In the Laboratory
Micelles in the Physical/Analytical Chemistry Laboratory. Acid Dissociation of Neutral Red Indicator
W
Kathryn R. Williams* and Loretta H. Tennant Department of Chemistry, University of Florida, Gainesville, FL 32611-7200; *
[email protected] Because of the widespread scientific and commercial applications of micelles, the undergraduate laboratory curriculum should include experiments involving these systems. Previous contributions from this department (1, 2) have described physical chemistry experiments on diffusion coefficients and hydrolysis kinetics in aqueous and micellar environments. The experiment presented here demonstrates the effects of neutral (Tween-80) and anionic (sodium dodecyl sulfate, SDS) micelles on the acid dissociation equilibrium of neutral red (see structure below). The dissociation constant is determined by a modified spectrophotometric method. H3C
N
H2N
N H
Cl +
N
−
CH3
H3C
protonated neutral red 3-amino-7-dimethylamino-2-methylphenazine hydrochloride
Background The effects of micellar systems on acid/base equilibria have been studied since the first half of the 20th century. The work of Hartley and Roe (3) in the late 1930s attributed apparent pKa shifts of a solubilized indicator to changes in the local activity of H3O+ in the immediate vicinity of a charged micelle. About 20 years later, Mukerjee and Bannerjee (4) proposed that the effects also involve a second contribution due to the lower dielectric constant of the micellar environment. They showed that the pKa in the presence of surfactant, pKaS, is given at 25 °C by pKaS = pKai – Ψ/59.16
(1)
where pKai represents the intrinsic pKa in the dielectric constant of the micellar interface, and Ψ is an electrostatic potential (in mV), which has the same sign as the charged micellar surface (negative in SDS; zero for Tween-80). The latter term determines the local H3O+ activity, which for an anionic micelle is greater than in the bulk water. This forces the dissociation equilibrium to the left and raises the pKa, as shown in eq 1. Subtracting pKaW, the value in a totally aqueous environment, from both sides of eq 1 gives or
pKaS – pKaW = pKai – pKaW – Ψ/59.16
(2)
∆pKaS = ∆pKai – Ψ/59.16
(3)
(zero to ᎑1 or +1 to zero change). Many common indicators (e.g., the sulfonphthaleins) undergo valence changes of ᎑1 to ᎑2, and others (e.g., methyl red) are complicated by multiple equilibria or lack of water solubility. The indicator chosen for this experiment, neutral red (NR), is commercially available in sufficiently pure form and is readily soluble in water. Most importantly, it undergoes a single +1 to zero change: HNR+ + H2O
(4)
In the hydrophobic micellar environment, neutral NR is expected to be stabilized relative to HNR+. Thus, the dissociation should shift to the right leading to a decrease in pKa and a negative ∆pKai. Fernández and Fromherz (5) determined the pKa values of two types of coumarin dyes in several surfactants and showed that ∆pKai in a charged micelle could be evaluated as the ∆pKaS in a neutral surfactant. Thus, if Ψ is known, eq 3 can be used to calculate ∆pKaS in charged micelles. In the experiment, students measure pKa′ (prime denotes 0.10 M ionic strength, as explained below) in water and in micelles of neutral Tween-80 and anionic SDS. They can then compare the experimental ∆pKaS in SDS to the value calculated from the ∆pKaS for Tween-80 and the value of Ψ (᎑100 ± 5 mV) obtained from the literature (6 ). Experimental Method The written material for students reviews concepts of acid dissociation and activity coefficients.W All solutions are adjusted to an ionic strength of 0.10 M with NaCl. Thus, as the student write-up explains, the laboratory results give the negative logarithm of the concentration quotient, Ka′, for the equilibrium in eq 4, valid at µ = 0.10 M. As part of the data analysis, students use activity coefficients from the extended Debye–Hückel equation to convert K a′ for the totally aqueous system to the corresponding activity quotient. (This calculation is not possible for the surfactant systems owing to lack of appropriate activity coefficient data). The spectrophotometric determination of an acid dissociation constant is described in laboratory manuals (7, 8) and in previous articles in this Journal (9, 10). The absorbance is measured for solutions with the same total indicator concentration at low pH (Aacid), high pH (Abase) and a series of pCH values (᎑log[H3O+]) in the transition interval (ApcH). A plot of
log
The sign of the ∆pKai term depends on the charge states of the conjugate forms of the indicator. The effects of the micellar environment are best observed with a weak acid having a single ionization involving a neutral conjugate form
NR + H3O+
A p H – A acid NR c = log A base – A p H HNR+
(5)
c
versus pCH should have a slope of unity and a y intercept of ᎑ pKa′. To avoid errors due to a nonideal slope, students are instructed to use the x intercept, where [NR] = [HNR+] and pCH = pKa′.
JChemEd.chem.wisc.edu • Vol. 78 No. 3 March 2001 • Journal of Chemical Education
349
In the Laboratory
Table 1. Spectral Data for Neutral Red λ/ nm
System
2.07 × 104
2.93 × 103
0.02 M Tween-80 460
5.38 × 10
4
2.65 × 10
1.08 × 103
0.02 M SDS
4.64 × 103 1.27 × 104 542
2.87 × 104
1.33 × 103
453
458
3
Previous experimental designs have specified preparation of a series of solutions in volumetric flasks with constant indicator concentration and pCH values set by a nonabsorbing buffer (plus the two solutions for Aacid and Abase). At the University of Florida, students use a procedure modified for the Spectral Instruments 440 fiber-optic spectrophotometer. They prepare a solution of the indicator in the protonated buffer (NH2OH⭈HCl, pKa′ = 5.971 at µ = 0.10 M, for the aqueous and Tween-80 systems; H3BO3, pKa′ = 9.090 at µ = 0.10 M, for SDS [Williams, K. R. unpublished results]) and add known increments of standard NaOH to gradually raise the pCH. The absorbance at each increment is multiplied by the ratio of the total volume to the initial volume to correct for dilution. This procedure allows more data points to be acquired with less waste of time, glassware, and reagents. If a fiber-optic instrument is not available, the individual solution method can be used with a spectrophotometer having a standard cell holder.
