λmax wavelength is recorded at 520nm. Then, the λmax of solution 2 to solution 6 show a decreasing trend in the λmax wavelength toward the isosbestic point of the spectrum. This shows that adding donor will have effect on iodine and complexation of iodine and alcohol occurred. This shifting of λmax is due to the iodine-alcohol complexes absorb a shorter wavelength of light than iodine as mentioned. Therefore, solution 1 of only iodine always shows the λmax at higher wavelength than solution 2 to solution 6 due to the higher transition energy of the complexes. The shift become more drastic as higher concentration of alcohol added and more complexes are formed which give rise to higher absorbance at shorter wavelength and lower absorbance at higher wavelength of λmax of iodine. Hence, these changes resemble a shift of the absorption maximum.

Isosbestic point

Isosbestic point is defined as the specific wavelength at which two chemical species have the same molar absorptivity (ε) and the overall absorbance of a sample does not change during a chemical reaction or a physical change of the sample. The isosbestic point is taken as a criterion for the existence of two species in equilibrium where the concentration of both species is constant. Hence, in this experiment, the isosbestic point proved that there will be two principal species present in the solution and can be used to determine the total amount of the two species in equilibrium, since the two species have the same molar absorptivity constant at the same wavelength.

Absorbance = εA[A] + εC [C]

If at isosbestic point, εA,iso = εC, iso

Absorbance = εA,iso [A] + εC, iso [C]

Absorbance = εiso ([A] + [C])

Hence, when total concentration of acceptor [A] and complexes [C] is constant, the absorbance will always be constant at the isosbestic point as epsilon εiso is constant. Even when the reaction is not at equilibrium or finish, there will still be an isosbestic point as the total amount of concentration is not changed among the six solutions.

Possible limitations and sources of errors in experiment

In this experiment, there are sources of errors and caution to be taken.

Firstly, all the apparatus used must be dry as water will affect the interaction between the iodine and alcohol by acting as ligand and disrupt the coordination between the iodine and alcohol. This will cause the concentration of the iodine-alcohol complex to deviate greatly from the intended concentration used to calculate the stability constant.

Secondly, the sample solutions are to be left in stable room temperature for 30 minutes so as allow the complexation between the iodine and alcohol to be completed fully. The temperature have to be constant and stable because the stability constant and the equilibrium of the reaction is temperature-dependent where a small changes in temperature will affect the reading of the spectrum greatly.

Thirdly, the use of auto-burette helps to reduce systematic error caused by inaccurate measurement of relying the marking on the volumetric flask to make-up the volume.

Also the cuvette is to be first washed with cyclohexane, the inert solvent, and subsequently washed two times with the sample solution and the less concentrated solutions are to be measure before the higher concentrated solution so as to prevent any carry over effect or residual mixing from the left over in the same cuvette.

Lastly, the volumetric flasks are to be capped once solution are added in and mixed well so as to prevent the volatile n-butyl alcohol to evaporate and affect the concentration of alcohol donor in the volumetric flask and to obtain a homogenous solution.

Conclusion

In conclusion to the whole experiment, the stability constant of the complex between iodine and n-butyl alcohol is 0.917 (±0.0501). The complex has an isosbestic point at wavelength 492.10nm with the absorbance of 0.323Å. The εc, 440, εc, 460, εc, 520 have also been calculated to be 879.7589 (±29.2589), 1010.9579 (±29.4586), 233.7889 (±26.3779) respectively.

References

[1] G. D. Christian, J. E. O'Reilly, Instrumental Analysis, 2e, Allyn & Bacon, 1986.

[2] Atkins, P & dePaula, J. (2006). Atkins' Physical Chemistry (8th ed.). New York: Oxford University Press.

[3] T. Engel and P. Reid, Physical Chemistry, 2nd ed.; Person Prentice Hall, 2010.

[4] IUPAC. Compendium of Chemical Terminology, 2nd ed. (the "Gold Book"). Compiled by A. D. McNaught and A. Wilkinson. Blackwell Scientific Publications, Oxford (1997). XML on-line corrected version: http://goldbook.iupac.org (2006-) created by M. Nic, J. Jirat, B. Kosata; updates compiled by A. Jenkins.

[5] Housecroft, C.E. & Sharpe, A.G. (2008). Inorganic Chemistry (3rd ed.). Harlow: Prentice Hall.

[6] Thomas, M.J.K. (1996). Ultraviolet and visible spectroscopy (2nd ed.). Chichester; New York: Published on behalf of ACOL (University of Greenwich) by J. Wiley.

[7] Rao. C. N. R. (1961). Ultra-violet and visible spectroscopy. Great Britain, Page Bros. (Norwich) Ltd.

Appendices

Appendix 3

At Wavelength of 520

Absorbance

[I2]

εa

Y (a/[A]0)

[D]0

X (εa-Y/[D]0)

Solution 1

0.471

0.0005

942

942

0

N/A

Solution 2

0.417

0.0005

942

834

0.2

540

Solution 3

0.373

0.0005

942

746

0.4

490

Solution 4

0.318

0.0005

942

636

0.8

382.5

Solution 5

0.29

0.0005

942

580

1.2

301.6666667

Solution 6

0.263

0.0005

942

526

1.6

260

At Wavelength of 460

Absorbance

[I2]

εa

Y (a/[A]0)

[D]0

Solution 1

0.101

0.0005

202

202

0

Solution 2

0.167

0.0005

202

334

0.2

Solution 3

0.208

0.0005

202

416

0.4

Solution 4

0.265

0.0005

202

530

0.8

Solution 5

0.314

0.0005

202

628

1.2

Solution 6

0.341

0.0005

202

682

1.6

At Wavelength of 440

Absorbance

[I2]

εa

Y (a/[A]0

[D]0

Solution 1

0.034

0.0005

68

68

0

Solution 2

0.095

0.0005

68

190

0.2

Solution 3

0.14

0.0005

68

280

0.4

Solution 4

0.202

0.0005

68

404

0.8

Solution 5

0.254

0.0005

68

508

1.2

Solution 6

0.286

0.0005

68

572

1.6

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