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ADSORPTION OF MOLECULES OF BIOLOGICAL IMPORTANCE AT THE GOLD-
SOLUTION INTERFACE
A Thesis
Presented to
The Faculty of Graduate Studies
of
The University of Guelph
by
HONG-QXANG LI
In partial fuU3hent of requirements
for the degree of
Doctor of Philosophy
August, 2000
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ADSORPTION OF MOLFCULES OF BIOLOGICAL IMPORTANCE AT THE
GOLD-SOLUTION INTERFACE
Hong-Qiang Li
University of Guelph, 2000
Supervisor:
Professor Jacek Lipkowski
The electrochemical and subtractively normalized interfàcial Fourier transfomi
i h e d spectroscopy (SNIFTIRS) techniques have been ernployed to study the
interfacial behaviours of benzoate, phenyManîne (Phe), and 1,S-dimy&y1-m-glycero-3phosphocholine (DMPC) at the Au(l11)-solution interface. Cyclic voltammetry (CV)and
differentiai capacity @C) have k e n used to qualitatively describe the characteristics of
the adsorption and desorption processes. The Gibbs excess
and Gibbs energy of
adsorption (AG) have been obtained fiom the quantitative analysis of the charge
densities, determined by use of the chronocoulometric technique. The results demonstrate
that, although benzoate adsorption starts at negative charge densities, it takes place
predominantly at a positively charged surface. At the most positive potentials, the Gibbs
excess of beazoate attains a limiting value of about 7.3 x 10-'O mol cm-'. This value is
consistent with packing density corresponding to a close-packed monolayer of vertically
adsorbed benzoate molecules. The adsorption of Phe is a three-state process. The
maximum surface concentration is determined to be about 2.0 x 1 0 " ~mol cnf2 in our
experiment. The calculated values of AG for Phe range nom -1 8 to -37 kJ mol-', which
are characteristic for weak chemisorption of small aromatic molecules. The
corresponding values of AG for benzoate are in the range of 4 4 to -51 kJ mol-', which
are much higher than those of Phe.
The SNIFTIRS studies uidicate that both benzoate and phenylalanine fàvor a fiat
orientation (z-bonding) at negatively charged surfiaces and a vertical (tiited) orientation at
more positively charged surfaces with the carboxylate groups attached to the electrode
&e-
N o dissociation is seen for benzoate while the decarboxylation on Phe is
observed at E > 300 mV. The integrated IR band intensities for both molecules agree well
with the Gibbs excess values determined Fom electmchemicai measurements.
A method of direct vesicle k i o n technique has been developed to deposit
phospholipid bilayers onto a gold electrode Surface. In addition, a modifïed LangmuirBlodgett (LB) technique has been used to deposit the DMPC bilayers. The properties of
the bilayer deposited by vesicle fiision were compared to the properties of the bilayer
formed v i a LB deposition. Atomic Force Microscopy @FM) was used to determine the
thickness of the bilayer. The film pressure of DMPC bilayer at the gold-solution interface
is determined to be 60 rnN 6'.The bilayer on gold shows very good barrier properties
although it has some defects. SNLFTLRS experirnents were performed to determine the
orientation of DMPC molecules in the bilayer. The bipolar shape of -P02- stretching
bands in the SNIFTIRS s p e c t m indicates that the DMPC bilayer displays a tail-to-tail
configuration with the polar heads t m e d towards the electrode surface. The peak
positions of h t h CH2 stretching and C=O stretching bands show Little variation with the
electrode potentials although their intensities increase as the electrode potential is made
more positive. The I
R intensities of CH2 symmetric stretches agree well with the charge
densities in the potential range investigated.
ACKNOWZEDGMENTS
1would like to thank rny s u p e ~ s o ProE
r
Jacek Lipkowski for his guidance and
help during the course of my PbD. program. I gained s lot trom his knowledge and
wisdom over the k t few years.
1 wish to thank other advkory and examination committee members, Prof. M.
Baker, Prof. M. Cocivera, Prof. M. McDermott, Prof. J. Pawkzyn, Prof- D. Thomas and
the chair, Prof. P. D. Josephy for their valuable tirne on rny research and thesis.
1 wish to thank Prof. S. Roscoe, Prof. AR. Merrill, Dr. S. Horswell and Mr. J.
Rubinstein for collaborations and discussions.
I wish to thank all my fi-iends and colleagues in the research group, Dr. G.
Szymanski, Dr. S. Wu, Dr. A.. Chen, V. Zamlyny, N. Li, 1. Burgess, T. Zhao, X. Cai. V.
Lee, C. Hill, Dr. A. Lachenwitzer and Ms. M. Dymarska.
I wish to thank all the supporting staff in the department, especially Mr. T. White,
Mr. Y. Savoret, Mr. B. McCluskey and Ms. K. Shiell.
Finally, 1 would like to express my deep gratitude to my wife Huiping, my parents
and other family members for their continued encouragement and suppon
Table of Contents
--
Acknowledgements
Table of Contents ---List of Fiwes --------------List of Tables
-
C _ - - - - - P - U I _ _ _
*---------------
------
---
List of Syrnbols and Abbreviations -------
----------------------__q_---_U_-------------
-CII_-------------
(9
---O
69
(+)
(24
List of Figures
Figure 1-1
--
0
Singer-Nicohn fiuid mosaic model for biomembrane structureFigure 1-2 --------------------------
Pl
Hypothetical structures of traasient and metastable membrane conformations that are
believed to be relevant to electroporation
Figure 2-1 ----------------------------------------
-
04)
Mode1 of the elec?rical double layer at the metal/soIution interface.
--
Figure 2-2
-(28)
____C__--
The surface electric field E/E, for plathum at (a) v = 500 cm-' and (b) v = 2000 c*' as a
fhction of the angle of incidence 0.
Figure 2-3 -------------
---- --------
----- ------
-------
@)
Instantaneous dipoles appearing on an adsorbed diatomic molecule oriented
perpendicular and paraliel to a metal surface.
Figure 2-5
----
----------.------
-------------------
--- -------------- (3%
Principles of subtractively nomiized interfaci d FIlR spectroscopy.
spectroscopy.
Figure 2-7--------------------------------------------------------------
(46)
Apparatus designed to deposit a bilayer of phospholipid on a planar substrate.
Figure 2-8
----------------------
(39
Demonstration of the Horizontal Touch technique. (a) The substrate is lowered to contact
the monolayer at the gas-solution interface, (b) The substrate is lifted horizontdy, (c)
The lipid is transferred to the substrate with non-polar tail attached to the substrate
surface, (d) Double horizontal touch (not a biomembrane model), (e) Lipids turned over
(possible?) to form the bilayer on the substrate during the double horizontal touch.
Figure 2-9
-
@O)
Proposed mechanism for the structural reorientation of a phospholipid monolayer on a
mercury surfêce fkom heads-down to tails-down fashion
-
Figure 2-10 --
--- (53)
Mechanisms of vesic1e spreading on substrates, (a) Rolling. (b) Sliding.
Figure 2-11
-
--
a
-
(58)
Impedance analysis of a supported bilayer deposited by vesicle fusion. AU dots are
experimentd data. The lines are the fïtting results of their corresponding RC models.
Figure 2-12
-_-
------------
(61)
Different structures of amphiphile aggregates in solution. Those structures can transfonn
fiom one to the other by changing the solution conditions.
Figure 2-13
--
------
------------
Graphic representation for the physical parameters of an amphiphile.
Pigure 2-14 ------------------------------ ----
A schematic procedure (a+b+c)
(63)
O
to make DMPC bilayers by LB deposition and
horizontal touch techniques.
Figure 3-5-------------------------------------
--------------------------
m
(a) A family of current transient curves measured by chronocoulometry. (b) A family of
charge density curves calculated fiom (a).
--
Figure 4-1 -----
Cr011
Differential capacity recorded at an Au(ll1) electrode in 0.1 M KC104(dotted line) and
in 0.1 M KC104 + 1.58 x 10-~M potassium bemoate (solid h e ) . The inset shows the
double layer section of CV m e s recorded in the same solutions at a sweep rate of 20
mV s-' .
Figure 4-2
------
---------------
--- ---------
(iw)
Charge density verstrs electrode potential plots for the Au(ll1) electrode in O. 1 M KC104
solution (dotted h e ) containing different molar concentrations of benzoate: 1.28 x IO-'
(filled circle); 2.56 x 10-'(open ckcle); 4.00 x 10-'(fiUed square); 6.66 x 10-~(0pen
square); 9.30 x 1oe5(filleduptriangle); 1.59 x lo4(open uptriangle); 2.86 x 10~(n11ed
down-triangle); 5.56 x 104(open down-triangle); 5.79 x 104(fUed diamond); 1-58 x 1o - ~
(open diamond); 2.9 1 x 10;'(£illed hexagon); 4.21 x 1oJ(open hexagon).
Figure 4 3 ------------------------------
---
-----
@os)
Dinerential capacity curves for the Au(ll1) electrode in 0.1 M KC104 + 1-58x 10;' M
benzoate solution determined fiom: ac impedance expriment (square); CV curve
(triangle); differentiationof the charge density cuve (circle). The insert shows the charge
density determined fiom: integration of the single-fiequency differential capacity curve
(square); integration of CV curve (triangle) ;chronocoulometry (circle).
Figure 4-4 --------------------------------------------------
--
008)
Plots of surface excess of benzoate on the Au(l11) electrode against (a) the potential and
(b) electrode charge density for 0.1 M KC104 with different molar concentrations of
benzoate: 1.28 x 10'~(nlledcircle); 2.56 x 10-'(open circle); 4.00 x 10-'(filled square);
6.66 x 10-'(open square); 9.30 x 10-~(nlleduptriangle); 1.59 x 104(open u p - k g l e ) ;
2.86 x lo4 (nlled down-triangle); 5.56 x 1 0 ~ ( o ~ edown-triangle);
n
5.79 x l ~ ~ ( f i l l e d
diamond); 1.58 x 105 (open diamond); 2.91 x loJ(filled hexagon); 4.21 x 10-~(0~en
Fitting the benzoate adsorption data to the "square root " isotherm (Eq.3) at (a) constant
potentials and (b) constant o ~ .
----
Figure 4-6
(iw
Gibbs energy of adsorption for benzoate on the Au( 1 11) electrode in 0.1 M KC104: (a)
AG vs E plots detenmined fkom Fig. 5(a) (square) and fiom Fig. 6(b)(circle); (b) A G vs GM
plot determined from Fig. 5(b).
Figure 4-7 ------------
-__U____----__-
vs ï of benzoate at constant electrode potentials/rnV (SCE).
Plots of the
Figure 4-8 --
-
_
_
L
_
_
-
--------
P
(i14)
Electrosorption vdencies determined fiom: (i) the dope of the
(3AG
(113)
GM
vs
r plots (circle);
vs E pIots (square) using AG values -determined f?om the analysis at constant
potential;
(3A G vs E
plots (triangle) using AG values determined fiom the analysis at
constant charge.
Figure 4-9
---
------------
-------
(im
Esin-Markov plots (E vs -RT Zn c) for benmate on the Au(ll1) electrode at charge
densities s h o w in the figure.
Pigure 4-10
---------
-----------------------
------------
(117)
The cornparison of Esin-Markov coefficients determined fiom the slope of the fitted lines
------.--
in Fig. 9 and by the differentiation of the I'vs o~ plots. The dotted line has unity slope.
----------------------
Figure 4-1 1
(120)
b e r layer capacities of the Au(ll1) electrode in 0.1 M KC104 containing different
molar concentratiom of benzoate: 5.79 x 104 (circle); 1.58 x 1oe3(square); 2.91 x 10'~
(up-triangle);4.2 1 x 1
Figure 4- 12 -----------
(down-triangle).
-----
--------------------
(121)
' ~the
) charge of adsorbed
Plots of the potential drop across the inner layer ( ~ 4 ~ vs
at constant charge density on the metal side of the intsrfàce as indicated in
benzoate (OIT)
the figure.
Figure 4-13
-------------
--
022)
Imer Iayer capacities at constant charge (rC) detennined from Fig. 12 (circle) and at
constant amount adsorbed (,C) detemiined from Eq. 8 (square) for the adsorption of
benzoate on the Au(1 Il).
vii
Figure 4-14 -
-024)
Plots of electrosorption valencies determined fkom the
vs
r plots (circle) and fkom the
ratio of ,C to rC (square).
Figure 4-15
---
---
---
W)
The smfhce dipole moment formed by benzoate (square) and CI-(circle) adsorbed on the
Au(ll1) electrode.
---------------
Figure 5-1 -------
(334)
Differential capacity (a), charge density (b) and Gibbs excess of benmate (c) a . the
Au(ll1) electrode d a c e : dotted line, aqueous solution of 0.1 M
K1o4;solid line,
aqueous solution of 0.1 M KC104 + 9.04 x 1o - M
~ benzoate.
Figure 5-2
--
------
(13Q
Transmission IR spectra of D l 0 (dotted line) and Hz0 (solid line).
Figure 5-3 ----------------------------------------Comparison of the transmission FTIR spectrum of benzoate in D2O and the S
038)
m
spectra of benzoate in DzOand in H20. Transmission spectnim was recorded for a D20
solution saturated with the benzoate. SNLFTIR spectra were recorded in O. 1 M KC104 +
9.04 x 10-~
M benzoate solution using Ei=750mV and Ez= 400 mV (SCE).
Figure 5-4 ----------------------------------------------------------
041)
SNIFTE spectra of CO2- asymmetric stretch and C-C stretch for benzoate adsorkd on
the Au(ll1) electrode fiom 0.1 M KCIO4 + 9.04 x 10'~M benzoate solution The spectra
were acquired using (a) s-polarized and (b) p-polarized S a r e d Light for a series of
whose values are indicated in the figure. The reference potential
sample potentials (E2)
EI=7SOmV (SCE).
Figure 5-5 ------ ------------Comparison of the
IR
------
---- -------------------- ----------------
042)
band intensity at 1547 cm-' (filled circles) with the surface
concentration of benzoate (solid line and trïangIes) Obtained by thermodynarnic analysis
of chronocoulometric data [l].
Figure 5-6 -------------------------------------------
---------------(144)
The SNIFTIR spectrurn Hz0 of 0.1 M KC104 + 9.04 x 1o5 M benzoate solution in H20
for the Au(ll1) electrode recorded in the Eequency range extended to 2400 cm-';
reference potential -750 mV and sample potential600 mV(SCE).
-
Figure 5-7
--
046)
SNIFllR spectra of C a - symmetric stretching mode for benzoate adsorbed on the
Au(ll1) electrode fiom 0.1 M KC1O4 + 9.04 x 1o5 M bernate solution using: (a) s- and
(b) ppolarized infi-ared beam and a variable sample potential (Ez) indicated in the figure.
Figure 5-8 ------------__---_____--_-_--------(148)
Dependence of the fiequency of the negative lobe of the bipolar band of the symmetric
CO; stretch on the electrode potential.
Figure 5-9
050)
u
I
_
-
Ratio of the integrated intensities for the SNIFTIRS bands acquired using p and s
polarized radiation; circles, bl band at 1390 cm-'; diamonds, ai band at the 1547 cm-'.
(a) Figure 5-10 -----------------------(153)
Two extreme surface coordiantions of benzoate on a Au(l11) electrode.
(b) Flat orientation; (b) Vertical orientation
Phe solutions (solid he), (a) cyclic voltammetry curves recorded using the sweep rate
20 mV s". (b) dserential capacity recorded ushg a 5 mV (rms) sine wave modulated at
25 Hz and a 5 mV s-' sweep rate.
Charge density c w e s determined fiom chronocoulometry in 0.1 M KC104(dotted h e )
and, for cIarity, a selected set of concentrations of Phe are shown ranging fkom
1.83 x 10" M to 1.06 x 10-~
M.
Figure 6-3 --------------------------
-- ------------------- ---------- -----------064)
Gibbs excess vs potential plots for Phe on Au(l11) at 1-26x 105 M (circle); 2.27 x 10-~
M (square); 3.52 x 10-~M (up-triangle); 6.36 x 13" M (dom-triangle); 1.O6 x 1 0 - ~
M
(diamond). Inset, for 1.06 x10-' M Phe solution, cornparison of the Gibbs excess and the
integrated intensities of the 1574 cm-' (open triangle) and 1410 cm-' (open circle)
SNIFTRS bands in the spectrum of Phe in D2O recorded using s-polarized photons.
Figure 6-4 ---------------------------------------------------------------------066)
For Phe on Au(1 11), Gibbs energy of adsorption and RTln ( cm5/55.5 ) plotted vs the
electrode potential.
Figure 6-5
---
-
--
------(169)
Transmission spectra of Phe and reiated compounds; (a) spectrum of Phe in a KBr pellet,
@) spectnrm of Phe in neutral solution of 90,pH=7, (c) spectnim of Phe in alkaline
solution of D20, pH=ll ,(d) spectnun ofphenylacetate in W .
---------
Figure 6-6 ---------
P
U
-
071)
Transmission spectra of; (a) Phe in alkaline solution of &O, p H 4 1, (b) Phe in acidic
solution of DtO@K
= 2),
(c) phenylacetic acid in &O.
Figure 6-7 -------
p
-
_
_
_
U
(173)
Transmission spectra of Phe in sqlutions of D20 with pH ranging fÏom 7 to 11. The pH
values are indicated at the corresponding spectm
Figure 6-8 -----
--_.__.
------
(174)
SNIETIR spectra over the range of fiequencies, 1300 to 1800 cm-' q u i t e d using the spolarized i n f k e d beam, for the Au(ll1) electrode in 0.1 M KCl04 + 1.O6x 10'2 M Phe
solution in D20. For each spectrum, the reference potential El was equal to -0.60V/SCE,
and the value of the E2 is indicated in the figure.
Figu re 6-9 -------------------------------------
-------- --------- ----- (178)
SNIFTIR spectra for Phe adsorbed at the Au(ll1) electrode fiom 0.1 M KC104 -f- 1.06 x
1oe2M Phe solution in DzO using p-polarized hfrared beam. For each spectrum, the
reference potential El was equal to -0.60 V/SCE, and the value of the E2 is indicated in
the figure.
Figure 6-10 -----------------------------
-------------------------------------- (181)
M Phe solutions, (a) integrated
For a Au(1ll) electrode in 0.1 M KC1O4 + 1.06 x 1 0 - ~
intensities of 1725 cm-' band(square, C=O stretch of the deuterated carboxylic group),
1410 cm-' band(circle, syrnrnetric 400-stretch) and 2343 cm-'(triangle, CO2 band). (b)
Gibbs excess vs electrode potential plot.
Figure 4-11-------------------------------------------------------------------
(183)
Dependence o f the fiequency of the minimum on the negative lobe of the bipolar bands
(closed points) at 1410 cm-' and (open points) at 1725 cm-' on the electrode potential.
Figure 6-12 ---------------------------------------------------------------------(184)
SNIFTIR spectra in the CO2 asymmetric stretch region resdting fkom the oxidation of
Phe adsorbed at the Au(ll1) electrode in 0.1 M KCI04 + 1.06 x 10" M Phe solution in
-
HzO, acquired using s-polarized idiared beam. For eachipectrum, the reference potential
El was equal to -0.60 V/SCE, and the value of the Ez is indicated in the figure.
Figure 6-13 ---
-
----
-----
-
(lm
Models d e m i h g orientation of Phe at the Au(ll1) electrode nirface; (a) at wgative
potentials, (b) at positive potentials.
Figure 7-1 ---------------
----
-----------------------
095)
Size distribution of DMPC vesicles measured by dynamic light scattering amlysis.
Figure 7-2---------------------------------0
Film pressures of DMPC monolayer at the gas-solution (0.05 M NaF) interface measured
by Langmuir-Blodgett microbalance,
Figure 7-3 ---
(199)
__U_--I_----
AFM images (lp
by lpm) of a g l a s slide before (a) and after (b) coatulg with 1 mg
mL-' DMPC vesicle sohtion for 1 hour at 30 OC.