Table 2. pKa´ Values for Neutral Red ( = 0.100 M with NaCl) pKa′ Exptl
⑀NR/ ⑀HNR+/ L mol᎑1cm᎑1 L mol᎑1cm᎑1
5.36 × 103 3.39 × 103 528
Aqueous
System
λ/ nm
⑀HNR+/ ⑀NR/ L mol᎑1cm᎑1 L mol᎑1cm᎑1
a
Lit. (12)
∆pKaS Exptl
Aqueous
6.66 ± 0.13
6.54b
—
0.02 M Tween-80
5.66 ± 0.10
5.69 c
᎑1.00 ± 0.16
0.02 M SDS
9.03 ± 0.17
9.17 d
+2.37 ± 0.21
aUncertainties
given as 95% confidence limits. bCalculated from the thermodynamic pK . a c5% (0.04 M) Brij-35, another neutral surfactant. d2% (0.07 M) SDS.
1.51 × 10
4
542
Hazards Hazards are minimal, because the students handle only dilute (0.10 M or less) aqueous solutions, which are prepared for them. The person who makes the solutions should use usual care to avoid injestion or skin contact with the reagents. A mask may be needed to prevent inhalation of fluffy SDS powder. Results Table 1 presents spectral data for neutral red in the three systems, and a typical plot of log[NR]/[HNR+] versus pCH is shown in Figure 1. The pKa′ values, presented in Table 2, agree well with published data. The signs of the ∆pKaS values agree with the predictions described above. Neutral NR is stabilized in the hydrophobic micellar environment of Tween-80, resulting in a ∆pKaS of ᎑1.00 ± 0.16. Using this value for ∆pKai and ᎑100 ± 5 mV for Ψ, eq 3 gives a theoretical value of ∆pKaS for 0.020 M SDS of +0.69 ± 0.18, which is obviously outside the error bars for the observed ∆pKaS of +2.37 ± 0.21. Drummond et al. (11) obtained a similar result and considered several possible sources for the discrepancy. They concluded that ∆pKaS in neutral micelles does not adequately represent ∆pKai for neutral red and other acids containing a protonated heterocyclic nitrogen. There must also be a specific molecular interaction, either ion-pair formation or hydrogen bonding to the sulfate oxygen, between the cationic indicator and the anionic headgroups. This information is given in the written material for students,W but some faculty may wish to withhold the explanation until students observe the discrepancy experimentally. Conclusion The spectrophotometric determination of pKa reinforces students’ knowledge of acid/base equilibria, activity considerations, and applications of spectrophotometry. The large pKa shifts produced by incorporation of neutral and negative micelles are readily observed and can lead to interesting discussions of LeChâtelier’s principle and the causes of unexpected experimental results. The experiment is suitable for upper-level physical chemistry and instrumental analysis laboratories, as well as honors courses in introductory analytical chemistry. W
Figure 1. Typical plot of log([NR]/[HNR+]) vs pcH for neutral red in 0.020 M Tween-80 (µ = 0.10 M); 䊉, 542 nm; 䊏, 460 nm.
350
Supplemental Material
The full version of this article, with detailed procedures and notes for the instructor, is available in this issue of JCE Online.
Journal of Chemical Education • Vol. 78 No. 3 March 2001 • JChemEd.chem.wisc.edu
In the Laboratory
Literature Cited 1. Williams, K. R.; Bravo, R. J. Chem. Educ. 2000, 77, 392–394. 2. Williams, K. R. J. Chem. Educ. 2000, 77, 626–628. 3. Hartley, G. S.; Roe, J. W. Trans. Faraday Soc. 1940, 36, 101–109. 4. Mukerjee, P.; Bannerjee, K. J. Phys. Chem. 1964, 68, 3567–3573. 5. Fernández, M. S.; Fromherz, P. J. Phys. Chem. 1977, 81, 1755–1761. 6. Hartland, G. V.; Grieser, F.; White, L. R. J. Chem. Soc., Faraday Trans. 1 1987, 83, 591–613. Students use a total ionic strength of 0.10 M. The Ψ value of ᎑100 ± 5 mV was interpolated from
7.
8. 9. 10. 11.
the data in the reference for SDS at ionic strengths of 0.065 M (᎑110 ± 5 mV) and 0.102 M (᎑ 95 ± 5 mV). Albert, A.; Serjeant, E. P. The Determination of Ionization Constants: A Laboratory Manual, 3rd ed.; Chapman & Hall: New York, 1984; Chapter 4. Sawyer, D. T.; Heineman, W. R. Chemistry Experiments for Instrumental Methods; Wiley: New York, 1984; pp 193–197. Tobey, S.W. J. Chem. Educ. 1958, 35, 514–515. Forst, W. J. Chem. Educ. 1959, 36, 289–290. Drummond, C. J.; Grieser, F.; Healy, T. W. J. Chem. Soc., Faraday Trans. 1 1989, 85, 551–560.
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