In the Laboratory
Surface pKa of Self-Assembled Monolayers
W
Penny S. Hale,*† Leone M. Maddox, and Joe G. Shapter School of Chemistry, Physics, and Earth Sciences (SoCPES), Flinders University of South Australia, GPO Box 2100, Adelaide, SA, 5001 Australia; *
[email protected] J. Justin Gooding School of Chemical Sciences, The University of New South Wales, Sydney, NSW, 2052 Australia
is vital in increasing the understanding of acid solutions and aids in providing a framework for the theory to be presented. The equation commonly used in the buffer region of the acid–base equilibrium is the Henderson–Hasselbach equation: pK a = pH − log
Theory
[A− ] [HA ]
(1)
This equation is based on the following acid equilibrium with water:
HA(aq) + H2O(l)
H3O+(aq) + A−(aq)
(2)
The titration method is commonly used to determine the solution pKa (3, 6). This involves monitoring the pH of a solution as base is added by titration. A plot of pH versus volume yields the familiar titration curve (Figure 1). When [A−] and [HA] are equal, it follows from eq 1 that the pH = pKa. In practice, this is measured as the pH at a point halfway to the stoichiometric end point. A theoretical argument expanding on this point is presented in many textbooks (7). 3- Mercaptopropionic acid (HSCH2CH2COOH, abbreviated as MPA) establishes self-assembled monolayers (SAM) on gold surfaces and has seen wide application in electroanalytical sensors (8). In a MPA SAM on a gold electrode the terminal carboxylic acid group of the SAM exhibits pH dependence. The carboxylic acid is fully protonated at low pH (below about pH 4). Therefore there are no electrostatic repulsive interactions between the SAM and the anionic fer-
10 9 8 7
pH
Understanding the way in which a surface responds to the solution environment allows the production of surfaces with controllable properties. The controllable (and reversible) activation or passivation of a surface is a key component of a number of different systems such as catalysis, biosensing, cell culture, and related areas. As the surface area-to-volume ratio become greater (such as in nanotechnology applications), the surface properties become increasingly important. The surface pKa of an acid may differ from the solution pKa. This may have large consequences on the surface sensitivity and selectivity in different solutions. Different methods such as contact angle titration (1), chemical force titration (using an AFM), quartz crystal microbalance (2), surface potential, electrochemical desorption (3), and spectroscopy (4) have been used to determine surface pKa. Recently a simple method was proposed by Zhao et al. (1) that is easily adapted to the undergraduate teaching laboratory. Only buffer solutions with different pH values, an anionic electrochemical probe to investigate the dissociation of a surface carboxyl group, and a standard potentiostat are required. The method involves forming a self-assembled monolayer of mercaptopropionic acid on a gold electrode and running a series of cyclic voltammograms of the ferricyanide–ferrocyanide redox couple in different pH buffers. Measuring the peak current (ip) for the couple as a function of pH allows the determination of the pKa of the surface. The experiment is used in the second year of a nanotechnology degree at Flinders University (5). The main goal of the experiment is to illustrate the difference between solution and surface properties such as pKa. This idea is essential in developing a strong understanding of the differences between nanoscale and macroscopic systems. A secondary goal is to become familiar with electrochemical instrumentation, such as a potentiostat, and to be comfortable using techniques such as cyclic voltammetry. This technique has the advantage that different aspects such as the dependence of the peak current on the scan rate of reversible redox reactions can be studied. Students are able to link theory and experiment to draw conclusions about a system.
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The theory of acids and bases is well established and the details are readily available from any chemistry textbook but the key point for students is how to get information from the equations provided and how to interpret the graphs in terms of the theoretical equations. A laboratory experience † Current address: Centre for Materials and Surface Science, Department of Physics, Faculty of Science, Technology and Engineering, La Trobe University, Victoria 3086, Australia.
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4 3 2 0
5
10
15
Volume of Base / mL Figure 1. Titration curve of a weak acid (MPA) against a strong base (NaOH).