Figure 7-4 ----------------------------------------------
-------------- P)
Extending force as a h c t i o n of separation distance between the AFM tip and a glass
slide coated by a DMPC bilayer using LI3 deposition.
Figure 7-5 -------------------------------
---------------------- ----------- P)
Extending force as a fûnction of separation distance between the AFM tip and a glas
slide coated by a DMPC bilayer using a direct vesicIe fusion technique. The
concentration of DMPC vesicles is 1 mg rnL1 DMPC in 0.15 M KCl solution.
Extending forces as a fünction of separation distance between the AFM tip and a glass
slide in an AFM liquid cell. The ce11 was filled up with 1 mg mL-' Egg-PC vesicles in
0.15 M KC1 solutim for 30 minutes, foilowed by gently flus~'zingthe cell with 90 mL of
0.1 5 M KCl solution.
Figure 7-7 ----------------------------------------------------------------------
@OS>
Cyclic voltammograrns of DMPC bilayers at Au(l1l)-solution interface.
Figure 7-8 -------------------------------------------------------------------
@7)
Forward differential capacity curves of DMPC bilayer at the Au(l1l)-solution interface.
Figure 7-9 -
-
-___I--
----- W)
Charge density vs electrode potential plots of DMPC bilayers at a Au(ll1)-solution
interface.
Figure 7-10
-----------
--------
@Il)
__---__._-__--C__
Cyclic voltammograms of Au(ll1) in 0.05 M NaF, 0.1 mM Femcyanide and 0.1 mM
Femcyanide with a DMPC bilayer formed by a vesicle fusion technique. Scan rate was
20 m v s".
fusion technique. Potentials are shown in graph.
Figure 7-12 ----------
----I__----UI_
-o
Correlation of the charge densities for DMPC bilayer on Au(l11) to the IR band intensity
at 2852 cm-'.
Figure 7-13
--
-------- ----------------
@17)
Dependence of CH2 stretching modes on the elecîrode potentials for DMPC büayers at
the Au(1ll)-solution interface. Inset: ratio of IR intensity at 2922 cm-' to the intensity at
2852 c d vs the electrode potential.
monolayers at the Au(l 11)-solution interface.
Figure 7-15---- ------------------ --------------
------- -------------------------P O )
Comparison of charge densities of DMPC bilayers with that of DMPC monolayers at
Au(1 l 1)-solution interface.
fiision; Triangle: rnonoiayers by horizontal touch.
Figure 7-17
---- --------- --------------------------------- -1
Comparison of the film pressures for DMPC monolayers at the gas-solution interface
with that at the metal-so lution (inset) interface.
Figure 8-1
_
C
_
_
_
_
I
-
---
-
--en
Chernical structures of nystatin and a "barrel stave" mode1 of the channel formed with
nystatins. The protuberance on the bottom represents the amino sugar and the shaded
intenor represents the hydrophilic polyhydroxyl portion of the molecule. The exterior
surface of the channel is completely nonpolarFigure 8-2 --------------------------------------
-
-------------
----e38)
Illustration of the structure of the cbannel formed by a dimer (right) of gramicidin A (lefi)
in the bilayerFigure 8-3 ------------------------------------------
-----
@8)
Proposed structures formed by an oligomer (bottom) of ahethicin (top) im a lipid
bilayer. Lefk The conformation in the absence of an applied trans-membrane voltage.
Right: The open channel conformation in the presence of an applied voltage. Middle: An
intermediate conformation.
xiii
List of Tables
Table 2-1 -----
----------_I---------
@a
Comparison of some physical characteristics of bilayer lipid membranes (BLM) with
natrucil membranes
Table 5-1 --------
a391
Vibrational fiequencies and assignments for benmate
Table 6-1 -
- - -
-
-
-
-
-
-
-
-
-
-
-
-
-
1
_
C
-
-
-
-
-
-
-
072)
Vibrationai fiequencies and assignments for transmission and SNIFTIR measurements
List of Symbols and Abbreviations
A
Anion
BLM
Bilayer Lipid Membrane
C
Cation
E
Electrode potential
F
Faradaic constant
AG
Gibbs energy
3
cument density
J
Joule
4
Total charge density
R
Gas constant
T
Thermodynamic temperature
AFM
Atomic force microscopy
CE
Counter electrode
cv
Cyclic vo ltammetry
DAC
~i~ital-to-analog
converter
DC
Dflerential capacity
DMPC
1,2-dimyristyl-sn-gtycero-3-phosphochoIine
DPPC
dipalmitoyphosphatidylcholine
FTIR
Fourier Transform h.f?ared Spectroscopy
G-S
Gas-so lution interface
HT
Horizontal touch
ml?
lmer Helmholtz plane
LEED
Low energy electron deaction
LB
Langmuir43lodgett
OKP
Outer Helmholtz pIane
PZC
Potential of zero charge
Phe
Phenylalanine
S-BLM
Supported Bilayer Lipid Membrane
SCE
Saturated calomel electrode
SNIFTIRS
Subtractively normalized Fourier transform inf?ared spectroscopy
STM
Scannùig tunneling microscopy
sw
SrnaIl Unilameiiar Vesicle
WE
Working ekctrode
r
Gibbs excess
S d c e tension
E3okzman.u constant
Electrosorption valency
Critical length of the akyl tail of lipids
Gibbs energy of adsorptio-n
Charge density on the metal side of the ideally polarized electrode
Surface area
Wavelength
Chernical potential of the appropriate species
Volume of the allcyl tail region of lipids
Film pressure
Variable in Parsons fiinction
The capacity of the uuier layer at a constant charge
The capacity of the ïnner layer at a constant amount adsorbate
Surface strain
Charge number of an ion
Zwitterion
Chapter 1Introduction
The siudy of the adsorption of biologid molecules is one of the most active
contemporary research areas, attracting the attention of scientists fiom dBerent
disciphes. When molecules are assembled on a surface to form a two dimensional phase
(monolayer or biiayer), they are expected to display unusual chemistry due to their
characteristic orientation in the nIm and due to the well-defined positions of lünctional
groups (e.g., carboxyl or amino groups). Very interesting chemistry happens at these
surfàces that would not happen if the same groups were distributed randomly as single
molecules in the bulk of a solution. Some of the properties of fihm of assembled
molecules are most interesting to chemists who want to create surfaces with well-dehed
chernical properties or who simply want to understand the contribution of the dBerent
segments of the molecule to its molecdar activity- These films also attract biologists
because they are abIe to strrdy how very subtle changes to the biological membrane (a
surface) can drastically change the tendency for particular proteins to bind to that
membrane and for enzymatic reactions to occur on that membrane. Some insight into the
Izature and function of biological membranes can be gained by studying synthetic bilayers
consisting of a single type of plzospholipid molecules or simple bilayer mixtures
consisting of phospholipids with proteins or peptides. Examining these d a c e s is also of
interests tc physicists because they can be used to measure intermolecular forces that are
in the nanonewton range. The synthetic membranes are also of intrinsic physical interest
as soft material systerns. They are of parîicuIar interest to electrochemists because
e1ectrochemical techniques aliow the control of the amount and orientation of the
adsorbai mo lecules by applying electrical fieIds. Moreover, the electrochernical studies
of biological molecules could contribute to the development of biosensors and
bioelectronics.
1.1 Objectives
The bio-fbnctionalization of solid surfaces is a central problem in today's
biotechnology. Much progress has been made with bio-functionalization of polymers Cl41. However, substrate-supported lipid membranes provide an alternative building
prïnciple yielding a homogeneous, cohesive coating with the naturd option to insert or
emkd membrane proteins. The goal of this project was to mimic biological membranes
on an electrode surfàce. The work presented in this thesis was of pioneering character and
its objective was to open a new direction of research. It is well understood that the
majority of bio logical membranes are composed primarily of phospholipids and proteins;
the amino acids k i n g the basic units of the latter. Consequently, one rnust know the
interfacial behaviours of phospholipids and h
o acids before they can be used to
control the properties of artificial membranes at electrode swfâces. In this thesis, the 1,2-
dimyristyl-sn-glycero-3-phosphocholine@MEC) was used to constnict the monolayer
and supported bilayers. DMPC is a satumted glycerophospholipid. It is rehtively stable in
air and the Liquid crystdgel phase transition temperature (Tm) of DMPC is
approximately 24 OC.Therefore, it is easy to produce DMPC films that are either in the
Iiquid crystalline or in the gel state. Phenylalanine (Phe) was selected for examination
because it plays important roles in the hydrophobie interactions arnong the proteins and
phospholipids. One of the objectives of this project was to compare the properties of the
DMPC bilayer deposited ~ ~ i reither
i g the conventional Langmuir-BIodgett (LB)method
or by a direct vesicle fusion technique. The main advantage of the vesicle fusion is the
ease of inserting protek, especially for the proteins which cannot tolerate the transfer
process during LB deposition.
The other objective of this thesis is to amplement and extend the systematic
studies of molecular adsorption on gold electrodes carried out in our Iaboratory. The
previous students in our laboratory have studied the interfàcial behaviours of inorganic
ions [S-71,small organic molecules [7-91 and surfactants [10,11]. My objective was to
extend t h research by providing new information concerning adsorption of organic ions
and zwitterionic moIecules. The benzoate ion and Phe have been selected as a mode1
organic anion and a zwitterionic molecule in my project. Both electrochemical and
spectroscopie techniques have been employed to investigate the surface concentration
and geometry of the adsorbed species. A well-defined single-crystal Au(l11) plane was
used to study the coordination of the organic ions and the zwitterionic molecules. This
gold electrode was chosen because it behaves as an ideal polarizable electrode in a very
broad range of potentials.
1.2 Fundamental Concepts
1.2.1 Amlno Acids
An amino acid is, by d e m i o n , an organic compound containing an amine group,
-mand a carboxylic acid group, -COOH in the same mo lecule. While there are many
forms of amino acids, au of the important amino acids found in living organisms are aamino acids wkch have both the -COOH and -N& groups attached to the same carbon
atom, defined as the a-carbonatom. The simplest amino acid is glycine, H2NCH2COOH.
It contains no asymmetric carbon atoms (tetrahedral carbon atoms with four different
groups aîtached). AU other amino acids contain an asymmetric carbon atom and are,
therefore, optically active. The optical isomers of each asymmetric carbon atom are
tradiitionally distinguished by the letter D or L- Only those amino acids which are of the L
form (Ieft-handed) at the a-carbon are found in proteins.
The general structure of the a-amino acids is R-CHN&-COOH. At low pH vaiues,
both functional groups are f U y protonated so that the amino acid molecule assumes the
cationic forrn, RCH(NI&>COOH. At high pH values, alI the acidic protons have been
removed and it assumes the anionic fonn, RCH(NH2)COO-. At pH near the isoelectric
point
@O,
the amino acid exists as the zwitterion, RCH(NH~~COO-,because the
carboxylic acid group is a much stronger acid than the ammonium group.
1.2.2 Lipids
A lipid is defïned as a water-insoluble biomolecule which has a high solubility in
nonpolar organic solvents such as chloroform The simplest lipids are the fats, which are
triesters made up of one glycerol unit and three fàtty acid units. The term fats is also used
as a general synonym for lipids, so the more precise term triacylglycerols or triglycerides
is preferable for the simplest lipids. TriacyIglycerols are used primarily for energy
storage in animals. Other categories of lipids, the phospholipids, glycolipids, and
choiesteroi, are the major constituents of biological ce11membranes.
Both phospholipids and glycolipids are amphipathic molecules, that is, molecules in
which one end is hydrophobie and one end is hydrophilic. The fatty acid chahs form the
hydrophobic end of the molecule while the polar glycerol-phosphate-alcohol or
sphingosine-sugar portion forrns the hydrophilic end of the mo lecule.
When the phospholipid molecules are spread at the gas-solution (G-S) intediace, the
hydrophilic head is toward the solution side and the hydrophobie tail extends to the gas
phase* In order to obtain a homogeneous munolayer at G-S interface, the temperature
rnust be carefùlly controiled above the phase transition tempecature(Tm) of phospholipids
where the phospholipids are in liquid crystalhe state. The liquid crystalline state is a
distinct phase of phospholipids Observed between the crystalline (gel) and isotropic
(liquid) states. The liquid crystal phase is characterized by molecules that have no
positional order but tend to point to the same direction. The phospholipids organize
themselves into layers with translational and rotational motion of the molecules.
Therefore, lateral diffusion within the phospholipid film is expected when the
temperature is above the Tm.
1.2.3 Bilayer Membranes
When amphipathic mo lecules are placed in an aqueous solution, their hydrophobic
tails attempt to orient themselves toward each other. They may achieve this by either
forming s d l spherical micelles or a plana. lipid biiayer. A micelle is a smaU structure,
generally less than two micrometers in diameter, while a lipid bilayer typically bas a
thickness of about 0.5 micrometers and can have an area of several square dlimeters.
Sonication or other treatrnent of lipids can produce spherical liposomes, or lipid vesicIes,
which contain aqueous solution within an ençlosing lipid bilayer. Liposomes are
considerably Iarger than micelles.
Both liposomes and planar bilayer membranes can be prepared f?om simple
solutions of phospholipids such as the diacylglycerol-3-phosphatidylcholines. Studies of
such artificially prepared membranes have provided much insight into the function and
operation of membranes in living cells. Simple bilayer membranes are highly permeable
to water molecules, while ions such as sodium and potassium can only traverse them
more slowly (by nine orders of magnitude) 1123. For small molecules, the rates of
perrneation through simple bilayer membranes increase with their solubility in nonpolar
solvents relative to their solubility in water. These Observations strongly suggest that
Figure 1-1 Singer-Nicolson fluid mosaic mode1 for biomembrane structure. (Adapted
IÎom S.J. Singer and G.L. Nicolson, Science, 175 (1972) 720-73 1 .)
transfer of substances through a membrane requires desolvation, transfer through the
anhydrous membrane interior and then resolvation on the other side of the membrane.
Natural biological membranes (Figure 1-1) consist rnainly of different lipids(such
as phospholipids, glycolipids, cholesterol) and proteins and thek structures can be
regarded as a two dimensional solution of ordered lipids and proteins that have a high
degree of motion. The protein hction varies fiom 20
- 80% depending on the type of
membrane [12]. The dBerences in transport behaviour between bio logicai membranes
and simple bilayer membranes are due to the incorporation of protein molecules or other
specificdy transporthg molecules directly into or onto the surface of a bilayer ce11
membrane. Considerable research in biology and in chemistry is k i n g directed toward
the transport properties of membranes.
1.2.4 Biomernbrane Electrochemistry
Electrochemical reactions play signifxcant roles in biological activities. The
electrochemical potential is the driving force for many biochemical reactions including
energy storage and substance metabolism. The fidamental membrane processes of
living cells, for example, generation of ion gradients, sensory transductance, conduction
of impulses and energy transduction, are electrical in nature. Each process involves
charge movement in a specialized protein structure, where part of the protein forms a
channel for conduction of ions. The opening of the channel is controlled by changes in
physical factors such as the electrical potential across the membrane or the binding of
signalling (e-g. neurotransmitter or hormone) molecules and ions to specific receptor or
enzyme sites. The electric field could also induce the redistribution of the biomembrane
-
coaiponents. Poo and Robinson Cl31 observed a field-dependent redistribution of
fluorescently labelled concanavalin A in the membranes of living muscle ceh. This
e x p e k e n t indicates that the electric field may change the conformation of molecules
and their properties in the membrane. Electroporation is another important fielddependent membrane phenornenon. Figure 1-2 shows a hypothetical mechanism relevant
to electroporation (induced by a short pulse, 1o4 c t,h
< 10-~s).Accordhg to the cartoon
in Fig 1-2(A), a fiee volume fluctuation is believed to be invoived in the transport of
nonpolar molecules across membranes [14] and a possible early precursor to a
hydrophilic pore. Fig 1-2(B) shows that an aqueous protrusion is envisaged as a more
direct precursor to a hydrophiüc pore [15]. Fig 1-2(C)depicts the hydrophobie pore, a
high-energy transient structure that is believed to be a direct precursor to a hydrophilic
pore [16]. Fig 1-20) shows the hydrophilic pore, believed to be the primary participant
in short-tem electrical behaviour and probably involved in rnolecuiar transport [16]. Fig
1-2(E) describes a composite pore with one or more proteins at the inner pore edge, a
speculative possibility that rnight account for a metastable pore that persists after the
potential has decayed because of reversible electrical breakdown. Finally, Fig 1 - 2 0
illustrates the composite pore due to a 'Yoot-in-the-door" mechanism, which involves
insertion of a linear charged macromolecule into a hydrophilic pore, such that screened
coulombic repulsion prevents shrllikage of the pore. This is another candidate for a
metastable pore, which can persist and assist in the transport of small ions and molecules
long after the potential has decayed to a srnall value in a reversible electrical breakdown.
The present view of electroporation assumes that transitions fiom A-+B+C+D
have a
nonlinear increased fiequency of occurrence as the potential is increased. Similar
transitions may lead to E and F.
nnnn
n
111111
n"huliU
U
Figore 1-2 Hypothetical structures of transient and metastable membrane conformations
that are believed to be relevant to electroporatioa (Adapted fcom: T.Y. Tsong and RD.
Astumian, Bioelectrochem, Bioelectroenerg. 15 (1986) 457-476.)
1.3 Scope ofthe Thesis
The research in this thesis has a multi-disciplinary character covering Biophysics,
Electrochemistry and Surface Science. A brief introduction is given in Chapter 1 where
some fidamental concepts are presented. Chapter 2 describes the theoretical background
and gives a review of the literature necessary for the interpretatim of the experimental
results in the thesis. There are four different topics covered by Chapter 2: 1) the electrical
double Iayer and the thermodynamics of the metal-solution interface, 2) a brief
introduction to in-situ FTIR spectroscopy of electrochemicd intefices, including a
review of the surface selection d e s and the data acquisition procedures, 3) a review of
the electrochemical studies on amino acids and 4) a detailed description of the theories
and experimental techniques conceming supported phospholipid bilayers and their
possible applications.
The experimentai procedures are presented in Chapter 3. These include the
description of the electrochemical cell and the procedures for cyciic vo ltammetry (CV),
dserential capacity @C), chronocouIometry (step experiment), the film pressure
measurements using the Langmuir microbalance, the operation of subtractively
norrnalized interfàcial Fourier transform i . e d spectroscopy (SNIFTDRS), the
preparation of small damellar vesicles (SUV) and the construction of phospholipid
monolayers and bilayers.
The experimental results will be described in the next four chapters. Chapter 4 and
Chapter 5 discuss the electrochemical and
FTIR spectroscopie studies respectively of
benmate adsorption on the Au(11Z) electrode surface. Phe fias k e n examined using the
same techniques and a neutral solution where the Phe exists as a zwittenonic foim in buk
solution These results will be descriid in Chapter 6. Chapter 7 compares the properties
of a DMPC bilayer formed by LB deposition with that formed by the direct vesicle fùsion
technique. P r e h i m r y I
R studies of the DMPC biiayer on gold will also be inciuded.
Chapter 7 describes a new methodology exploration The research described in this
chapter opens a new research direction to be followed in our laboratory,
Chapter 8 summarizes the conclusions of this thesis. New research directions will
be proposed based on the experimental remlts obtained in this thesis.
References
1. W. Hanke, CRC Crit. Rev. Biochem 19 (1985) 1.
2. A. Ottova-Leitmannova and H.T. Tien, Prog. Surf. Sci 41(4) (1992) 337.
3. V. Tvarozek, KT. Tien, 1. Novotny, T. Hianik, J. Dglugopolsky, W. Ziegler, A.L.
Ottova-Leitmannova, J. Jakabovic, Rehacek and M. Uhlar, Sensors and Actuators
B, 19 (1994) 597.
4. E. Sackmano, Science, 27 1 (1996) 43.
5. Zhichao Shi, PhD. Thesis, University of Guelph, 1996.
6. Shijie Wu, PhD. Thesis, University of Guelph, 1996.
7. Aicheng Chen, PàD. Thesis, University of Guelph, 1998.
8. Lorne Stolberg, PhD. Thesis, University of Guelph, 1990.
9. Dongfang Yang, Ph.D. Thesis, University of Guelph, 1995.
10. Dan Bizzotto, PhD. Thesis, University of Guelph, 1996.
11. Vlad Zamlynny, M.Sc. Thesis, University of Guelph, 1998.
12. L. Stryer, Biochemistry, New York, San Fransisco, W. H. Freeman and Company,
1995.
13. M.-M. Poo, and K R Robinson, Nature, 265 (1977) 602.
14. RO. Potts, and M.L. Francoeur, Proc. Natl. Acad. Sci U.S.A. 87 (1990) 3871.
15. Y.A. Chizmadzhev, V.B. Arakelyan, and V.F. Pastushenko, Bioelectrochem
Bioenerg. 6 (1979) 63-70.
16. I.G. Abidor, V.B. Arakelyan, L.V. Chemomordik, U . k Chizmadzhev, V.F.
Pastwhenko, and M.R Tarasevich, Bioelectrochem Bioenerg. 6 (1979) 37.
Chapter 2 Review of Theories and Techniques
2.1 Electrical double layer and thermodynarnics of the electrified
interface
2 1 Introduction
At the solid-solution interface, the properties of the electrolyte solution are altered
due to interactions between the solid and the electrolyte. The surface of a so1id electrode
whose potential is controlled by a potentiostat is charged. The presence of these
electrostatic forces causes long range changes in the electrolyte composition at the
interface. The charge on the metal is screened by the charge of the ionic cloud on the
solution side of the interface which leads to charge sepration in the interfacial region and
formation ofthe so called electrical double layer. Many models [l-31 have been proposed
to explain the behaviour of the double layer. Below we introduce several models that are
rnost fiequentiy used to explain the phenornena taking place at the electrified interface.
2.1.2
The electrical double layer
The first mode1 of the 'electrical double layer' was proposed in the 1850's by
Helmholtz [4]. Ke assumed that there is no electron tramfer across the interface and that
the solution is a solution of a strong electrolyte. The interactions between ions in the
solution =d the ekctrode were assumed to be electrostatic in nature, and resulted fiom
m
an excess or
the fact that the electrode holds a charge density which arises ~ o either
deficiency of electrons at the electrode surface. In order for the interface to remah
neutral, the charge held on the electrode is balanced by the redistribution of ions close to
the electrode surface.
The ions attracted to the charged surface are assumed to approach the electrode and
to form a layer balancing the electrode charge. The distance of closest approach is limited
to the radius of a solvated ion. The plane in which the ions are located at the distance of
the closest approach is termed the outer Helmholtz Plane (OHP). The overall result is two
layers of charge (the double layer) and the potential &op across the interface is confined
only to this region. This model is analogous to an electrical capacitor which has two
plates of charge separated by a layer of dielectnc (solvent). The potential drop occurs in a
linear m e r between the two plates. In fact, when irnpedance analysis is performed on
electrochemical systems the response due to the electrolyte redistribution at the interface
is modelled in terms of capacitive elements.
The model of Helmholtz does not account for many factors, such as screening out
of the ionic charges due to thermal motion of ions, the possibility of a specîfïc adsorption
on to the surface and the interaction between solvent dipole moments and the electrode.
A more advanced model was proposed by Stem [SI to address some of these limitations.
Stem combined the compact layer mode1 of Helmholz with the dinuse layer model of
Gouy-Chapman [6,7j.In Stem's mode&,the ions are allowed to move in solution so that
the electrostatic interactions are in competition with Brownian motion. The result is that
ions are located not in a plane but in a region close to the electrode surface containing an
excess of one type of ion called the dinuse layer. The potential drop occurs over this
extended region. Grahame 18'91 developed Stem's model m e r by introducing the idea
of two planes of closest approach, the inner Helmholtz plane @I
atl')
a distance xi for
specifically adsorbed ions and the outer Helmholtz plane (OHP) at a distance x2 for nonspecifically adsorbed ions. Thus the OHP forms a boundary between the inner and diffuse
Meral
Plane\
Inner
Helmholtz
f plane/
Outer
HeLmholu
Plane
Figure 2-1 Mode1 o f the electrical double layer at the metal-solution interface.
layer. B o c k et al- [IO] took the solvent molecules into account. They proposed that the
solvent molecules are oriented according to the charge density on the metal and form a
layer of oriented dipo les together with the specifically adsorbed ions (see Figure 2- 1).
Many modifications and improvements [Il-141 have been made to these early models
and a detailed review of these models is given in reE [15].
2.1.3
Therrnodynamics of the electrified interface
(9 Derivation of the electrocapillary eauation:
The thermodynamics of the electrified metal-solution interface has been thoroughly
reviewed before [1649]. The thermodynamïc theory of the interfacial phenomena was
derived fiom the GÏbbs adsorption isotherm
-dy = c r$pi+ ç mI
I
pj
where y is the interfacial tension (or surface tension), T's are the Gibbs excesses (or
surfàce excess) of the uncharged species i and ionic species j, pi and
correspondhg chemical and electrochemical potentials. The application of the Gibbs
adsorption isotherm to the idealIy polarizabIe electrode (initially Hg) l a d s to the
eIectrocapillary equation
-dy = sdT - Vdp + o,dE, + T&pi
where s is the entropy, T for the temperature, V for the volume, p for pressure, o~ for the
charge density on the metal side of the interface, & for the electrode potentid with
respect to an intemal reference electrode wfiich is reversible either to a cation or an anion
in the electrolyte, risis the relative Gibbs excess of species i with respect to a reference
component (water for aquaeous solution) and pi is now the chernical potential of either a
neutral species or a charged species i in the solution. The extension of Gibbs isotherm to
the electrocapillary equation is based on two assumptions. i) The real electrode system
has a well-deked thermodynarnic state (it is in thermal, mechanical, electrostatic, and
physicochemical equilibrium). ii) The system is electrically neutral when considered as a
whole, but there is a charge separation across the interface. At constant temperature and
pressure, equation 2.2 can be reduced to equation 2.3.
Note that it is practically dinicult to employ a reference electrode that is reversible
with respect to the cation or anion of the electrolyte investigated. When an excess of
electrolyte is used, however, the activity of the reference cation or anion can be easily
kept constant. Consequentiy, the potential of the Ïntemal reference electrode versus an
extenial reference electrode will be constant. & = &+ constant and dE+ = Gt
where
kf
is the
electrode potential measufed versus the extemal reference electrode. In this
case, the & in equation 2.3 c m be replaced by
bfor simply E. In aqueous solution, the
risin equation 2.3 is defined as
where ri and rwt,represent the Gibbs excess of species i and water respectively, xi and
x
,,
are the mole fiactions of species i (10'~at most) and water (55.5) in the bulk
solution. TWt, is estimated to be 1014molecules
assurning a close-packed monolayer
on the rnetal surface [20].Consequently, the second term on the nght side of equation 2.4
is often four to five orders of magnitude less than Ti. The difference between the relative
and absolute surfêce excesses can be neglected. To facilitate fuaher discussions, we
would rewrite equation 2.3 in a simple form:
The charge density at the metal side c m be obtained f?om the variation of y with
respect to potential at constant T, p and solution composition. This is cailed the Lippman
equation:
Further differentiation of the Lippman equation l a d s to the relationship between the
interfacial tension y and the differential capacity C as shown in equation 2.7.
The above therxnodynamic description o f the metd-solution interface is applicable
to a lïquid electrode scch as Hg. When the electrocapillary equation is applied to a soïd
electrode such as gold, an additional term called surface stress, Y must be introduced to
the electrocapillary equation Thus the equation 2.2 should be rewritten as equation 2.8
for a solid electrode.
Here Y
=y
+ ûy/a~,,where
ce is the elastic surface strain However, it was recently
shown [2 11 that the additional term in equation 2.8 has a very srna11 contribution to the
change of surface tension To a good approximation, the equation 2.2 can also be used for
solid electrodes as well. However, for a solid electrode it is not possible to measure y
directly. Ln this case one usualiy uses capacity measurement to calculate
and y by the
integration method:
It was shown that this method worked weli for mercury electrodes. It couid be also
extended to the measurements at solid electrodes [22j. Lipkowski et al. [23,24] have
developed a chronocoulometry technique to measure the equilibrium charge density at
solid metal electrodes. The surface tension can be obtained by the integration using
equation 2.10.
{ii) Determination of the film pressure at constant wtentid and at constant charge:
The film pressure at constant potential is d e h e d as
where y w and y* are the interfacial tensions of the electrode in the absence and in the
presence of organic molecu1es For a Liquid electrode, these values c m be detennined fiom
independent electrocapillary measurements. For a solid electrode, the absolute value of y can
not be detemiined. In this case, however, the equation 2.10 can be applied The Ef in
equation 2.10 is chosen as the potentid at which no adsorption takes place. The integration
constant y ~ isf independent of the presence or absence of investigated species in solution at
this point. Thus the term Y E wdl
~
be elimuiated when equation 2.10 is combined with
equation 2.12. nierefore, the fih pressure of adsorbed molecules on a solid electrode can
be mathematically calculated as follows:
where the subscripts 0 and 8 4 indicate the presence and absence of the organic molecules
in the bulk electrolyte.
The film pressure a .constant charge is d e h e d by the following equation:
lj=y+Et~~
The total dEerential of the Parsons fùnction is equal to:
dg=dy+EdaM+cMdE
(2-16)
I f we replace dy in equation 2.16 by the electrocapillary equation (equation 2S), a new
equation results:
This equation is similar to the electrocapillary equation, but now the charge density rather
than potential is the independent electncai variable. Equation 2.17 indiaes that ~(CTM)can
be calculated by the integration of E with respect to the charge at a constant chernical
potential of al1 species in the solution,
where
00
is the charge density at Et where no adsorption takes place.
5(00)
is therefore
independent of the presence and the absence of organic molecules. The film pressure at
constant charge cm be then calculated by:
(iui Determination of Gibbs excess:
The expression for the relative Gibbs excesses can be derived nom the electrocapillaiy
equation (Eq. 2.5) using potential as the independent electrical variable, and fiom Eq. 2.17
using charge density as the independent electrical variable. If it is known that the h i t y
coefficients do w t change with the concentration of the analyzed species, then
dp = RTdlnc
(2.20)
This condition is satisfied for ionic species if an excess of inert electrolyte is used and for
neutral moiecules if their concentration is very iow in bulk solution. The Gibbs excesses at
constant potential and at constant charge are described by equation 2.21 and 2.22
respectiveIy.
where c is the bulk concentration of the organic species. The reiative Gibbs excesses can
aIso be plotted as a function of the bulk concentration of the organic species at a constant
potential or at a constant charge density. In this way, the adsorption isotherm is determined.
Cross dserentiation of equation 2.5 and equation 2.17 gives equation 2.23 and 2.24
respectively.
The correlations of equation 2.23 and equation 2.24 may be used to examine the
consistency of the thermodynamic data Actually equation 2.23 can be used to determine
the electrosorption valency ( Z ) as defined by equation 2.25.
The expressions in equation 2.24 give the Esin-Markov coefficients.
(iv) Determination of the Gibbs energies of adsorption:
The Gibbs energy of adsorption can be calculated fiom a fit of the experimental
M a c e excess data to the equation of a particular adsorption isothenn The adsorption
#
isotherms that are fiequently used are (i) Henry, (hi Langmuir and (iïi) Frumkm ûotherm. In
the Limit of zero coverage d l isotherms reduce to the Henry isutherm. Under this condition
the film pressure depends Luiearly on the bulk concentration of the organic species. When
the standard state is unit mole hction of the organic species in the b u k o f the solution and
monolayer coverage ( 0 4 ) of the "ideal noninteracting" adsorbate, the equation of state
correspondhg to the Henry isotherm is given by equation 2.26 [26].
n = RT&
(2.2 6)
/ 55.5
The adsorption equilibrium constant P in Eqs 2.26 on be deterrnined corn the initial dope
of n versus cI55.5 plot The zero coverage Gibbs energy of adsorption c m then be
detemiined fiom the value of fl using the expression of
AG'^ = -RTLn@
The advantage of
this method is that no mode1 of adsorption needs to be assumecl SirnSarly, AG'-
can be
detemiined f?om the initiai dope of the film pressure at constant charge vs. concentration
plots according to the fo110wing equation:
here charge is used a s the independent electrical variable.
The energy of adsorption may also be descriid in tenns of the f i 1 1 coverage Gibbs
energy of adsorption AG'^=^). The AG'^=[ values c m be detemiined &y plotting the film
pressure as a function of the logarÏthm of the mole fiaction of absorbate in the bulk solution
and extrapolating the linear segment of this plot to zero film pressure. T h e intercept of the
ln(d55.5) axis multiplied by RT gives the &II coverage Gibbs energy o f adsorption [27,28].
The difference between the magnitudes of the zero coverage and f Ü U coverage Gibbs
energies of adsorption illustrates how the properties of the d a c e phase deviate f?om the
properties of a perfect solution (a perfect two-dimensional
of solvent and the
adsorbate mo lecules).
The Gibbs energies of adsorption can also be detemiined using an empirical
method. In that case, the sUTface pressure
(a)data are
fitted to the "square root "
isotherm [29,30].
i n ( k ~ c +) h ~ = h c BO'^
~
(2.28)
where c is the bulk concentration of adsorbate,
P
=
exp(-AGkT) is the adsorption
equilibrium constant, B is a constant and 0 is the surface pressure. Wheo analysis is
carried out a . a constant potential, <D is set as y~ - ye, where y is surface tension. If we
take charge as the independent electrical variable, cD is set as
- 50, where 5 is Parson's
function n i e plots of the u > l R vs. ln(kTc/O) are almost linear. The Gibbs energies can be
calculated by extrapolating their linear segments to zero d a c e pressure in accordance
with equation 2.29.
Fi(kTc/Q?)],=
-In P
The standard state here is an "ideal"
(2.29)
r = 1 ion
for the adsorbed species and an
"ideal" c = 1 M for the bulk species.
2.1.4
Mode1 of the double layer
In order to interpret the structure and the charge distribution at the interface, a
specifc physical model has to be used. Here the thermodynamic data have been
processed by the Grahame-Parsons model [3 11 of the double layer. The inner-layer
capacity C' c m be calculated fiom the relationship
where cdiS the capacity of the difhise Iayer, C is the overall capacity determined by the
differentiation of the charge density curves and d r h Mis the siope of the
r versus CM
plots. The value of F(ar/hM)is very close to one at most cases. So the second term on
the right side of the Eq. 2.30 could be neglected. The inner-layer capacity is now
approximately equal to the overall capacity. The inner-layer capacity is a fünction of two
variables, the charge on the metal and the arnount of adsorbed anions. Therefore it can be
expressed in terms of the two components, the inner-layer capacity at a constant charge
rC and the inner-layer capacity at a constant e u n t adsorbed C
. as follows [3 11:
The rC is determineci fiom the dope of the plots of the potential drop across the inner
layer ( A + =
~E
~~
û~and
- 42) v e r w the charge (-FT)of adsorbed adsorbate at constant
92 is the outer Helmholtz plane potential. The potential $2 can be obtained by
equation 2.32 in aqueous sdution
a~+ a
i = - 11.74(c4IR sinh(19.46~4~)
where z = z+ = lz-1, crd is in pC
cbin mol dm-3anci t$2 in volts.
(2.32)
Ci=
zFT, is the charge
corresponding to the specifically adsorbed anions. Once rC has been determined, ,C can
be easily calculated by equation 2.3 1. Equations 2.33 and 2.34 show that rC and ,C may
be considered as integrai capacities respectively,
K=&/(x2-xf)
,C=E/X~
(2.34)
where E is the pemiittivity, x, and xz are the distance of the b e r and the outer Helmholtz
planes fiom the metal surface respectively. The ratio of the ber-layer capacities at a
constant arnount of adsorbed anion and a constant charge is equal to the electrosorption
valency [3 1]:
The electrosorption vaiency may fûrther be used to calculate the surface dipole that
represents a dipole formed by an adsorbed anion and its image charge in the gold
electrode. Tne magnitude of the surface dipole is a direct measurement of the polarity of
the bond formed between the anion and the electrode surface. The surface dipole is
described by equation 2-36 132-341,
where eo is the charge of an electron (1.6 x 10-l9C) and E is the permittivity of the inner
layer which is regarded as the same 8s the perrnitîrvity of the vacuum (8.85 x 1612
T'
m-').
2.2 In situ FTIR in electrochemistry at the metal-solution interface
2.2.1
Introduction
Before the appearance of in situ vibrational spectroscopy, WNisible spectroscopy
was undergohg rapid development as an in situ tool for the investigation of electrode
processes 135,361. The fiactional change in reflectivity can readily be detected as low as
104 [37]. However, WNisible spectroscopy has many inherent drawbacks due to its
inability to i d e n t e adsorbed molecules that do not absorb the light in this range and the
difllculty in detennining the electronic structure of the electrode without a pre-defined
model. The red breakthrough in the development of in situ vibrational spectroscopy
came in 1973 when the first Raman spectra were obtained by Fleischmann [38,39].
M a r e d spectroscopy was first introduced hto etectrochemistry by Hansen [40], Kuwana
and Osteryoung [41]. However, the early studies were based on the intemal reflectance
techniques. Bewick et al. developed the external reflection Wared method [42-441 and
Pons et ai. developed the Fourier transforrn version of this technique [45,46]. The
modulation of the photon polarization in electrochemicai FTIR was introduced by Habib
and Bockris in 1985 1471. This technique allows one to detennine the character of the
surface CO-ordination of the adsorbate. Recently, Faguy et al. developed a high
throughput reflection accessory for infrared spectroelectrochemical studies, which
dramatically improves the sensitivity and reproducibility of in situ measurements (48-5 11.
Korzeniewski et al- [SS] designed a jacketed cell to control the temperature in the cell
used for in situ FTIR studies. Stable spectra could be obtained up to 70 OC. A brief
introduction of the in situ IR spectroelectrochemistry presented below is based on
selected publications [48-681.
2.2.2
Surface selection rriles
Infkred reflectance spectroscopy is based on the specular reflection of the incident
light on the substrate sudàce. On metal surfaces the component of the electric field
parallel to the surface is almost pedectly screened [66], whereas the component
perpendicular to the surface is enhanced by a factor of nearly 2 for angles of incidence
near to grazing. Figure 2-2 shows the dependence of the electric field strength E/E, on the
angle of the incidence 0 at a platinum surface for v
= 500
and v
= 2000
cm-'. Here, E. is
the amplitude of the oscillating electric vector in the incident beam and E is the resultant
amplitude at the surface; p and s refer to the components of the electric vector parallel
and perpendicular to the plane of incidence. The resulting s component of the electric
field Es is close to zero for any angle of the incidence as shown in Figure 2-2. The p
component can be M e r split into two components Epl (perpendicular to the metal
surfàce) and Epii(paralle1 to the surface). This figure shows that Epllis also very small,
whereas Ep1 increases with increasing 0, reaching a maximum between 80 and 90" but
falling rapidly to zero at 90". These results show that only the p component of inçared
radiation has appreciable field strength at the electrode surface, it alone carries
information on vibrations of surface species. It is interesthg to note that the s
components are even more effectively screened out at 500 cm? than that at 2000 cm-'.
The response of the metal surface that screens the parallel field wiil dso screen out any
dynamic dipole moment appearing on the molecde in the direction parallel to the d a c e .
This may be visualized in terms of an induced image dipole in the metaL Figure 2-3
shows two instantaneous dipdes, which are perpendicular and parallel to a metal surface.