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In the Laboratory
4
2
Current / µA
tion 4 is derived in the Supplemental Material.W) Hence, a plot of peak current i versus pH can elucidate the pKa of a surface. The secondary goal of this laboratory was to introduce cyclic voltammetry (CV) as an electrochemical technique to probe surface properties. The idea that electrochemistry can probe surface properties has been illustrated before (10) and is common in research but not in undergraduate programs. Students were required to perform a number of standard CV tests such as proving the reversibility of a reaction by measuring peak potential and peak current as a function of scan rate. They were also required to estimate the area from the Randles–Sevcik equation,
A B C D E F G
0
-2
-4
-6 -250
-50
150
350
550
750
(
Figure 2. Cyclic voltammograms of ferricyanide in 0.2 M PBS at the SAM-modified electrode over a range of pH values: A = 3.37, B = 4.05, C = 4.20, D = 4.46, E = 5.08, F = 6.44, and G = 7.98. Scan rate 100 mV/s.
ricyanide and a good electrochemical response is recorded at the electrode. As the pH increases the carboxylic acid groups begin to deprotonate (9) and hence the increased electrostatic repulsive forces result in a decrease in current response. At high pH (above pH 6), the carboxyl groups in the SAM will be fully dissociated. As the repulsive forces between the SAM and the ferricyanide will be at a maximum, the current response at the electrode will be at a minimum. Hence the variation in peak current with solution pH can be used to determine the pKa. The charge transfer is inhibited under conditions where the monolayer is negatively charged (i.e., high pH) and hence the degree of inhibition will increase as the charge on the SAM carboxyl group increases. This results in a change in the shape of the cyclic voltammogram such as a decrease in peak current and an increase in peak separation with pH, as shown in Figure 2. The method for determining surface pKa uses the following argument: if the total current through the electrode is assumed to be composed of two independent parts, one through the dissociated SAM [A−] and the other through the non-dissociated SAM [HA], then the current can be described by i = i A − [A −] + iHA [HA ]
(3)
where iA− and iHA are currents of the probe on the SAM fabricated by [A−] and [HA] structures, respectively. In this case, [HA] refers to surface concentration or coverage, not solution concentration. By setting the monolayer coverage to 1, the coverage of the surface components is [A−] + [HA] = 1. Using this expression, plus eqs 1 and 3, the following equation can be obtained (1), pK a = pH − log
iHA − iA −
−1
i − iA −
(4)
where iA− and iHA can be determined by the average values of the high pH (> 6) and low pH (< 4), respectively. (Equa780
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3
Ip = − 2.69 × 105 A n 2 C D
(Potential vs Ag/AgCl) / mV
1
2
ν
1
2
µA
(5)
where Ip is the peak current, C is the bulk concentration (mol兾L) of active species, ν is the scan rate (V兾s), D is the diffusion coefficient (cm2兾s), A is the electrode area (cm2), and n is the number of electrons involved. The students then compare the estimated area to their measured geometric area. Materials and Methods All chemicals were supplied by Aldrich. The buffer solutions were made from 0.2 M phosphate (Na 2HPO 4, NaH2PO4) and 0.01 M KCl, 0.5 mM K3Fe(CN)6 in a range of pHs between 2 and 9. At least eight solutions were made with more solutions between pH 4 and 6 (around the pKa value). The SAM was made by immersing a clean gold electrode in 0.1 mM MPA in ethanol and leaving it submerged overnight.1 (This step was completed for the students by the laboratory staff.) The titrations were performed on 1 mM MPA in water against 5 mM NaOH. The reference was a saturated Ag兾AgCl reference electrode supplied by BAS. The gold electrode was also supplied by BAS and was a 1-mm wire set in epoxy with a typical roughness factor of 1.5. The surface area of the electrode was measured using a ruler or vernier calipers before cleaning thoroughly. The gold electrode was polished thoroughly with wet and dry sandpaper before rinsing with ethanol and deionized water. It was then immersed in fresh piranha solution2 for two minutes and rinsed again. The electrochemical setup used a three electrode cell and a BAS 100B potentiostat. Cyclic voltammetry of the clean gold electrode was performed to show the reversibility of the ferricyanide redox reaction and to prove the pH independence of clean gold. This is achieved by measuring the peak current at different scan rates and at different solution pH. The former should be dependent on the square root of the scan rate and the latter should be independent. Peak potential versus scan rate should also be independent. Ideally, the above step would be done on the same gold electrode on which the SAM would be formed. To achieve this, the electrode would have to be cleaned and analyzed as the final step in the experiment, not the first step. It is presented in this format because this sequence follows more closely the rationale behind the experiment. Alternatively, a separate gold electrode could be included and analyzed prior to the instructor-prepared SAM on a second gold electrode.