For the dipole-oriented nomal to the surface, the image dipole is reinforced whereas, for
the parallel dipole, there is a very effective caricellation. Hence, optimal absorption is
obtained when the component of the dipole derivative with respect to the normal
coordiite of an adsorbed species is in the sarne direction as the incident p-polarized
radiation. Oscillators with their dipo le derivative components parauel to the surface will
have a smaller probability of S a r e d absorption. Therefore, the relative intensities of
30
60
90
Angie of incidence, @/des
Ang ie of incidence, €3 / deg
Figure 2-2 The d a c e electric field E/Eo for platinum at (a) v
= 500
cm? and (b) v =
2000 cm-' as a fünction of the angle of incidence 9. (adapted fiom ref
î.n.f?ared absorption bands can give information regarding mo lecular orientations on the
electrode surfàce. This distinction between absorbances due to oscillators oriented
parallel and perpendicdar with respect to the surface is known as the surface selection
rule. Another consequeme of these feaîures is that the maximum absorption of the
i n f k e d bearn will occur for p-poiarized light. This would justw the use of a polarizer in
reflectance spectroscopy. Distinction between adsorbed and dissolved species may be
facilitated by modulation of the polarkition and measuring a diffirence s p e c t m
between reflectances at both s- and p-polarizations. Randomly oriented solution species,
which can absorb both s- and p-polarized light equally well, can be disthguished fiom
adsorbed species, which can absorb only p-polarized radiation.
Figure 2-3 lnstantaneous dipoles appearing on an adsorbed diatornic molecule oriented
perpendicular and parallel to a metal surface. (adapted fkom ref [55])
Recently, Osawa et al. have found that ïnfiared absorption of molecules is
remarkably enhanced when they are adsorbed on or near very thin films of silver, gold,
copper, and indium. This phenornenon is called surface-enhanced i n . e d absorption
(SEIRA) [64,67,69].
They have shown that only the vibrational modes which give dipole
changes perpendicular to the metal surbce are infrared active. The enhanced infiareci
absorption arises fiom an electnc field at the metal nnfacc produced by the incident
infrared radiation through the excitation of a localized plasmon of the metai. Since the
enhancement is the largest at the metal surface and decays sharply within a distance of a
few mowlayers fiom the surface, the solid-liquid interface can be investigated
selectively. They have also demonstrated that monolayers at the interface can be detected
at a very high sensitivity without interference nom the buk solution by this technique.
2.23
Sampling Method
The sampling methodç in FTIR c m be classified into two groups, extemâl reflection
170-741 and attenuated total interna1 reflection (ATR) [75,76]. ATR sampling has not
k e n used widely in electrochernistry because preparation and use of the thin metal film
working electrode has k e n dificult. In the extemal reflection arrangement, the infiared
beam is directed through a polarizer and onto the fiont surface of a highly polished disk-
shaped working electrode. A special thin-layer electrochemical ce11 is employed that
permits the inhred beam to enter and strike the Eont surface of the poIished disk at an
incident angle of about 60' to 70' with respect to the d a c e normal. The beam is
reflected out of the ceil and is colIected by optics that image the radiation onto a detector.
The optical configuration is different for static linear polarization (SLP) and polarization
modulation experiments. In the SLP expriment, the electrode potential is constant while
the po1arizer is set to ailow the radiation to be either s- or p-polarized. The polarization
modulation layout includes a photoelastic modulator (PEM), which rapidly switches the
polarkation state of the radiation while the electrode potential is controlled to a specific
value. The optical configuration and experimental procedures used in experiments with a
static photon polarization are described below.
2.2.4
Ce11 design
Different designs of the thin-layer cell for extemal reflection idtared spectroscopy
have been reported in [44,49,50,59,64,65,68]. Figure 2 4 shows a drawing of a typicaI
t hree-electrode cell. The working electrode is a disk-shaped metal approximately 6- 10
mm in diameter- The electrode material should have high reflectivities in the hfkared
spectral region. Pt, Au and Ag have k e n employed most fkequently. The fiont face of the
#
disk is plished to a &or
finish; typically with a slurry of alumina or diamond paste.
The disk is mounted in a plunger. The counter electrode is typicaily platinum wire or
gauze that is looped around the working electrode and positioned just behind the metal
disk. The reference electrode is mounted in an extemal cornpartment that connects to the
main charnber through a Luggin cappillary. An infrared-transparent window is attached
to the fiont of the cell. The working electrode is positioned in order that the polished
fiont face is adjacent to the optically Bat surface of the window and during S a r e d
spectroscopy experirnents, it is pushed against the IR window to trap a thin Iayer of
solution (typically less than 10 p)between the window and the electrode. Incident IR
radiation passing through the window and the thin-Iayer of solution onto the electrode
d a c e is reflected back out of the thin-layer ce11 and is subsequently focused ont0 the
detector. The solvent absorption is minimized using the thin-layer ceU. The choice of
window materials depends on the spectral region of interest and on the solvent. Examples
IR Radiation
Working
Eiectrode HoIder
Figure 2-4 The design of the spectroelectrochemical cell for reflection infiared
spectroscopy. (Reproduced fiom A. Bewick et al. J. Electroanal. Chem 160 (1984) 47.)
of IR window matenals for the m i d - ï d k e d region include C e , ZnSe, Si and KRSS (a
mumire of 42% TlBr and 58% TII). For the f x - h h r e d region polyethylene is usually
employed. It is obvious that the eIectrode surface must be planar and parallel to the
IR
window in order to obtain a thin and unifonn layer- Recently, Faguy designed a novel
spectroelectrochemical cell using the l
R window, a ZnSe hemisphere, to collimate the
beam 148-5 11. Such a design provides near-crit ical angle of reflect ion, Iower e s t - d a c e
reflection losses, and reestablishes the instrument focal point. As a result, enhancements
of the signal-to-noise ratio of better than five fold are achieved 1501.
Korzeniewski [52] designed an insulated infiared spectroelectrochemica1ce11 that is
capable of operation at variable temperatures. The cell is built around a g l a s chamber
that holds electrolyte solution and the workïng and counter electrodes. The fiont of the
chamber consists of an O-ring seal joint onto which the cell window mounts. The
reference electrode is held in a separate cornpartment behind a wetted stopcock and
comects to the glas chamber through a capillary tube. The glass chamber is insulated
and secured inside an aluminum jacket. The cell temperature sensing and heating
functions are built into the working electrode. A thennocouple is used for temperature
measurement. The cell performs well for the study of CO adsorption.
2.2.5
Spectral acquisition
Acquisition of in situ S a r e d spectra has been achieved by several methods.
Potential difference techniques reduce background solvent absorption and have k e n
most widely applied. The sknplest approach involves the use of a ETR spectrometer with
a static Iinear polarization. A sequence of interferograms is collected while the electrode
is held at a constant potential (El). The interferograms are added, average4 Fourier
transfbrmed to a single beam spectnim and electronicdly store& The electrode is then
stepped to a new potential CE2), the data collection and processing is repeated and the
resulting single beam spectrum is stored in a new fiIe. The single beam spectra encode
the wavelength-dependent reflectivity of the system, R A few number of interferograms
are obtained at two different potentials El and E2, and the process is repeated untii the
desired S N is obtained. A potential difference spectrum is computed fiom the ratio of
single beam spectra recorded at two different electrode potentials. The two single-beam
spectra are subtracted and divided by the intensity at one potential to obtain the final
normalized spectrum. The result is a dserence spectnun with magnitude AR/R
= CR(E2) - R(E1)I 1NE11
=
where R(EI) and
(2.3 7)
R(Er) are the electrode reflectivities at potentials EI and
E2,
respectively. The rneasured spectrum shows only changes in reflectance caused by
changes in potential. The instrumental drift and al1 the IR absorption by environmental
species are cancelled. The method described above is called subtractively norrnalized
interfacial Fourier transform ïnhred spectroscopy (SNIFTIRS). Figure 2-5 descn'bes the
principles of SNPTIRS. The s-polarized photon wilI have a change of phase of 180"f i e r
reflection The electric fieid of the s-polarized photon at the electrode surface is close to
zero. Thus there is no interaction between the s-po1anZed photons and the molecules at
the electrode surface. For the p-polarized photon, however, the electric field at the
eiectrode surface will be enhanced and paraIIel to the plane of incidence. There wiI1 be
an interaction between the p-polarized photon and the dipole perpendicular to the
electrode surface. The IR adsorption at the surface only occurs under this condition.
There will be no adsorption when the dipole is parallel to the electrode surface even if a
p-polacized beam is applied. Figure 2-6 summarizes the characteristic shapes of the
SNIFTIR spectra For s-polarized beam, there will always be a positive peak if the
molecules are IR active and are adsorbed at the electrude surface. For p-polarized beam,
the shape of the spectra can be either positive or negative or bipolar depending on the
intensity of the interactions between the molecules and the electrode, the orientation of
the rno Iecules and the alignment of the dipo les at the electrode surface.
The SNIFTIRS technique uses potential modulation and static linear beam
po1ar:zation. However, in situ spectra can also be recorded with the photon polarization
modulation techniques [44,58,70,77]. Here the dBerentia1 reflectance (ARIR) spectra are
computed as the ratio of (IcIp)/(Is+Ip),where 1, and 1, are the intensities of reflected sand p-polarized radiation respectively. The sum ((I,tI,))
and dserence ((L-1,)) spectra
are determineci by Fourier transformation of the interferogram signal before and after
demodulation, respectively. With polarization modulation, it is possible to record an in
situ h.f?iired spectnun. of an electrode at a single potential. However, because s- and p-
pohrkd radiation do not sample quivalent regions of the thin solution layer between
the working electrode and window, the dserential (demodulated) spectral signal contains
a strong background nom the bulk solvent and electrolyte. In addition, the photoelastic
rnodulator (PEM, which switches the polarization state of radiation at 74k Hz) has a
wavelength dependence that adds a slowly varying sinusoida1 component to the
differential reflectance spectra [7 2,781. Therefore, polarization modulation spectra
recorded at two different electrode potentials are typically ratioed to eliminate the bulk
features and correct the spectral baseline.
A change in reffectance can result fiom a change in potential for a number of
reasons. These include changes in the coverage of species adsorbed on the electrode,
changes in the nature of bonding of adsorbed species, migration of ions into or out of the
optical path and reactions that generate species in the t h layer or on the electrode
surface- Al1 these changes may result in a change in band intensity and shape or the
appearance of new bands. A band position may change due to changes in the fiequency
of the oscillator as the strength of bonding to the surface is changed with electrode
potential. Consequently, a bipolar shape in a SNIFTIR spectrum is observed. In addition,
the reffectivity of the electrode material itself rnay change as a resuit of the potential-
induced reconstruction/construction C79-8 11. The adsorbate can also induce nanoscale
restructuring of the electrode d a c e in UHV and gas-phase environments [82-841, as
well as in solution 185,861. The electrode restnicturing could occur even without altering
the substrate unit ceK
2.2.6
Applications
h f k e d spectroscopy is applied most ofien in electrochemistry to study reactions
and adsorption at conventional polycrystallhe working electrodes. Its use in single
crysîal efectrochemistry has becorne possible with the development of bench-top methods
for c1eaning and ordering single crystal surfaces. Atomically well-defined electrodes have
enabled the study of surface geornetric and electmnic uIfiuences on processes such as
electrocatalytic reactions, electrodeposition, assembly of organic [87] and phospholipid
188,891 monolayers and structure dynamics of phospholipid bilayers [90-931. Generdy
speaking, infirared spectroscopie techniques have been applied to elucidate the following
issues in electrochernistry:
Monitoring of the potential-dependent coverage of molecular/ionic species on the
electrode [60,94]
Determination of the surface CO-ordinationand geometry of adsorbed molecules/ions
[48-53,943
Identification of the
IR active intermediates and products produced f?om
electrocatalytic reactions [61]
Probing of the interactions arnong the molecules/ions and with the electrode (i.e.
adsorbate-induced nanoscale restructuring) [62]
In this thesis research, the SNIFTIRS technique has k e n used to explore the
potential-induced adsorption/desorption and the orientation changes of benzoate,
phenylalanine and DMPC at Au(ll1) electrodes.
2.3 Interfacial behaviour of proteins and amino acids at the soiidsolution interfaces
The interest in the interfkcial bebaviour of proteins at solid surfaces originates fkom
the need to better uoderstand the mechanisms of processes associated with their use in
advanced technical applications and industrial problems. On one han&
the use of
immobilized enzymes in analyti d techniques, in biotechno logy and in chromtography,
requires such knowledge [95]- Adsorption of proteins and enzymes on various kinds of
substrates is widely used for the punfication, identification, fixation, and separation of
these materiais [96]. On the other hand, it is desirable to minimize protein adsorption in
food industries to avoid fouling problems [97,98]. Many papers have been published to
investigate the fouling process [99-1011. In general, the driving forces for protein
adsorption at the solid-liquid interface c m be divided into four categories [102]:
(i)
Protein-surface interactions. These interactions have various origins such as
electrostatic attraction or repulsion between the protein and the interface andor
Van der Waals attraction The most important contribution to the long range
electrostatic interaction is the Coulomb interaction due to the net charge of the
surface and the proteins. The net charge on the protein results from dissociation of
various amino acids located mostly on the periphery of the molecule. Any nonhomogeneity in the charge distribution, the so-called "mosaic
charge
distribution," leads to a dipole moment and can also contribute to the electrostatic
interaction which is further complicated by the incorporation of ions into the layer
between the interface and the protein [Z 03- 1051.
ci)
Dehydration of interfoces on the outside of the protein and on the solid surfiace.
Dehydration o f hydrophobic interfices promotes protein adsorption while the
dehydration of hydrophilic interfaces opposes it. As protein mo lecules have large
radü (2 to 4 nm) compared to that of a water molecule (0.14 nm), adsorption
results in the release of a large number of water molecules. Therefore, this
interaction is very important and often ovemdes the other driving forces.
(Ui)
Shtcturd changes in protein molecules upon adsorption.
The densely folded structure of globular proteins in solution is rnainly attributed
to the fact that hydrophobic interactions in the interior of the molecule (i-e.
prevention of hydration of these residues) are stronger than the intrarnolecular
electrostatic repulsion in the periphery of the molecule and the loss of entropy
J
upon folding. When the protein is brought into contact with a hydrophobic
interface, the balance between these forces may change because the hydrophobic
interior of the mo lecule can prevent hydration in an alternative way by positioning
itself agauist the interface. This unfolding will be promoted by electrostatic
repulçion between the charges on the adsorbed proteh Unfolding of the adsorbed
protein wiI1 increase the surface area occupied per protein molecule. It will also
reduce the electrostatic repulsion between like charges in the adsorbed protein,
because these become more distant, and increase the entropy of the protein, The
extent of unfolding of the protein also depends on the characteristics of the
sorbent [106]. A hydrophobic surface promotes unfolding. Also, charges on the
interface can compensate for those on the adsorbed protein, thereby changing the
intramolecular electrostatic repulsion and altering the degree of unfolding.
(iv)
Surface coverage-dependent lateral interactions due to acczrrnulation of protein
molecules at the inteface.
The electrostatic repulsion between e q - y charged molecules and the resulting
loss of entropy of the system oppose completion of monolayer coverage, Dipoledipole interaction between the proteins on the surfke can be either repulsive or
attractive depending on the degree of aRïgnment of the molecules.
The relative importance of the various drEving forces is dif5erent for various classes
of proteins. Therefore, it is very hard to distïnguish the comrnon driving forces behind
their adsorption and predict the behaviour of a particular protein at any solid-liquid
interface. Roscoe and CO-workers have used electrochemical and spectroscopie
techniques to determine the mo lecular conformations and orientations of adsorbed protein
molecules. Their studies have k e n made on P-lactoglobulin [lO7,lO8], a-hctalbumin
[109], K-casein [ I l O], ribonuclease [108], lysozyme [108], BSA [log, 1 111, insulin [1121,
cytochrome c [113], myoglobin [Il31 and hemoglobin [113] at a platinum electrode. The
surface charge density resulting fiom protein adsorption was foimd to be sensitive to the
conformational behaviour of the proteins. The carboxylate groups were deterrnined to
play a major role as the surface-active functional group of the proteins at anodic
potentials. The adsorption process occurring on the metal electrode surface after addition
of proteins could be described by equation 2.38 [IOq
P + nM
+ P(M),,& + ne
(2.3 8)
where P represents the protein adsorption on the metd (M) surface with n representing
the number of sites for carboxylate interaction with the metal d a c e , accompanied by
the transfer of a total of n electrons.
Amino acids are the basic units which build up the proteins. Studying the interfacial
behaviour of amino acids will contribute significantly to the understanding of biological
processes such as conformation changes and metahlism pathways of proteins.
Unfortunately, the eady electrochemical studies did not focus on the biological
implications. Their original intention was to examine the electrochemical properties of
molecules carrying dBerent fùnctional groups Cl 14-1151. It was Gouy [Il61 who first
noticed that the amino acids behaved differently fiom other organic molecules. He found
that the interfacial tension of the mercury-0.5 M Na2S04 interface was lowered by
glycine more at extreme potentials than near the maximum of the electrocapillary curve.
In the early 1970s, Parson's group [114] studied the adsorption of glycine at the mercury-
water interfàce in acidic, neutrd and basic solutions. In neutral solution, ther, was a
minimum in the adsorption, as wodd be expected for a highly polar adsorbate
(zwitterion). However, the high dipole moment of glycine was not apparent in the f o m of
a potential shift at constant electrode charge. Hence, it was suggested that the molecule
was adsorbed with the dipole essentially parallel to the electrode s d a c e . Glycinate (in
basic solution) and glycinium (in acidic solution) ions behaved as typical anions and
cations, respectively. Horanyi and
CO workers
[117,1183 used a radiotracer method to
study the adsorption of arnino acids on platinun electrodes. It was found that in acidic
solution there was no significant ciifference between the adsorption behaviour of glycine
and acetic acid. This means that, in this case, the amino group has little effect. At low
concentrations, only loosely adsorbed species c m be observed. However, strong
chemisorption occurred in alkaline medium Alanine [Il51 ciiffers somewhat fiom
glycine. Both loosely physisorbed and strongly chemisorbed species were present in
acidic sofution while o d y strong chemisorption occurred in aïkaline solution. The
radiotracer has also been used to examine the adsorption of phenylalanine and tyrosine on
a bismuth electrode [119,120]. 330th amino acids fàiled to desorb at anodic potentiak in
acidic solutions. This is probably due to the interaction of the additional n-electron of the
adsorbed arornatic compounds with the electrode surface. The limiting surface
concentration for both amino acids is about 2 x 10''~and 3 x 10'" mol
CM < O
at c r >~ O and
respectively. There is a third fbnctional group, -S-CH3 in methionine. An attempt
to anchor the zwitterionic methionine to the mercury electrode [12 11 did not lead to the
insertion of a large density of these dipolar ions into the layer immediately adjacent to the
metal surface. It is likely that the arnino acid group, whether in its ionic or zwitterionic
form, is located on the solution side of the outer Helmholtz plane where the charges are,
to a large extent, screened by supprting electrolyte in the diEuse layer. Therefore, it can
be assumed that the interaction of the S-atom with a metal surface plays a predominant
role in the adsorption process. The adsorption of methionine on platinum displays sirnilar
phenornena [122]. Roscoe [123,124] has carried out mechanistic studies on the
adsorptionfoxidation of amho acids at the platinum electrode. According to her
experirnents, the followllig rnechanisms have been proposed for the adsorption of amino
acids on a metal electrode (M). In basic solution, where the amino acid is anionic:
In acidic solution, where the amino acid is cationic:
Although the electrocatalytic properties of gold are not a s good as those of platinum, the
oxidation of phenylalanine at the Au(l11) surface has been obsenred in the present work.
The production of CO2 indicates that the oxidation process of phenylalanine follows the
above mechanism. The detailed description of phenylalanine adsorption is presented in
Chapter 6.
2.4 Interfacial electrochemistry of phospholipid bilayers
2.4.1 Introduction
Owhg to the complexity of biological membranes, it has long been desirable to
investigate artificially constituted membrane systems. The study of these simplified
mode 1 systems provides insights for understanding sirnilar processes in real bio logical
membranes. The reconstituted planar bilayer lipid membrane @LM) that separates two
aqueous solutions was £irst reporîed in the 1960s [125]. BLM is also an acronym of Black
Lipid Membrane. The ultra-thin Lipid films of bimo1ecuIa.r thickness appear "black"
because the light reflected fiom the fkont interfàce undergoes a phase shifi of half a
wavelength and interferes destmctively with the light reflected fiom the back interface,
which experiences essentially no phase shifk 11261. During the last thirty years, the
conventional BLM has k e n used for models of biological membranes. In particular, the
BLM has been used to elucidate the rnolecular mechanisms of biomembrane fùnction
such as ion sensing, material transport, excitability, gated channels, antigen-antibody
binding, signal transduction and energy conversion [l27-l3 O]. In order to improve the
mechanical stability and to produce the large area membrane, a self-assembled BLM on a
soWgel substra.te was introduced in 1976 as an attempt to develop a -del
system for
use as a solar energy conversion device [13 1-1351. The supported BLM (s-BLM) was
then formed on metallic wires [136], conducting SnOz glass [137], gel substrate 11381401 and on microchips [141]. The s-BLM has a long-term stability and a desirable
rnechanic property. It opens a window for laboratory studies o f membrane properties and
for practical applications in biosensors and rnolecular recognition devices.
2.4.2 Construction of the lipid bilayer on a solid substrate
1. Langmuir-Blodgetf technique
The supported BLM can be formed by using a modified Langmuir-Blodgett @B)
technique Cl421 pioneered by Takagi, Azurna and Kishinoto in 1965 [143] and later
rnodified by others [144-1473. We modified the method fkther in our experirnent. Our
setup is shown in Figure 2-7. In this apparatus, a glass cylinder across the G-S interface
rnaintains a section of the G-S interface ftee kom the phospholipid monolayer, thus
producing a Iipid-fkee wuidow. A few dmps of BLM-fonning solution (n-hexane or
chloroform as lipid solvent) are introduced onto the surface of the aqueous solution
outside of this glass cylinder. M e r evaporation of the solvent, the aqueous surface is
coated with a lipid monolayer. A suitable substrate (gold or g l a s slide) is raised and
lowered through the monolayer of the lipid spread at the G-S interface. Interactions
between the substrate and the lipid molecules in the mono layer lead to the tramfer of the
monolayer fiom the G-S interface to the surface of the solid substrate. If a movable
barrier is employed to ccimpress the monolayer to maintain a constant surface pressure
during tramfer, the substrate may be raised quite rapidly. If no barrier is ernployed, the
procedure must be performed slowly (-1 mm min-') so that a complete mowlayer at the
G-S interface may be maintained by addition of lipid to the interface fiom excess lipid in
the form of aggregates [148]. In the first case, the transfer ratio could be used as a
parameter representing the effectiveness of transfer. It is the ratio of the area of barrier
movement required to compensate for the Ioss of molecules fkom the monolayer at
constant film pressure to the geometric area of the solid substrate. This ratio is unity for a
complete transfer,
a
The formation of a bitayer using the LI3 technique depends on how the substrate is
moved through the G-S interface. In order to deposit the fist leafiet of the bilayer with
the polar head of the lipid oriented towards the hydrophilic substrate, the substrate must
be raised through the G-S interface (step (i) in Figure 2-7). In order to deposit the second
leaflet of the bilayer, the mono Iayer-coated substrate must be lowered through the G-S
interface (step (ii in Figure 2-7). This second step leaves the bilayer-coated substrate in
the aqueous phase. RemoWig the s-BLM from the solution would necessitate a third pass
through the monolayer at the G-S interface, where the integrïty of the bilayer could be
compromised. Removing the sample through the Lipid-fiee window (step (üi) in Figure 27) overcomes this problem. Deposition by the LB technique allows a large degree of
CyLinder
window
Figure 2-7 Apparatus designeci to deposit a bilayer of phospholipid on a planar substrate.
control over the experimental conditions for the spreading and the orierntation of
phospholipids in each leaflet. It may produce a heterogeneous bilayer. T h e major
disadvantage of this technique is that it complicates the transfer of proteins to ahe bilayer
since many membrane proteins do not tolerate transfer fiom a monolayer at the G-S
interface [149f.
The general advantages of LB techno logy are [1501:
Formation of continuous coatings, in contrast to vacuum-deposited and spincoated films.
Very precisely controlled thickness and very good unifonnity.
Coating o f curved surfaces with equal uniformity.
On suitable substrates, the films c m bridge pores up to the micrometer range.
The general disadvantages are [150] :
The coating involves the dipping of al1 objects hto an aqueous subphaseIn many cases the substrate bas to be modified to enhance its adhesion properties.
The dipping speed is quite slow.
It is very dficult to h d a smooth substrate in order to form a perfect monolayer
without defects.
2. Horizontal touch
The horizontal touch (KT), or Schaefer method [15 1,1521 is a variation mf the L-B
technique. Here, the horizontaily oriented substrate is 10wered or raised vertkally . The
solid substrate is brought into horizontal contact with the compressed monolayer at the
G-S interface (Figure 2-8). Tt is then lifted fiom the surface and the monolayer 5s
1
substrate
1
1 substrate 1
1
substrate
1
1 substrate 1
Figure 2-8 Demonstration of the Horizontal Touch technique. (a) The substrate is Iowered to
contact the monolayer at the gas-solution interface, (b) The substrate is lified horizontally, (c)
The lipid is transferred to the substrate with non-polar tail attached to the substrate surface, (d)
Double horizontal touch (not a biomembrane model), (e) Lipids t m e d over (possible?) to fom
the bilayer on the substrate during the double horizontal touch.
transferred in the sanie orientation as on the water s d c e , although a reorientation of the
monolayer may occur [153]. Bizzotto et aL have successfully demonstrated that films of
insoluble surfactants [154,155] and 4-peotadecylpfidhe [156] can be transferred to an
electrified metal-solution interface using this technique. A tramfer ratio of unity is
achieved at potentials around the potential of zero charge. The HT rnethod is very simple
and can be the choice for autornated systems. It fonns a good monolayer on a solid
substrate and is very usefil when the substrate needs to be chemically modified by a
layer of amphiphiles. However, challenges arise if it is used to construct a lipid bilayer.
The molecules in the first monolayer have to f i p over in order to form a leaflet of the
LBM (polar head towards the substrate). Extra operations on the lipid molecules are
required. Some of the Iipids may be desorbed from the substrate, thus producing more
defects in the film.
It is now known that the electric field could be the driving force to turn the
phospholipid on the electrode d a c e . Nelson and coworkers [157-1591 studied the
potential-induced reorientation of pho spholipid mono layers on a hanging mercury drop
electrode. Based on the analysis of the capacibnce-potentid curves and cornputer
simulations, they proposed a flipping mechanism for the adsorbed phospholipid
mowlayer before the formation of bilayers. Chen and A b m a (1601 employed rapid scan
cyclic voltarnmetry to investigate structural transitions of a phosphoiipid monolayer on a
mercury surface. Their observations suggested a mechanism of O pposite flipping (Figure
2-9). Initially the system is in a high (metastable) energy state when the polar head of the
phospholipid is in contact with the hydrophobic mercury surface. As the potential is
scanned negatively, the etectrode surface is more negatively charged. The electrostatic
repulsion between the electrode and hydrated polar head increases so that the molecules
are repelled fiom the electrode surface. An active cluster forms near the electrode
surface. However, since this intermediate structure alters the energetic state within the
rnonolayer, it could initiate a cooperative process with neighboring molecules until ail of
them were reoriented into the tails-dom structure. If the potential becomes more
negative, a phase transition of lipids may occur and the formation of micelles is assumed
[16 1,1621. However, some of the membrane proteins may not tolerate the harsh
1- Initial Çtate
11, Activated Clusters
I
W. Final State
III, Flipping and Cooperativeness
Figu re 2-9 Proposed mechanism for the structural reorientation of a phospho lipid
monolayer on a mercury surface Tom heads-down to tails-down fashion. (Adapted fiom
reference [1601.)
conditions that are needed to turn over the lipids. The vesicle spreading technique is one
possibility to solve the problem.
3. Vesiclef i o n
Phospholipid bilayers can be deposited by both indirect and direct vesicle
spreading. In the former, the biomimetic bilayer membrane is formed by spontaneously
coating a covalently tethered hydrophobic layer [l63-165]. Here the thiol group of
alkanethiol mo~eculesattaches to the metal surface by a covalent bonding to produce a
hydrophobic surface. The covalent association with the surface is insensitive to changes
in buffer, pH, ionic strength, or lipid composition. The phospholipid vesicles in aqueous
media are then exposed to the alkanethiol-coated surface. The vesicles will rupture and
the lipids will spontaneously assemble into a second layer over the akanethiol
mono layer. The driving force for the self-assembly of phospholipid layer is the reduction
of the fiee energy of the allcanethioVwater interface. The hybrid bilayer preserves the
ciyoamic nature of the abnethiof rnonolayer and has good mechanicd stability- The
plananty and stability of these bilayers facilitate the use of atomic force microscopy for
surhce imaging. However, the real biomembrane should be mobile while the fluidity of
the hybrid bilayer is poor and the insertion of tram-membrane proteins is dificuh.
A supported LBM may also be constmcted by direct fusion of unilamellar
phospholipid vesicles to an appropriate hydrophilic substrate [166]. This approach ailows
for easy incorporation of proteins into the bilayer of the s-LBM via incorporation of
proteins into the bilayer of the vesicles [167].In order for a s-LBM to form by vesicle
fusion, four processes must occur: vesicle approach to the substrate; vesicle adsorption;
vesicle rupture; and fusion of adjacent patches of bilayer on the substrate. Approach and
adsorption of the vesicles will occur spontaneously if the substraîe is suitably
hydrophilic. Rupture of the vesicles may also occur spontaneously for highly strained
vesicles that have a diarneter at the lower linnit for vesicle stability [167]. Altematively,
rupture may be induced by the addition of a monovalent salt (-1 00 mM), or a divalent
salt (-1 rnM) to the solution, or by repeated heating and cooling of the adsorbed vesicles
through the chah melting temperature [167]. Two extremes of the mechanism of lipid
spreading on a substrate have k e n proposed [168]. These two extremes, spreading by
sliding and spreading by rolling, are illustrated in Figure 2- 10. Lipid sliding will occur on
substrates that are very hydrophilic and do not dehydrate the membrane. On these
substrates, the planar bilayer is separated fiom the surface by a 1O - 20
A film of water
[168,169]. In this instance, the s-BLM retains many properties of fiee membranes
including laterai fluidity of both leaflets and fiee diffusion of molecules confined to the
bilayer. The advancing edge of a sliding bilayer is thought to consist of a cylindrical
hemimicellar type structure [168]. Spreading by rolling will dominate on surfaces that
strongly dehydrate the membrane. It does not allow for lateral reorganization of the lipid
on the substrate and results in many defects and patches of overlapping bilayer 11671.
Spreading by sliding will result in a complete, continuous bilayer with "self-healing"
properties.
One disadvantage associateci with vesicle spreading is the cornplexity of the
procedure. The homogeneous vesicles have to be made before the s-BLM is constructed.
In addition, direct interactions may occur between protein incorporated into
proteoliposomes and the hydrophilic surface [170]; the incorporated proteins often
remain active, but they lack laterai rnobility in the s-LBM [17 11. One approach that has
Figure 2-10 Mechanisrns of vesicle spreading on substrates. (a) Rolling. (b) Sliding.
k e n shown to prevent the lateral immobilization of membrane proteins in s-LBM is
tethering the proteins using glycan-phosphatidyl inositol 11721. Aiso, interaction between
membrane proteins and the substrate may be prevented by constnicting the s-LBM via
the fusion of vesicles to a monolayer deposited fiom a Langmuir-Blodgett film [171].
4. Self-assembly [l36,173- 1771
A supported bilayer lipid membrane could also be formed by two consecutive self-
assembling steps. First, the lipid molecules are placed in contact with a hydrophilic metal
surface and monolayer of lipid molecules is irreversibly bound to its surface with the
nonpolar tail towards the solution. Second, the adsorbed lipid monolayer is immersed in
an aqueous solution and interacts with the hydrophobic chains of other lipid mo lecules. A
bilayer wiil f o m spontaneously at the electrode-solution interface. The formation of both
layers is randorn and more defects are expected. Self-assembly is no longer used for
bilayer construction,
2.4.3 Electrochemical studies of the phospholipid bilayer and their applications
Electrical measurements are the methods most used in characterizing the properties
of the BLM. They provide simple means of monitoring changes and interactions between
the BLM and its mudifiers. Both ac and dc methods have k e n used for the measurements
of various electncal parameters, such as the membrane resistance, capacitance, potential,
dielectric breakdown voltage, etc. [178-1841. Because of the very Iow conductivity of the
unmodified BLM, great care is needed in the ùisulation of the membrane chamber,
electrodes, and connections to avoid current leakage. It is generally advisable to enclose
the whole set-up in a Faraday cage made of copper xreen, to avoid electrostatic
interence. Table 2-1 shows sorne physical data of a BLM in cornparison with the
natural membrane. The BLM has proven to be an excellent mode1 for biomembranes.
The electrochemical investigations of iipid adsorption onto the metaI d a c e began
thirty years ago Cl86,187. The dropping Hg electrode was initiaIly used as a substrate for
adsorption of phospholipids. The surface concentration of lipids cl88] on the electrode
was calculated fkom the radioactivity measurement of the radioactively labelec! lipid
molecules. The results showed that the concentration of lipids on the electrode is very
similar to that at the G-S interface around the pzc. Nelson et aL 1157-159, 189-1971 ako
studied the interfacial behavior of lipids at a mercury electrode. They found that the lipids
are oriented with their hydrophobic tail towards the electrode. A minimum of capacity
Table 2-1 Comparison of some p hysical characteristics of bilayer lipid membranes
@LM) with natural membranes Cl851
Property
Natural Membranes
BLM
niickness (A)
60 -90
Electron microscopy
X-ray difEaction
---
Optical rnethds
40 - 80
Capacitance
40 -130
Potential difference (mV)
O - 140
Resistance (ilcm2)
103- log
Breakdown voItage (mV)
100 - 550
Capacitance (pF cm-2)
0.3 - 2.3
Refiactive index
f .37 -1.66
Interfacial tension (ergs cm-')
0.2 - 6-0
Water permeability (lo4 cm sec-')
8 -50
c'Excitability"
Observed
Observed
Ion selctivity and specificity
Observed
Obsewed
Excitation by light
Observed
Observed
was observed of 2 pF cm-2 at the pzc for DOPC [189]. This value is typical for most
lipids Cl591 regardless of the po Iarity of the head group or the saturation and chain length
of the tail. The effects of choIesterol [189], eIectroIyte composition, amho acids and
chelating agents [191,192] on the permeability of the lipid membrane were abo studied.
They assumed that the lipid could tum over with the polar head towards the electrode
surface, followed by the formation of a lipid bilayer between the electrode and the
electrolyte solution at very negative potentials [157- 1591. The electrochemical properties
of the s-BLM are consistent with those of conventional BLMs [180, 18 11.
An electrochemical mode1 [198] has been proposed to describe the properties of
sBLM investigated with the help of impedance analysis. The impedance is deterrnined by
contributions
of
(1)
the
Gouy-Chapman-Stem
layer
(GCS
layer)
at
the
membrane/electrolyte interface, CGCS,(2) the bilayer impedance (including defects), &,
and (3) the electrolyte resistance, &, which is generally ohmic. It is important to note that
ligand binding does not o d y change the membrane impedance but also affects the
impedance of the GCS layer owing to the change in the surface charge of the membrane.
However, the contribution of the GCS Iayer can often be neglected because its specific
capacitance is a factor of 50-100 larger than the membrane capacitance. The effective
membrane capacitance Cm is given by
Cm = OC, + (1 - O ) h s
(2.45)
where Cm, accounts for the contriiution of the perfectly tight areas of the membrane
covering 1008 percent of the surface area and Cd
=
( ~ - B ) C G C ~accounts for the
contribution of defects covering an area fiaction of 1-8.
The fiequency dependence of the complex impedance (Z(o)) is defined as:
Z(o)= Re(Z(o)) + Zrn(Z(o))
= (Z(o)lexp(i+)
(2-46)
Where 4 is the phase angle between the stimulating ac voltage and the response current.
The fkequency dependence of log(IZ(o)l) and of $(a) and the simultaneous least square
fits of both data sets to several network models is shown in Figure 2- 11. Figure 2- 11a
shows a three-parameter network where Ri and Rz are the membrane resistance and
electrolyte resistance respectively and Cr is the membrane capacity. There is a signifcant
difference between the fitted line and the experimental data which accounted for by the
fact that the capacitance of the real membrane is deterrnined by the area fiaction 0 of the
defect-fiee membrane and by a small area fkaction 1-e covered with defects. The former
is denoted by Cr in the model of Fig. 2- 11b and the latter by a series connection of Cz and
RI. We see that the deflection of 4(o) in Fig 2-1 la at high fiequencies is significantly
minimized in Figure 2-1 1b but that the fit for +(a) tends to 4 2 for o + O is in contrast
to the observed data. This means that one has to introduce an additionai resistance
(Figure 2-1 lc) which describes the effect of small ohmic currents in the area covered by
defects because Z(o) obviously becomes ohmic for o + O. These currents indicate that
some electrochemistry occurs at the defect sites.
More recently, Gu et al. [199] developed a novel method for accurate
rneasuements of the membrane capacitance and resistance by the cyclic voltaminetry
technique. Their proposed equivalent circuit for the s-BLM is very similar to Figure 21lc giving good agreement between the simulation and the on-Iine measurement.
From the viewpoint of membrane biophysics and physiology, biological membranes
represent the basic structure of most sensors in nature. They act as the gatekeeper for ion
and nutrient exchange, provide the site for a host of ligand-receptor interactions f 127,
1301 and conduct the signal and the information fiom one side to the other [200-2041. By
definition, the biosensor is "a device that recognises an analyte in an appropriate sarnple
and interprets its concentration as an electrical signal via a suitable combination of a
bio logical recognition system and an electrochemical transducer." B iosensors can be
1
10
f
100
1000
[Hz1
Figure 2-11 Impedance analysis of a supported bilayer deposited by vesicle fusion [198].
Al1 dots are experirnental data. The Lines are the fitting results of their correspondhg RC
models.
ciassified as potentiometric, amperometric, optical or other physicochemica1 types.
Potentiometric devices measure the potential between the testing electrode and the
reference. Amperometric sensors rnonitor currents generated when electrons are
exchanged between a biological system and an electrode. Optical biosensors correlate
changes in concentration, mass or number to direct changes in the characteristics of light.
Other physicochemical sensors detect biologicaI interactions through changes in
enthalpy, ionic conductance and mas. The development of eiectrochemical biosensors
has grown rapidly in the last twenty years [205-2081. However, the problerns associated
with reliability and commercialisation of biosensors are still numerous, such as the longterm stability of enzymes, the biocompatibility for in vivo applications and the selectivity.
Those problems could be overcome by embedding the biological transducers into a sBLM. Entrapping isolates the transducer fiom the harsh condition in the solution, and
greatly improves the stability of the biocomponent such as an enzyme. Further, the
presence of the lipid bilayer reduces the interference and effectively excludes hydrophilic
electroactive species fi0m reaching the detecting surface, such as the metal electrode,
which may cause undesired reactions. A p a t e r selectivity, precision and accuracy of the
assay could be achieved [137,205-2081.