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In the Laboratory
Hazards MPA is toxic. Gloves and safety glasses should be worn and the chemical should be prepared in a fume hood. Since the chemical is pre-prepared as a dilute solution for the students, they have minimal exposure and are not at risk. Piranha solution2 reacts violently with organic material and has been known to explode when stored in closed containers. Safety precautions such as wearing protective eyewear and gloves should be adhered to. Only small quantities should be prepared at any one time (about 10 mL). The solution should be disposed of by diluting with copious quantities of water and flushing into the fumehood sink. Results and Discussion As shown in Figure 1, a typical titration curve for a weak acid against a strong base results in an end point that is slightly basic, between pH 7 and 9. The solution pKa therefore is between 3.5 and 4.5. Since the literature value is 4.85 (1) this allows for some error in the technique, which students can discuss in relation to the theory of acids and bases and the assumptions made when estimating from the graph. The range of CVs for the SAM is presented in Figure 2. The peak current from each of the CVs can be determined and plotted against pH as shown in Figure 3. This plot looks similar to the equivalence part of a titration curve, and if a theoretical curve is fitted to the graph, using eq 4, the pKa can be elucidated. This results in a surface pKa of 6.0 ± 0.5, which is comparable to the literature value of 5.6 ± 0.1, for surfaces and 4.85, for solution (1). The students will notice that the more solutions they have between pH 4 and 6, the better their result and the smaller the corresponding error range will be. This laboratory can be accompanied by discussion on surface properties and methods of measuring them. It was incorporated easily into the existing course on nanotechnology, which includes other surface analysis methods such as www.JCE.DivCHED.org
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2.4 2.2 2.0
I p / µA
A series of CVs were performed on the SAM on the gold electrode in a range of buffers from pH 2 to pH 9. The solutions were degassed with nitrogen for 10–15 minutes before each scan to stop oxidation of the SAM. At least eight different solutions were tested and their pH remeasured after each scan. It was important to use more solutions about the region of the pKa and fewer solutions around the end points (conc acid or conc base). At one pH the electrode was scanned at a range of different scan rates to show the independence of the reversible peak potential and the dependence of the peak current on the scan rate. The area can also be measured at this point by knowledge of the diffusion coefficient for ferricyanide and compared to the geometric area measured previously. Students worked in pairs with one person in control of the CV and the second doing the titration. The titration followed the method of Cawley (6). Since the solutions for the CV had to be degassed, this allowed plenty of time for both students to be involved in each experiment and for discussions to take place about each technique and the validity of their results. Students could also swap places and so both became comfortable with the electrochemical apparatus.
1.8 1.6 1.4 1.2 1.0 0
2
4
6
8
10
pH Figure 3. Relationship between the peak current and the pH value of the solution.
contact angle and ellipsometry techniques. However, it would also fit into any physical chemistry course as a basis for understanding acids and bases or as part of a surface analysis course similar to that developed at Michigan State University (10). Using other techniques, such as contact angle to determine pKa can also extend the experiment and give independent results for the same constant. W
Supplemental Material
Instructions for the students, notes for the instructor, and the derivation of eq 4 are available in this issue of JCE Online. Notes 1. MPA is light sensitive; solutions should be stored in the dark. The MPA-modified gold surface should be stored in the dark and under vacuum or nitrogen to prevent oxidation of the monolayer as well. 2. Piranha solution is 30% H2O2 and conc H2SO4 mixed in a 1:3 ratio.
Literature Cited 1. Zhao, J.; Luo, L.; Yang, X.; Wang, E.; Dong, S. Electroanalysis 1999, 11, 1108–1111. 2. Sugihara, K.; Shimazu, K. Langmuir 2000, 16, 7100–7105. 3. Munakata, H.; Kuwabata, S. Chem. Comm. 2001, 1338–1339. 4. Yu, H. Z.; Xia, N.; Liu, Z. F. Anal. Chem. 1999, 71, 1354– 1358. 5. Shapter, J. G.; Ford, M. J.; Maddox, L. M.; Waclawik, E. R. IJEE 2002, 18, 512–518. 6. Cawley, J. J. J. Chem Educ. 1993, 70, 596–598. 7. Atkins, P. W.; de Paula, J. Physical Chemistry, 7th ed.; Oxford University Press: Oxford, 2002. 8. Gooding, J. J.; Hibbert, D. B. Trends Anal. Chem. 1999, 18, 525–533. 9. Kim, K.; Kwak, J. J. Electroanal. Chem. 2001, 512, 83–91. 10. Severin, K. G.; Blanchard, G. J.; Bruening, M. L. Book of Abstracts, 218th ACS National Meeting, New Orleans, Aug. 22-26: CHED-245, 1999. http://poohbah.cem.msu.edu/courses/ cem419/experiments.html (accessed Feb 2005).
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