In the potentiometric biosensor, a potential
across the BLM is ïneasured under conditions where no charges are flowing. In the
amperometric measurements, a current is measured when a voltage is impressed across
the BLM [205,209-2121. A variation of the above is conductometric BLM-based devices
in which the conductivity of the BLM is determined. The electrical parameters (such as
resistance and peak current) cari also be detemiuied by the cyclic voltarnmetry technique
[199]. Besides the development of biosensors, the s-BLM has also found other practical
4
applications. These are outside of the scope of this thesis but interested readers rnay
follow the topics bebw.
a
Long lasting mode1 of biomembrane 12131
a
Bipolar redox electrodes [2 141
a
Ion selective membrane sensors [Ml ,2151
a
Electrochemical biosensors for imrnunology [2 16-2181
a
Molecular sensors for bioassay [139,207]
a
Electron transfer [21 1,2191
2.4.4 Phospholipid aggregates and phase transitions 12201
A BLM is heterogeneous fiom the point of view of its interaction with two different
contacting phases. The interfacial properties of membranes can be understood to a large
extent in terms of the laws of chernistry and physics of the interface. The major force that
governs the self-assembly of amphiphiles into well-defined structures (Figure 2-12) such
as micelles, bilayers, etc. is the hydrophobic interaction between the hydrocarbon tails
and the hydrophilic nature of the head groups. The hydrophobic interaction induces the
molecules to associate. The head groups impose the opposite requirement that they
remain in contact with water. These two interactions compte creating two opposing
forces [221]: the one tending to decrease and the other tending to increase the interfacial
area a per molecule (the head-group area) exposed to the aqueous phase. The attractive
interaction arises mainly fkom the attractive hydrophobic or interfacial tension forces
acting at the essentially fluid hydrocarbon-water interface [222,223]. It rnay be
represented by a positive interfacial fiee energy, characteristic o f interfaces with y
= 50
erg cm-2,though there are indications that this value may be much srndler and closer to
Figure 2-12 DifEerent structures of amphiphile aggregates in solution D20]. Those
stmctures can transform fiom one to the other by changing the solution conditions.
20 erg cm2 12241. Thus the interfacial hydrophobic fkee-energy contribution to poN
(interfacial fiee energy per molecule in an aggregate) may be written as yu.The repulsive
contributions are too complex and diffcult to fornulate explicitly. niese uiclude the
electrostatic repulsion between charged head groups, hydration forces and steric head
group and chah interactions [225]. However, one c m borrow the ideas of the twodimensional Van der Waals equation of state for monolayers [226].
(n + a / d ) ( o- b) = KT
where R is the füm pressure, cr is the area occupied per molecule at the interface, K is the
Bolt~nannconstant, T is the absolute temperature, a and 6 are the two-dimensional
analogues of the van der Waals constants. The first tenn in any energy expansion is
expected to be inversely proportional to the area a.Therefore, poNmay be written, to e s t
order, as
poN= y a + W a
(2.48)
where K is a constant. It is assumed that both these forces act in the same plane. The
minimum fiee energy is, therefore, given when apoN/ûa= O, Le.
pO~(min)
= 2ya0
(2.49)
where a,= dWy, a, will be referred to as the optimal surface area per molecule defined
at the hydrocarbon-water interface. The interfacial energy per molecule may now l
x
expressed in the more convenient form
ON =
2yao + y ( a - ~ , ) ~ / a
(2.50)
in which the unknown constant K bas been eliminated, so that koNas a function of a is
now in terms of the two known or measurable parameters y and a*.We see, therefore,
how the concept of opposing forces leads to the concept of an optimal area per head
group at which the total interaction energy per lipid molecule is a minimum.
The above three equations, while crude, nevertheless contain the essential features
of inter-lipid interactions in micelles, bilayers and membranes. They imply that the
interaction energy between lipids has a minimum at a certain head-group area a,,about
which the energy varies parabolically, i.e. ehsticaLly. Indeed, the area compressibility
modulus k may be readily obtained fkom equation 2.49 since, by defrnition,
Elastic energy =
kT(a*',
IL
which immediately gives a value of k = 2y = 100 erg cm2 per monolayer, while for
bilayers k = 47 c 200 erg cm2, in agreement wÏth measured values on lipid bilayers and
-
biological ce11 membranes where k is usually in the range 100 200 erg cm-2.
HaWig established the equations that adequately describe the interaction between
lipids within aggregates, we have yet to establish into which type of structure different
lipids will assemble. Geometric considerations must now be appiied to determine the
most favoured lipid structure, i.e. the shapes and s b e s of rnacromolecules formed by
these aggregates. The geornetric or packing properties or lipids depend on their area a,,
the volume v of their hydrocarbon chain, which wvill be assumed to be fluid and
interfaciaL
(hydrophobie)
attraction
Figure 2-13 Graphic representation for the physical parameters of an amphiphile 12201.
incompressible, and the maximum length that the chains can assume. We shall cal1 this
the maxMum or critical ch& length I, (see Figure 2-13). The critical length is a semiempirical parameter, since it represents a somewhat vague cut-off distance beyond which
hydrocarbon chains can no longer be considered as fluid. However, as may be expected,
it is of the same order of magnititiïde, though somewhat less than, the fülly extended
length of the c h a h [221,2271. According to Tanford [2211, for a phospholipid composed
of n carbon atoms,
v
= (27.4 +- 26.9n) A3
2, < lm, = (1-5 + 1.265n) A
(2.52)
(2.53)
It has been shown previously 12271 that, for iipids of optimal area a, hydrocarbon
volume v and criticaI chah Iength Zc, the value of the dimensionless packing parameter
da& will determine whether they will fonn spherical micelles (da& < 1/3), nonspherical micelles (1/3 < da& < %), or bilayers (1/2< da& cl).
Splierical
For lipids to assemble into spherical miceiles their optimal
surface area a, must be sufficiently large and their hydrocarbon volume v nifficiently
s
d such that the radius of the miceHe will not exceed the critical chain length Z,. From
simple geometry the micellar radius R will be given by
R=3v/uo~Zc
i-e. da& < 1/3, and the aggregation number will be
M = 4 7 r ~ ~v/ 3= 47r
a.
(2.55)
Cyiindrical micelles: Most Lipids that form spherical micelles have charged head
groups, since this leads to large head-group area G,Addition of sa1t partially screens the
electrostatic head-group repulsion and thereby reduces ao.Those lipids which possess
smaller head-group areas such that 1/3 c vhJCc '/2 c a ~ opack
t
into spherical micelles,
but cm f o m cylindrical (rodlike) structures. Thus single-chained lipids possessing
charged head groups in hÏgh salt or those possessing uncharged (nonionic or Zvcritterionic)
head groups fa11 into this category (e.g. SDS in high sa!t). The aggregation nurnber of
rodlike aggregates is very sensitive to the total lipid concentration C. According to the
equation
<Ni= ~C/CMC
their mean aggregation number
should increase proportionally with
(2-56)
above CMC,
which is indeed found to be the case experirnentally [228].Their sensitivity to the ionic
strength arises fiom its effect on decrûashg a,. As might be expected, the mean
aggregation number should hcrease dramatically wit h increased ionic strength. For
example the aggregation number of SDS micelles in 0.6 M NaCl is -1000 compared to
-
60 in water 12291.
Bihyem: Lipids that form bilayers are those which cannot pack into small micellar
structures due to their small head-group area a, or - as is more cornmon- because their
hydrocarbon chains are too bulky to fÏt into such small aggregates while mahtaining the
surface area at Ïts optimal value. For bilayer-forming lipids the value of da& must lie
between '/z and 1, and this requires that for the same head-group area a,and chain length
I, and their hydrocarbon volume v must be about twice that of micelle-forming lipids (for
which da& is normally in the range 1/3 to %). Therefore, lipids with two chains are
likely to form bilayers, and indeed most of them do.
Vesicies: Under certain conditions it becomes more favourable for closed spherical
bilayers (vesicles) to form, rather than infinite, planar bilayers. This arises since in a
closed bilayer the energeticaily unfavourable edges are eliminated at a nnite, rather than
infinite, aggregation number, which is entropically favoured. Thus, so long as the curved
bilayer lipids can maintain their areas at their optimal value, vesicles should be the
preferred structures. It is a simple matter to show [22TJthat for % < da& CI the srnallest
vesicle that may be formed without forcing the head group area a in the outer monolayer
to exceed a, is given by
Rc = Zj(1- da&)
And their aggregation number is
N = 4 x [ +~(~~ - t ) ~ ] / a ,
(2.58)
t king the bilayer hydrocarbon thickness given by t E- 2 da,.We may note that if da& >
1 (i.e. a, < 42
A2 for double-chained amphiphiles)
such lipids cannot even pack into
bilayers since their head-group area is too small: instead, they form inverted micellar
structures or precipitate out of solution (e.g. cholesterol).
Factors af5ecting changes from one structure to another cm be summarized as:
(1) Lipids with smaller head-group areas @igh du&) should form large vesicles, less
curved bilayers, or inverted micellar phases. For anionic lipids this can be brought
about by increasing the salt concentration, particularly &+,or lowering the pH. This
also has the effect of condensing the chais.
(2) Increased unsaturation, paaicularly of cis double bonds, reduces Ic and thus increases
da&. This leads to larger vesicles and, in general, to more inverted structures.
(3) Increasing the temperature increases the hydrocarbon chah motion, involving trans-
gauche isomerization, and thereby reduces their limiting length 2,. This again leads to
an increased da&.
Figure 2-14 Photographs (magnif~cation10x) of an n-hexadecane droplet (-70pL),
deposited at mercury pooVelectrolyte interface at 4 0 0 mV, take at constant potentids of
(a) -550, @) -1300, (c) -1400, and (d) -1450 mV. (Adapted fiom reference [230].)
Recently, Ivokvic and Zutic [230] studied the potential dependence on the
wettability of hexadecane. The microscope images show that the nonpolar hexadecane
droplet forms a planar-convex lem (Figure 2-14a) on the surface of pooled mercury at 550 m V (Ag/AgCl, potential of zero charge). At -1400 mV (Ag/AgC1) the droplet
changes to a sphere with a very small contact area (Figure 2- Mc). This simply indicates
that the electrode potential could be another factor that effects the phase transition of
Iipids at the metal-solution interface. However the details of the potential-induced phase
transition are still unknown.
References
1. D. Mohilner in AJ. Bard (Ed.) Electroanalytical Chemistry, Vol. 1, Marcel Dekker
Inc, New York, 1966, pp241-
2. J.0 'M. Bockris and A K N . Reddy, Modem Electrochemistry, Vol. 2, Plenum Press,
New York, 1970, pp623.
3. C.M.A. Brett and A.0. Brett, Electrochemistry, Principles, Methods and
Applications, Oxford University Press, Oxford, 1993, pp39.
4. H.L.F. Von Helmholtz, Am. Physik., 89 (1853) 2 11.
5. 0, Stern, 2. Elektrochem, 30 (1924) 508.
6. G.GouyyJ.Phys.,9(1910)654.
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227. J.N. Israelachvili, D.J. MitchelI and B. W. Ninham, Biochim. Biophys. Acta, 470
(1977) 285.
228. N.A. Mazer, G.B. Bemedek and M.C. Carey, J. Phys. Chem 80 (1976) 1075.
229. H. Wemerstrom and B. Lindman, Phys. Rep. 52 (1975) 3.
230. N. Ivosevic, and V. Zutic, Langmuir, 1998, 14(1) (1998) 23 1.
Chapter 3 Experimental Methodology3.1 Electrochemical cet1 and single crystal electrode
A glas three-electrode ce11 [1,2] was used
for the electrochemical
measurements. For studies of DMPC phospholipid, a jacketed glass ce11 was used in
order to control the temperature of the buk solution. The working electrode was a single
crystal Au(ll1) plane which was prepared in our laboratory following the procedures
described by Hamelin [3]. The surface of a new electrode was fürther r e h e d by
electropolishing in 1 M HC104 using an anodic polarization at a constant current of 0.1 A
cm2for about 4 minutes followed by dissolution of the thin layer of gold oxide in 10%
HCl solution [4-61. The electropol i s h g was repeated several times unt il a characteristic
cyclic voltammograrn was obtained. The working electrode for SNTFTIRS was a diskshaped Au(ll1) crystal of diameter
- 1 cm The counter electrode in the electrochemical
cell was a gold c d , and a platinum foil for the SNIFTIRS experiments. The reference
electrode was a saturaîed calomel electrode @CE, Fisher Scientific). The reference
electrode for the electrochemical ce11 was housed externally in a 250 mL beaker which
wntained the saturated KCI (99%, Aldrich Chemïcais Company) solution and was
connected to the electrolyte in the electrochemical cell through a salt bridge. One end of
this bridge fomed a Luggin capillary. It was inserted into the electrochemical ce11 close
to the working electrode. The other end of this bridge was a capillary tube which was
dipped into the saturated KCl solution in the beaker. A three-way stopcock was welded in
the middle of this bridge. The comection was achieved by sucking both KCI and
electrolyte solutions out via one way of the stopcock. This conf~gurationprevented the
4
KCI solution fiom flowing fiom the beaker into the cell, and thereby avoided the
contamination of the electrolyte solution, while ensuring the electrkal contact between
the beaker and the cell. The solution in the ce11 was bubbled for 30-40 minutes with argon
(BOC GAC) to elïminate oxygen Argon was also passed over the top of the solution
through another port during the experïments. An additional port on the side of the cell
was used to add the organic reagents. The ce11 was placed inside a Faraday cage to
elirninate eIectricd and magnetic interference. The g l a s ce11 and its accessories were
cleaned in hot acid (1 : 3, Nitric acid:Sulfuric acid) for 30 minutes and washed
thoroughly with pure water (R > 18 ml2 cm). The working electrode was flame-annealed
and quenched with pure water to ensure the cleanness o f the sufice.
3.2 Chernicals and solutions
Benzoate
Phenylalanine
Figure 3-1 Chemical structures of the adsorbate moIecules.
DMPC
Al1 solutions were prepared with water purified by a M X - Q water system (R > 18
r d 2 cm). The KC104 (ACS Certified fiom Fisher) was purified by calcinating at 300 OC
and recrystalized twice fÏom Milli-Q water. Potassium Benzoate (Aldrich Chernicals Co,,
99%), Sodium Floride (Merck, Suprapur), Phenylacetic acid (Sigma, > 99%), Lphenylalanine (Sigma, > 99%) and 1,2-dimyristyl-sn-glycero-3-phosphocholine(DMPC)
(Avanti Polar Lipids, Birmingham, AL) were used without fiirther purification. The
DMPC phospholipid was dissolved in chloroform (15 mg ml-')and stored in a freezer.
The supporthg electrolyte for benzoate and phenylalanine was 0.1 M KC104 solution.
The suppoahg electrolyte for DMPC studies was 0.05 M NaF.
3.3 Experimentd set-ups
3.3.1 Electrochemical setup
The working electrode (WE) was placed in the ceii and positioned in the hanging
meniscus configuration [7], The height of the meniscus was on the order of 2 mm. The
cell was connected to the potentiostat (PAR, model 173, USA) and the WE potential and
current in the form of analog voltages were acquired by a PC based data acquisition
board (RC Electronics ISC- 16).
The current flowing through the WE when a trianglular voltage sweep was applied
to the electrode resulted in a cyclic voltammûgram. To perform ac voltamrnetry, a
sinusoidal potential generated by a fùnction generator was fed to the external input of the
potentiostat and was superimposed on the voltage rarnp generated by the po tentiostat.
The curent flowing through the WE was phase-analysed with respect to the reference ac
signal incorning fiom the tùnction generator, by the lock-in amplifier (LIA) (PAR model
129A). The output of the Iock-in amplifier consisted of voltage signals proportional to the
real and quadrature components of the WE current. These signals were acquired and
digitized by the RC bard and stored in the computer.
To perform chronocouiometry, the WE potential was pre-set to the base value by
the potentiostat, and then stepped to the variable potential. The cornputer controlled the
duration and the amplitude of the potential steps. The step was generated by the DAC
output of the RC board and then attenuated by the SA and fed to the externa1 input of the
potentiostat. The output signal f?om the potentiostat was proportional to the WE current.
The current was acquired and digitized by the data acquisition board and processed by
home-made software.
3.3.2 Setup for SNIFTIRS
A FTiR spectrometer is mainly composed of two systems; an optical bench and a
computer. The optical bench system includes the source, interferometer, detector and a
set of mirrors. It produces the spectroscopie signal in the form of an interferogram that
contains information about al1 frequencies present in the spectrum The computer reads
the interferogram and uses a Fourier transform to convert the interferogram to the
Eequency domain spectrum. Figure 3-2 shows schematics of the instrumentation for a
SNIFTLRS experirnent. The electrically heated source produces IR radiation. The beam
exits the interferometer and passes through a polarizer. The s- or p-polarized beam is
deflected by Ml and M2 to
reach the electrode surfàce, The beam interacts with the molecules in the thin layer
between the IR window and the electrode where it is absorbed by the sample. The
reflected beam fiom the electrode surface is directed to the detector by M3 and M4. The
mercury cadmium teliuride (MCT) detector produces an electricai signal proportional to
the power of the incident radiation. The sample in the thin Iayer between the IR window
and the eiectrode absorbs some of the radiation, changing the profile of the interferogram
The interferogrms for both the background and the sample are stored in the computer,
and are used to calculate the transmission or potential difference spectra. The computer
also controls the potential on the potentiostat. The detailed diagram of the
electrochemical cell used for SNIFTIRS is shown in Figure 3-3.
SNIFTIRS was carried out using a Nicolet 20X/C FTIR apparatus equipped with a
MCT-B detector cooled by liquid nitrogen. The sample cornpartment of the apparatus
was purged throughout the experiment using COz and H20 f?ee air provided by the
Puregas Heatless Dryer. The electrode potential was controlled by a PAR 173
potentiostat.
The SNLFTIR spectra were acquired using a multiple potential step (MPS)
procedure. In the MPS the electrode potential was stepped periodically between the
reference potential El and the sarnple potential E2.During each step, n interferograms
(usually 100) were acquired at the potential El and E2. The acquisition was delayed for 30
seconds afier each potential change to d o w the interface to reach a thermodynarnic
equilibrium at the new potential. The change o f the electrode ptential was synchronized
with the acquisition of the interferograms by connecting the extemal trigger port of the
PAR 173 potentiostat to the communication port of the DX 486 computer. This procedure
repeated m times ( u d y 20) so that N = n x rn
= 2000
interferograms were acquired at
each of the two potentids in order to improve the signal-to-noise ratio. The
interferograms were added, Fourier transformed and used to calculate a relative change of
the electrode reflectivity which is defined by equation 2.37. The shape of the final spectra
could be mono-pofar or bipoiar depending on the different modes of interactions among
the eiectrode, the beam and the adsorbed molecules (see Figure 2-6). The spectra here
Working Electrode
Reference Electrode
-
Counter Electrode
Thin layer of
Electrolyte
Window
Figure 3-3 The electrochemical ce11 for SNIFTIRS with CaFz prism configuration.
(Adapted fiom ref [SI.)
were rewrded at a resolution of 4 cm-'. Both a CaFt prism and a ZnSe hemispherical
window were used in this thesis- For the former, the IR incident beam was normal to the
prism, but at an angle about 70' to the normal of the electrode s d a c e after it passed
through the thin layer (cl 0 p)between the window and the electrode surface. For the
latter, the IR beam was adjusted to an incident angle of 33' at the hemispherical surface,
but at an angle of about 80' to the normal of the electrode surface. The use of the ZnSe
window ailowed a lower cut-off frequency of 1000 cm-' for the SNIFTIRS studies of
DMPC at the gold electrode. The hemispherical shape also collimated the incident beam.
Thus the throughput was irnproved- Although the hernispherical window usually shows
better sensitivity and signal-to-noise ratio, it is not suitable for the s-polarized photons
because, at the angle of incidence for p-pdarized photons, only about 30% of s-polarized
photons enter the thin layer. For this reason only p-polarized spectra were collected for
DMPC molecules at the electrode surface. However, the p-polarized spectra provided
enough surface information for our current interest.
3.4 Deposition of the DMPC monolayer and bilayer
3.4.1 Preparation of SUVs and detemination of vesicle size
S W s (small unilamellar vesicles) were prepared roughly according to the Barenholz
procedure [9].Briefly, 1 mL of a IS mg
stock solution of DMPC in chloroform was
evaporated under a strearn of nitrogen ont0 the walls of a test tube. The dried
phospholipid was desiccated under vacuum for 1
- 2 hours to remove any trace of
solvent. 2.00 mL of 0.150 M KC1 or 0.05 M NaF solution was added to the tube and the
phospholipid was suspended by gentle swirling. The suspension was sonicated to clarity
in a bath sonicator at 30°C. The solution was ailowed to warm up durhg sonication and
solution temperature was maintained at 30°C after sonication.
Determination of the distniution of vesicles sizes was perfonned by dynamic light
scattering (DLS) as per Hallett et al- [10,11]. Light fkom a helium-neon laser (Model
125, Spectra Physics, Mountainview, CA) was focused into a scatterhg chamber. The
scattered light was collimated by two pinholes and detected by a photomultiplier (Model
9863, EMI Electronics Ltd., Hayes, England), a quantum photometer (Model 1140,
Princeton Applied Research, Princeton, NJ), and analyzed with a correlator (LangleyFord Instruments, Amherst, MA) us h g in-house fitting sohvare. These rneasurements
were performed at room temperature (below the TM for DMPC).
3.4.2 Preparation of DMPC monolayers and detemination of the film pressure
DMPC is an amphiphiIic molecule. It forms a monolayer at the gas-water intefice
by self-assembly. The polar heads of phospholipids are anchored to the surface of the
water solution and the non-polar taiis extended into the gas phase. 1 @ of 15 mg d-'
DMPC stock solution (in chloroform) was injected ontc the surface of aqueous 0.05 M
NaF solution (T = 30 OC) in an electrochemical celi whose inside diameter was about 4
c m Argon was gently flowing over the water sufiace for about 30 minutes to ensure that
the chioroform was evapotatec! completely. The film pressure of DMPC at the gas-water
interface was measured with a Wilhelmy plate attached to a LB microbalance. The
microbalance was controlled by software developed for the Langmuir trough (TSV,
LB5000). The output fiom the cornputer was a set of film pressure data as a function of
tirne. The cleanness of the solution and glassware is critical for these measurements.
I
The monolayer of DMPC ai: the metal-solution interface was deposited using the
single horizontal touch technique [12-141 described in section 2.4.2. The monolayer at
the metal-solution interface is characterized by recording CV, differential capacity and
charge density curves. The charge densities on the DMPC monolayer covered electrode
are detennined using the chronocouIometry technique (see section 3.5.3 for more details).
The film pressures were calculated Born the charge density data using equation 2.13 and
the procedure described in section 2.1.3. If the maximum value of the film pressure at the
metal-solution interface is very close to the pressure of DMPC monolayer at the gassolution interf~ce,the transfer of DMPC is considered as successfiil3.4.3 Deposition of DMPC bilayers
Three different methods have been used in this thesis to deposit the DMPC bilayers,
(0
double horizontal touch, (ii) combined LB and horizontal touch and (iü) vesicle
h i o n . A comprehensive description of these procedures has been given in section 2.4.2.
This section is only described the experimental procedures.
Double horizontal fouch
-A
DMPC monolayer was prepared as described in
section 3.4.2. A Au(ll1) electrode was lowered to contact the solution covered by a
monolayer of DMPC and then lifted up to the argon environment. The electrode was then
horizontaily touched again to the film covered gas-solution interface- Unfortunately it is
not possible to deposit a bilayer of DMPC ont0 the Au(i 11) electrode swface using this
method.
LB deposition - The electrode was bent so that the Au(ll1) surface was at a 30'
'angle to the normal of the solution surface. The bent electrode was immersed into the
0.05 M NaF solution (figure 3-4a). Its potential was set at 200 mV (SCE) by a
potentiostat. A compressed DMPC monolayer was formed at the gas-solution interface as
described in section 3A.2. n i e electrode was pulled out of the water surface by a Dipper
on the microbalance at a vertical speed of 0.2 mm min-'.
A monolayer was then
transferred ont0 the electrode surface with the polar heads attached to the s d a c e (figure
3-4b). The electrode was straighted out and Iowered to touch the film covered water
(a)
Figure 3-4 A schematic procedure (a-b-x)
0)
(4
to make DMPC bilayers by a combination
of LB deposition and horizontal touch techniques.
sudàce again. A DMPC bilayer with a tail-to-tail orientation of the phospholipid
molecules is deposited at the metal-solution interface using this approach (figure 3-4c).
The apparatus in figure 2-7 was used to coat the g l a s slides. To accomplish the
deposition, the siides were raised and lowered through the gas-solution interface using
the Dipper above.
4
Vesicle fuson - The DMPC vesicles were allowed to spontaneously spread on the
substrate surface either by the immersion of the substrates into the SW solution (1 mg
DMPC/mL, rnaintained at 30 OC) for approximately 1 h or by addition of vesicle solution
(200 ILL)of 15 mg/mL DMPC to the electrolyte solution (-30 mL) in the electrochemical
ce11 (also maintained at 30 OC). To mat a glass slide in situ in the Atomic Force
Microscopy (MM) liquid ce11 the cell was first filled with water. Next, the ce11 was
filled with 1 mg m ~ SUV
"
solution in 0.15 M KCL md allowed to be present in the cell
for 30 minutes. Finally, the cell was flushed gently with 90 mL of 0.15 M KC1 solution to
remove the vesicles.
3.5 EIectrochemical methods
3.5.1 Cyclic voltammetry
Cyclic voltaxnmetry has b e n used for the following four purposes. i ) Cyclic
voltammograms of the pure suppoaing electrolyte were used to d e t e d e the potential
range in which the electrode is ideally polarizable. The range for Au(1ll) in 0.1 M
KC1O4 and 0.05 M NaF was -800
- 600 mV (SCE) and -800 - 550 mV (SCE)
respectively. ii) The oxide formation peaks of single crystal electrodes are characteristic
of the crystallographic orientation, therefore the cyclic vo ltammogam could be used as a
fmgerprint to ver@ the quality of the electrode surface investigated. üi) The shape of the
cyclic voltammogram gives information on the presence of oxygen, hydrogen and other
irnpurities in the buk solution. The cyclic voltammogram displays 'Yilts" character when
a significant amount of oxygen is present in the solution A proton reduction peak below
-700 mV (SCE) appears on CV in acidic solution and unusal features indicate the
presence of impurities. iv) The cyclic voltammogram recorded in the solution containing
organic molecules is used to obtain qualitative information about the adsorption of
organic molecules. It helps to choose the sarnpling potentiak in chronocuulometry and
SNIFTIRS. In the cyclic voltammetry studies, a PAR (Princeton Applied Research)
Model 173 potentiostat (with a built in PAR Model 176 I/E converter) comected to a
home-made programmable voltage generator was used. The current was digitized and
stored in a personal cornputer (IBM PC microcornputer with a Computerscope ISC-16
multichannel board). All the CV curves shown in this thesis were recorded at a sweep
rate of 20 mV s"
.
3.5.2 Differen tial Capacity
The double layer at the metal-solution interface can be viewed a s a capacitor. A 5
mV (root mean square) and 25 Hz ac signal fi-om the fiinction generator was
superimposed on a potenta ramp (5 mV/s) fiom the potentiostat. The resulting ac current
was analyzed by a lock-in amplifier into the reai and imaginary components which were
fed to the mmputer and used to calculate the differential capacity according to the
foHowing equation:
where C is the capacitance, V,, is the potential of the ac signal, i,is the rea1 component of
the ac current and ,i is the irnaguiary component of the ac current. The reai component of
the ac current (i,)is in-phase with the Va,signal, and the irnaginq component of the ac
current (i,)is 90" out-of-phase with the Va,signaI.
The differential capacities measured in the absence of the organic adsorbate were
used to deterrnine the potential of zero charge (pzc). The pzc was determined fiom the
#
position of the diffuse layer minimum on the dserential capacity curves for the pure
electrolyte solution The measurements should be repeated in two different
concentrations of the supporting electrolyte. The pzc of the Au(ll1) electrode in this
thesis is 290 mV (SCE). The dfierential capacity curves were also used to identiQ the
total desorption and the omet of adsorption of the organic motecules.
3.5.3 Chronocoulometry
Chronocoulometry was used to measure the difference between the charge densities
on the metal side of the interface AG^) at a potential Ei, where the organic mo lecules are
adsorbed, and at a potential Er,where total desorption of organic mo lecules takes place.
In chronocoulometric experiments, the potential of the electrode was held at
Ei for a
period of time long enough for the equilibrium between the interface uld the buk of the
solution to be established (usually 1 to 5 minutes depending on the concentration).
Solutions with adsorbate concentration lower than 104 M were stirred. The s t i r ~ gwas
interrupted 60 sewnds before the data were acquired. The potential was then stepped to
Er and the current due to double layer charging was measured and stored in the
microcornputer using the ISC-16 board. The above process was repeated after another
equilibrium has k e n reached at Ef.The potential Ei varied with 25 mV hcrements each
tirne to cover the whole double layer region, and thus a family of current transients was
deterrnined (Figure 3-Sa). The current transients were digitally integrated to give the
charge transients (Figure 3-5b). The chronocou10metric curves display a fast rlsing
section followed by a plateau. The charge in the plateau region fiequently displays weak
tirne-dependence due to the presence of instrumental offsets a d o r parasitic faradaic
currents (reduction of traces of oxygen or evolution of hydrogen). This dependence
usually has a Iinear character and the contribution Eom the offsets and/or parasitic
currents can be eliminated easily by extrapolation to t = O. The intercepts give the
required values for the charge difference. In order to perform the thermodynamic analysis
the above measurements have to be repeated in solutions with different bulk
concentration of organic n o lecules (usually fi0rn 1O" to 1o5 M) .
(a>
(b)
Figure 3-5 (a)A family of current transient curves measured by chronocoulometry. @) A
family of charge density curves calculated fÎom (a).
The absolute charge density can be calculated fiorn the measured value of relative
charge density, A o M with the help of the p z determined independently f?om the position of
the diffuse layer minimum in the dserential capacity curve for a dilute solution of the pure
.J
supporting electrolyte. Once the pu:is known for a solution in the absence of organic solute,
the absolute charge density at potential E is calculated fiorn the relative charge density a? the
p z according to the following equation:
It is known that at Efno adsorption of organic solute takes place. It is reasonable to assume
that the absolute charge density at & is the sarne for a solution in the presence and in the
absence of organic molecules. Therefore, the absolute charge densitiies at any Ei for both
solutions in the absence and presence of organic molecules can be calculated fiom AsM(Ei)
and A c T ~ ( ~ z cusing
)
the following equation:
o ~ ( E i=) A g ~ ( E i ) A- CTM@C)
(3.3)
The absolute charge density has been used to do the therrnodynamic analysis described in
section 2.1.
References
1. R Parsons, Proc-R.Soc. London Ser. A, 261 (1961) 79.
2. B.B. Damaskin, O.A. PeLrü and V.V. Batrakov, Adsorption of Organic Compounds on
Electrodes, Plenum Press, New York, 1971, p70.
3. J. Torrent and F. Sanz, J.Electroanal.Chem. 286 (1990) 207.
4. J. Torrent and F. Sanz, J.Electroanal.Chem 359 (1993) 273.
5.R. Payne, J.Electroanal.Chem. 41 (1973) 277.
6. B.B. Damaskio, A.. Fnimkin and A. Chizhov, J. Electroanal. Chem. 28 (1970) 93.
7. D. Dickertmann, J.W. Schultz and F.D. Koppitz, Electrochim. Acta, 21 (1976) 967.
8. B. Beden, C. Lamy, Spectroelectrochemistry: Theory and Practice (R-J. Gale Ed),
Plenum Press, New York, Vol. 22, 1988.
9. Y. Barenholz, D. Gibbes, B.J. Litman, J. Goll, T.E. Thompson and F.D. Carlson.
Biochemistry, 16 (1977) 2806,
10. F.R Hallett, T. Craig, and B. Nickel. Can. J. Spectrosc. 34 (1 989) 63.
11. F-R Hallett, I. Watton, and P Krygsman, Biophys. J. 59 (1991) 357.
12. D. Bizzotto, J. Noel, J. Lipkowski, Thin Solid Films, 248 (1993) 69.
13. D. Bizzotto, J. Noel J. Lipkowski, J. Electroanal. Chem. 369 (1994) 259.
14. D. Bizzotto, A. McAlees, J. Lipkowski, and R McCrindle, Langmuir, 11 (1995)
3243.
Chapter 4 Electrochemical Studies of Benzoate Adsorption on
the Au(ll1) Electrode
4.1. Introduction
This work is a part of systematic study of the adsorption of anions [l-71and
neutral organic molecules [8] at the gold-solution interface. Benzoate is an organic anion
and its interfacial behaviour should display the characteristics typical of both an anion
and an organic molecule. Benzoate is reported to be an anti-corrosion agent [9,1 O], and
hence information conceniùig its surface CO-ordinationmay fmd practical applications.
There are a number of papers devoted to benzoate adsorption at electrode surfàces.
Shilotkach and Gil'rnanshina [ I l ] studied the adsorption of benzo ic acid and benzoate on
the bismuth electrode by radiotracer and dserential capacitance measurements. Benzoic
acid adsorption at poIycrystailine gold was described by Zelenay et al.[12]with the help
of a radiochernical technique. The behaviour of benzoic acid at gold electrodes was also
investigated by conductivity measurements [13], surface enhanced Raman spectroscopy
Cl41 and potentiai difference IR spectroscopy 1151. There are a few papers devoted to
benmate and benzoic acid adsorption at polycrystaiiine Pt electrodes 115-181. In addition
benzoate adsorption on Cu(ll0) has been studied in vacuum employing IR spectroscopy
[l9,2O].
The objective of these studies is to describe the coordination of the benzoate ion
to the Au(1 Il) electrode surface by a combination of eIectrochemica1 and IR
spectroscopie techniques. The structure of the benzoate is relatively simple. However, it
contains both hydrophobie ( b e n n e ring) and hydrophilic (carboxylate) groups which ore
characteristic for phospholipids and some of the important amino acids. Studying the
adsorption of benzoate will surely provide informative insights to understand the
interfacial behaviour of complicated bio logical molecules such as phenylalanine (see
Chapter 6). In this chapter we will describe thermodynamics of benzoate adsorption at the
gold single crystal electrode ushg cyclic vo ltammetry, dinerential capacity and
chronocouiometric measurements. The results of spectroscopie experiments are presented
in the next chapter (Chapter 5). To facilitate a cornparison of the present results with the
data for adsorption of inorganic anions acquired earlier, the structure of this chapter and
the method of data presentation foUows previous publications [1-71 in our lab.
4.2. Experimental
The experimental methods and procedures have k e n described elsewhere [22,23]
and also in Chapter 3 of this thesis.
43. Results and Discussion
4.3.1. Electrochernical rneasurements
The adsorption of benzoate at the Au(ll1) electrode surfàce was i8itially
characterized by means of cyclic voltammetry and differential capacity measurements.
The inset to Figure 4-1 shows cyclic voltammetxy curves recorded in the double layer
region of the Au(l11) electrode, in O. 1 M KClO4 (dotted line) and in the presence o f 1.58
x 1o5 M potassium benzoate (solid line). The pair of quasi-revenible peaks between 300
and 400 mV results fiom adsorption and desorption of benzoate ions. The peak
-800
-600
-400
-200
O
WmV vs SCE
Figure 4-1 DifEerential capacity recorded at an Au(1 I l ) electrode in 0.1 M KClO4
(dotted line) and in 0.1 M KClO4 + 1.58 x 1oa3M potassium benzoate (solid Iine). The
inset shows the double layer section of CV curves recorded in the same solutions at a
sweep rate of 20 mV S-'.
recorded in the positive voltage scan is quite broad compared to the voItammetric peaks
associated with of adsorption of inorganic anions, for example Cl- [4]. The difEerential
capacity curve presented in the main section of this figure shows that this peak has a
shoulder at potentials between 100 and 200 mV. These features suggest that benzoate
adsorption is a two-state process. In fact the segment of the CV curve recorded during the
negat ive vo Itage scan displays two well-separated peaks, Apparently, desorption of the
benzoate ion proceeds through two well-reso lved steps. The reflection IR spectroscopy
studies described in ref. [21] demonstrate c1early that the shouIder or a srnall peak seen
on the CV or the dserential capacity curves at E -150 mV corresponds to a flat nbonded d a c e orientation of the adsorbed mo1ecuIes~The taIl peak seen at -300 mV
corresponds to a vertical or tilted orientation in which the benzoate rnolecule orients with
the carboxylic group towards the metal and the arornatic ring towards the solution. The
orientation of the adsorbed benzoate changes with the eiectrode potentiaL The IR studies
also showed that no oxidation or dissociation of the benzoate ion takes place when the
electrode potential is lower than 600 mV.
It is now well established that the Au(l11) electrode surfke is reconstmcted
when it is negatively charged and that the reconstruction is lifted at positive charge
densities (for review see ref [24,25]). It is also known that adsorption of inorganic anions
assists in lifting the surface reconstruction [24-281. For korganic anions the characteristic
tdl peak in CV or dzerential capacity curves corresponds to coupIed phenornena of
ionic adsorption and lifting the surface reconstruction [26-281. We have attempted to
determine the structure of the electrode surfàce in the presence of benzoate with the help
of STM. Unfortunately, we did not succeed in acquinng atomic resolution images of the
gold surfâce in the presence of benzoate ions. ~ o n s e ~ u e e t it
b ,is unclear how tmzoate
adsorption affects the surface crystallography of the Au(l11) electrodeThe CV or the daerential capacity curves in the Presence of potassium h x ~ o a t e
merge with the correspondhg cunres of the pute supporthg electrolyte at E c -400 mV.
This behaviour indicates that the benzoate ions are totalb desorbed fiom the Au(l I 1)
sUTface at potentials more negative than -400 mV. We \uere therefore able to determine
the charge density O f the electrode surface fiom chronocoulometric measurements using
the procedure described in ref [22,23]. The reference potential @otential of total
desorption) in these experïments was chosen as -750 mV. Figure 4-2 shows a family of
charge densities (aM)versur potential plots determined hthese experkents- The charge
density increases when the electrode potential increases and when the bulk knzoate
concentration increases. In addition, the potential of zef0 charge shifis in the negative
direction as the buk concentration of benzoate increases. This khaviour is t~picalfor a
specfic adsorption of anions [6-91.
The IR spectroscopie studies descnkd in 1211 damnstrated that h z ~ a t e
assumes a fiat x-bonded coordination at E < 200 mV ida verticai tiited orientation
when E > 200 mV. The charge densities in Fig. 4-2 show that the flat orientation and the
l~
surface,
vertical (tilted) orientation are observed at a negatively and ~ o s i t i v e char@
respectively. The orientation of the adsorbed benzoate is apparently c h g e d by moving
across the potential of zero charge (pzc).
In Figure 4-3, the dinerential capacity d e t e d e d fiom the ac im~edance
measurement is cornpared to the capacity calculated fkofnthe ~ositiv*gohg section of
voltammograms and capacity calculated by differentiation of charge densit~data
-800
-600
-400
-200
O
200
400
600
E/mV vs SCE
Figure 4-2 Charge density versus electrode potential plots for the Au(ll1) electrode in
0.1 M KC1O4solution (dotted line) containhg dii3erent molar concentrations of benzoate:
1.28 x 10.' (filled circle); 2.56 x 10-'(open circle); 4.00 x 10-~(filledsquare); 6.66 x 10"
(open square); 9.30 x 10-'(filled up-triangle); 1.59 x 104(open up-triangle); 2.86 x 104
(filled down-triangle); 5.56 x 1o4(open dom-triangle); 5.79 x 10~(filleddiamond); 1.58
x 1o - (open
~
diamond); 2.91 x 1 ~ - ~ ( f i l l ehexagon);
d
4.21 x 1 0 - ~ ( o ~hexagon).
en
-600
-400
-200
O
200
400
WmV vs SCE
Figure 4-3 DiEerential capacity curves for the Au(ll1) electrode in 0.1 M KC1O4 + 1.58
x 105 M benzoate solution determined fiom: ac impedance experiment (square); CV
curve (triangle); differentiation of the charge density curve (circle). The insert shows the
charge density determïned fi0 om: integration of the single-frequency differenti d capacity
curve (square); integration of CV curve (triangle); chronocoulometry (circle).
determined fiom chronocoulometric experiments. The differences between the capacity
curves calculated fiom CV and chronocouIometric data are quite small. The differences
between the two Iater curves and the dserentiai capacity measured fkom the ac
impedance are more pronounced. The inset to Fig. 4-3 compares the charge densities
determined by chronocoulometry to the charge densities detennined by integration of the
differenti d capacity corresponding to the ac impedance or cyclic vo ltarnmetry
measurements. The charge density data measured by chronocoulometry are apparently
higher than charges determined by the two other techniques. These daerences ïndicate
that the state of adsorption equiiibrium has not been attained during the measurement of
the CV curve or the ac impedance. For this reason the quantitative data analysis reported
below was based on the data derived fiom chronocoulometric experiments.
4.3.2. Gibbs excess and Gibbs energy
The film pressure of adsorbed benzoate c m be calculated fiom the charge density
data obtained fkom chronocodometry measwements Pig. 4-2). The area wntained
between the curve corresponding to a given benzoate concentration and that of the pure
electrolyte solution is equal to the film pressure 122,231:
The subscripts 8 and 0=0 denote the presence and absence of benmate in the bulk
solution and Ef and Ei correspond to final (complete desorption) and initial (variable)
potentials, respectively. The dflerentiation of .rr vs. Zn c plot gives the relative Gibbs
excess (171.
where c is the bulk concentration of benzoate in O. 1 M KC&
A farnily of the ï YS. E plots for various buk benzoate concentrations is presented
in Fig. 4-4(a). The surfàce concentration of benzoate changes graduaiIy with potential
until it attains a plateau. The maximum surface concentration of benzoate corresponds to
7.3 x 10-Iomol cme2.This value is in a reasonable agreement with the number 8.5 x IO-"
mol cm-2 for the maximum surface concentration of benzoic acid at polycrystalline A y
detemiined using radiotracers by Zelenay et al. [12]. It is ody somewhat lower than the
theoretical value dculated fiom the cross sectional area for the vertically adsorbed
molecule (0.198 nm2) that is 8.4 x 10-'O mole
At potentials lower than 200 rnV,
where benzoate molecules assume n-bcmded orientation, the surface concentration is
lower than 20% of the limiting surfàce coverage. The transition from the flat to vertical
(tilted) orientation is graduaL In contrast to the adsorption of berizoic acid on
polycrystalline Au 1121, no well dehed plateau attriiutable to a flat (n-bonded state) is
seen on the r versus E plot for benzoate on the Au(Z11) electrode. The asymmetry of the
r versus E plot is the only indication of the progressive reorientation of the adsorbed
benzoate ion with potential.
Fig. 4 4 b ) shows the surface excess data plotted against the charge density on the
metal. The middle sections of these plots are fairly linear. Their slopes give the Esin-
Markov coefficients that wiil be discussed in section 4.3 -3.
In order to detennine Gibbs energies of adsorption, the surface pressure (O) data
are fitted to the "square root " isotherm [Zg730].
-200
O
200
400
WmV vs SCE
600
80
70 -
60 E
50 O
9
-E 40 -
30 -2 20
Y
10 -
0-
Figure 4-4 Plots of surfàce excess o f benzoate on the Au(ll1) electrode against (a) the
potential and @) electrode charge density for 0.1 M KClo4 with different molar
concentrations of benzoate: 1.28 x l0"(filled circle); 2.56 x 10~~(open
circie); 4.00 x IO-'
(filled square); 6.66 x 10~~(open
square); 9.30 x loJ(filled up-triangle); 1.59 x 1o4(open
up-triangle); 2.86 x 1 o4 (filled down-triangle); 5.56 x lo4(open down-triangle); 5.79 x
104(filled diarnond); 1 S 8 x 1o5 (open diamond); 2.9 1 x I 0'.'(filled hexagon); 4.2 1 x 1 o5
(open hexagon).
~ n ( m ) + ~ n=pI . u Q > + B ~ + ~
where c is the bulk concentration of bermate, f3
=
(4-3)
exp(-AG~T) is the adsorption
equilibrium constant, B is a constant and 0 is the surfhce pressure- m e n analysis is
carried out at a constant potentiaI, @ is set as ye=o - ye, where y is surface tension. If we
take charge as the independent electrical variable,
is set as
ce-o- 59, where 5 is Parson's
fûnction Fig. 4-5(a) and 4-5(b) show the plots of the d Rvs. h&Tc/@)- The plots are
fairly linear and when extrapolated to zero surface pressure their intercept with the
ordinate is equal to:
lim[ln(kTc/Q)
= -ln
P
(4-4)
The Gibbs energies of adsorption detemiined by this method are plotted against potential
in Fig. 4-6(a) and against charge density in Fig. 4-6@). The standard state is an "ideal" ï
= 1 ion
for the adsorbed species and an "ideal" c = 1 M for the buk species. The
extrapolation to zero surfàce pressure ïs very long- ConsequentIy, even a small change in
the slope may result in a significant error in the intercept. Therefore it is usenù to
compare the Gibbs energies determineci independently by changing the electrical
variables. For that purpose, the Gibbs energy determiined at constant charge is converted
to the corresponding AG vs. E plot (circle in Fig. 46(a)) using the DM vs. E plot for pure
supporthg electrolyte. For potenrials lower than 250 mV the agreement between the two
sets of data is very good For E > 250 mV the two sets of data points diverge. This is the
region of higher coverage, where the extrapolation is much longer. These differences
indicate that the Gibbs energy data contain a systematic error.
Figure 4-5 Fitting the benzoate adsorption data to the "square root " isotherm (Eq. 4.3) at
(a) constant potentids and (b) constant q.
*
LO
O
4
rn
@Y
O
ln
Cr)
CV
T
F
Figure 4-6 Gibbs energy of adsorption for benzoate on the Au(ll1) electrodz in 0.1 M
KC104: (a) AG vs. E plots determuied fiom Fig. 4-5(a) (square) and fiom Fig. 4-6
(b)(circle); (b) AG vs. aMpIot determined IiomFig. 4-Se).
P
4.3.3. Electrosorption valency and Esin-Markov coefficient
The fkst derivative of AG vs. E gives the electrosorption valency (y').
The
electrosorption valency can be determined independently fiom the dope of the charge
density vs- Gibbs excess plots:
Figure 4-7 shows a plot of a~ vs.
r for various electrode potentials. The data display a
fairly linear relation between charge and coverage. The slopes of these plots can be
compared to the first derivatives of the AG vs. E plots. The electrosorption valencies
calculated by these two independent routes are plotted in Fig. 4-8(a). To assess the values
of Gibbs energies detennined at constant charge and constant potential, the two sets of
Gibbs energy data were merentiated separately. Electrosorption valencies determined
fiom the diEerentiation of the Gibbs energies at constant potential are much lower tban
the values obtained fÎom the dope of o~ vs.
r plots. In contrast, the derivative of the
Gibbs energies at constant charge gives electrosorption valencies that are in a better
agreement wiîh the slopes of
vs.
r
pbts. This analysis indicates that the Gibbs
energies detennined at constant charge may be trusted more than the energies determined
using potential as the independent eIectrical variable. It is usehl to stress that the squareroot isotherm is an empirical isotherm that is not based on a well-defined physical model.
It is a convenient procedure to linearize the experirnental data Consequently, the Gibbs
energies determined by this method may be afEected by a systematic error. The above
analysis shows that the cross dzerential relationships could be a quaiity test for Gibbs
energy data,
Figure 4-7 Plots of the
vs. ï
of benzoate at constant electrode potenitials/mV (SCE).
113
O
100
300
400
Figure 4-8 Electrosorption valencies determined fiom: (0 the slope o f the
vs. l? plots
-100
2010
E/mV vs SCE
(cùcle); (ii) AG versus E plots (square) using AG values determined fiom the analysis at
constant potential; (iii) AG versus E plots (triangle) using AG values determined fiom the
andysis at constant charge.
The second cross differential relationship gives the Esin-Markov coefficient:
The derivative on the left side of Eq. 4.6 is given in terms of directly rneasured quantities
such as E and
CM.
The derivative on the nght side of Eq. 4.6 is given in terrns of the
surface excess, a quantity whose calculation involves one integration and one
differentiation step. This relation rnay therefore be used to check whether the caicuiations
of the Gibbs excess are fiee of systematic errors. Figure 4-9 shows the plots of E versus -
RT(h c) for various electrode charge densities. The plots are fàirly linear and thek slopes
give the Esin-Markov coefficient. Independently the plots of r vs. m. in Figure 4-4b can
be Mted by a polynomid and differentiated numerically to give (X/ ao,,,)
. The two
derivatives determined independently are plotted against each other in Fig. 4- 10. The
points representing calculated values are closely scattered around the dotted h e whose
slope is ULLity. This behavior indicates that the
r
values are fiee fiom major data-
processing errors.
The middle sections of T vs. CTM. in Figure 4-4b are fàirly linear and may be used
to calculate the number of electrons flowing to the interface per one adsorbed benzoate
ion at a constmt benzoate concentration in the bulk. The reciprocal of the slope of the
middle section of the
r
vs. o~ plots is equal to -1.1 electron per ion. This value is
somewhat higher than electrosorption valencies reported in Figure 4-8.
12
14
16
18
20
-RTln(c,,)
22
24
26
28
30
/ kJ mol-'
Figure 4-9 Esin-Markov plots (E vs. -RT In c ) for benzoate on the Au(ll1) electrode at
charge densities show in the figure.
6
7
8
9
-1 ~ ' ~ ( d ~ / d moi
&/ p ~ "
Figure 4-10 The cornparison of Esin-Markov coefficients determined f?om the siope of
the fitted Iines in Fig. 4-9 and by the differentiation of the r versus o~ plots. The dotted
line has unity slope.
4.3.4. Mode1 of the double layer
In order to extract information conceming the structure of the interface and the
charge distribution at the interface a specific physical model has to be used to interpret
the thermodynarnic data Following our earlier work in this series Cl-7,281, we will
employ the Graham-Parsons model [3 11 of the inner layer. The inner-layer capacity
ci
can be calculated fiom the overall electrode capacity C determhed by differentiation of
the charge density curves and using the theory of the diffuse layer with the help of the
formula [23] :
--[+))Lao
1 = 1
-
CI c
where
cd
cdiS the capacity of the diffuse layer and a r l i ? is~ the
~ slope of the r versus CM
plots in Fig. 4-4@). The value of F(arIi3oM)is close to unity for d concentrations of
benmate and for most of the electrode charge densities. In most cases, the second tenn on
the rïght side of the Eq. 4.7 can be neglected and the inner-layer capacity may be
approximately taken as equal to the overall capacity. Fig. 4- 11 shows the inner-layer
capacities calculated with the heIp of Eq. 4.7 pbtted against the charge density on the
metal, Those values of capacity are ahost independent of the bulk concentration of
benzoate and are essentially equal to the differential capacity of the overall double layer.
The inner-layer capacity is a fim~tionof two variabies, the charge on the metal
and the amount of adsorbed anion. Therefore it can be expressed in terms of the two
components, the inner-layer capacity at a constant charge rC and the inner-layer capacity
at a constant amount adsorbed ,C as foliows [3 11:
The capacity rC is equal to the slope of the plot of the potential drop across the inner
layer ( ~ 4 =~ E"- E,
- 42) versus the charge (-FI) of adsorbed benzoate at constant CM
(42 is the outer Helmholtz plane potential). The values of ~ 4 are
~ plotted
"
aga% -FTin
Fig. 4-12. Although these relations are quite short, they are fairly h e m and their slope
&. The values of rC were then used to calculate ,C with the help of Eq. 4.8. The
and
rC deterrnined following this procedure are plotted against CM in Fig. 4-1 3. The shape of
these two curves is quite similar. They display a maximum at small positive charge
densities. This behaviour is quite similar to that observed for specifk adsorption of
inorganic anions on Au(l11) [28]. However, we note that for benzoate, the values of the
diffierential capacity at constant charge and at constant amount adsorbed are significantly
lower than for inorganic anions, This behaviour may be explaineci by the iarger size of
the benzoate ion.
The capacities at constant charge and constant amount adsorbed may be
wnsidered as integral capacities described by [3 1,331:
r C = I~( x -~x ~ )
,C=E/X~
where E is the permittivity, x: and
x2
(4-4)
(4.1 O)
are the distances of the b e r and the outer
Helmholtz planes fiom the metal surface, respectively. The capacity at a constant charge
rnay be affected by both a change in the permittivity and the position of the inner
Helmholtz plane
XI.
The interpretation of ,C is easier, although in the present case the
thickness of the inner layer xt increases by moving fkom negative to positive charges due
-30 -20 -10 O
10
20
30
40
50
60 70
80
o,/ FC cm"
Figure 4-11 Inner layer capacities of the Au(1ll) electrode in 0.1 M KC104 containhg
different rnolar concentrations of benzoate: 5.79 x 1o4 (circle); 1 -58 x 1 o ' ~(square); 2.9 1
x 1 O" (up-triangle);4.2 1 x 1o5 (dom-triangle).
O
-5
-10
-15
-20 -25
-30 -35
-40 -45
-m/ pc cm"
Figure 4-12 Plots of the potential &op across the inner layer (
~ 4 ~versus
- ~ ) the charge of
adsorbed benzoate (-ET)at constant charge density on the metal side of the interface as
indicated in the figure.
O
10
G,
20
/ p~ cm"
Figure 4-13 huer layer capacities at constant charge (rC) determined firom Fig. 4-12
(circle) and at constant amount adsorbed (,C) determined fiom Eq. 4.8 (square) for the
adsorption o f benzoate on the Au(111).
to the reorientation of the adsorbed anion However, an increase of x2 should cause a
decrease of the capacity. Therefore, the maximum on the ,C curve displays chiefly a
change of the inner layer permittivity. The changes of the ùiner layer permittivity reflect
changes in the orientation of the surface water dipoles. The maximum on the ,C curves is
usually interpreted as being due to the field induced solvent reorientation. It is observed
at the charge where the disorientation of water dipoles is at maximum. The maximum
solvent disorientation should take place around the pzc, in Fig 4- 13 it is observed at -8
p C cm-'.
The ratio of the inner-layer capacities at a constant amount of adsorbed anion and
a constant charge is equal to the electrosorption valency w , 3 3 ] :
Electrosorption valencies determined fiom the inner-layer capacities using Eq. 4.1 1 are
shown in Fig. 4-14, They are compared to the electrosorption valencies calculated fiom
slopes of
vems
r
plots in Fig. 4-7. The agreement between the electrosorption
valencies calculated by the two methods is very good and indicates that no major errors
were made in the data processing.
The poiarity o f the eIectrosorption bond may be conveniently described in teof the surface dipote [33-371. n i e surface dipole represents a dipole formed by an
adsorbed benzoate ion and its image charge in the gold electrode. Its magnitude is a direct
measure of the polarity of the bond formed between the benzoate and the electrode. The
surface dipo le can be calculated by [33] :
-10
-8
-6
-4
-2
O
O,
2
4
6
8
10
12
14
/ pC cm"
Figure 4-14 Plots of electrosorption valencies determined fiom the CM vs. r plots (circle)
and f h m the ratio of ,C to rC (square).
-5
O
5
o.,/ PC cm"
Figure 4-15 The surfàce dipole moment formed by benmate (square) and Cl-(circle) [4]
adsorbed on the Au(l11) electrode.
where eo is the charge of an electron (1.6 x 10'19 C ) and
E
is the permittivity of the inner
layer. Usually, the permittivity of vacuum (8.85 x 10'12
(?T' II?) is
used U1 these
ca1cuIations. The surface dipoles for benzoate at Au(1 l 1), cdcrilated fiom the
electrosorption vafencies are plotted in Fig. 4- 15.
For cornparison, the variation of
surface dipole for Cl-, taken ftom our earlier work [4], is also shown in Fig. 4-1 5. The
two curves display very simi1a.r behaviour. The surface dipole is much larger at a
negatively than at a positively charged surface. Apparently, the polarity of the
chernisorption bond changes signifïcantly with the charge on the metal. At a positively
charged surface the surface dipole attains a very low value of 0.3 C m ion-'. Such a low
value of the surface dipole may either indicate charge transfer fiom the benzoate ion to
the metal or an efficient screening of the charge on the anion by the charge in the metal.
4.4. Summary and Conclusions
We have described adsorption of benzoate at the Au(ll1) electrode s u f k e in
terms of Gibbs excesses, Gibbs energies of adsorption and numbers of electrons flowing
to the interfâce per one adsorbed benzoate ion at constant potential (electrosorption
valency) and at constant buk benzoate concentration (reciprocal of the Esin-Markov
coefficient). Benzoate is an aromatic anion and its surface coordination changes fiom a
flat (n-bonded) state at a negatively charged surface to a vertical (tilted) state at positive
charge densities. However, benmate adsorption takes place predominantly at the
positively charged surface. The change in the surface orientation takes place at very low
surface coverages ( l e s than 20% of the limiting value). Consequently,the adsorption of
benzoate at the Au(l11) surface are dominated by the vertical tilted CO-ordinationand the
rnajority of the thermodynamic data correspond to this surface orientation. The
orientation change of adsorbed benzoate ions will be characterized by subtractively
norrnalized FTIR spectroscopy described in the next chapter.
OveralL the adsorption of benzoate resembles closely the adsorption of inorganic
anions. At positive charge densities the charge of the benzoate ion is effectively screened
by the charge on the metal and the electrosorption bond has a very low polarity.
AlternativeIy, the low poIarity of the electrosorption bond may be explained by a charge
transfer from the benzoate ion to the metal. The benzoate ion is more buLky than
iwrganic anions such as halides or sulfate. The larger size of benzoate af5ects chiefly the
magnitude of the inner Iayer capacity. The components of the inner layer capacity at
constant charge and at constant amount adsorbed are lower for benmate than for
inorganic anions.
References
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Chapter 5 FTIR Studies of Benzoate Adsorption on the
Au(ll1) Electrode
The objective of this chapter is to complement our previous electrochemical
studies of benzoate adsorption in Chapter 4 [Il. In situ FTIR wïU be employed to
investigate the surface coordination of the adsorbed molecules. AIthough conventional
electrochemical methods
121 provide significant information concerning
the d a c e
concentration of adsorbed motecules and energetics of moIecuIar adsorption, it is s t U
desirable to supplement and confirm these results with techniques that provide insight hto
molecular structural behaviour. In situ surface spectroscopie techniques f3-51can provide
information on the vibrational properties of molecules adsorbed at the metal d a c e ,
which help to identm the molecular orientation and CO-ordinationto the adsorption site.
In fàct, surface enhanced Raman spectroscopy (SERS) 131 and in situ KR spectroscopy [o]
have aiready k e n successfiilly applied in Weaver's Iaboratory to study the CO-ordination
of benmate to a polycrystalline gold e1ectrode. These studies showed clearly that both in
situ IR and SERS provide the required information concerning the dissociative or nondiçsociative character of adsorption and the potential induced change of the orientation of
molecules adsorbed at the electrode surface. In addition, high resolution electron energy
loss spectroscopy (HREELS) and reflection-absorption infiared spectroscopy @AIRS)
have k e n used to investigate benzoate adsorption on copper 17-91 in vacuum These
studies provided usefùl information conceniing vibrational modes of the adsorbate-
4
substrate cornplex and the nature of the interaction of the benzoate ion with a metal
surfàce.
In this chapter, subtractively normalized interfacial Fourier transform S a r e d
spectroscopy (SNLFTIRS) [IO, 111 has been used to explore the orientation of adsorbed
benzoate on the Au(ll1) surface. We wiU demonstrate that no oxidation of benzoate
occurred on the Au(l11) surface over the potentid range investigated. In addition, we will
descn'be how the data f?om Chapter 4 has been used to analyse the reiationship between
the surface concentration of benzoate and the IR band intensity. Finaliy, we will perform a
carefùl analysis of the potential controlled reorientation of the adsorbed benzoate
molecules. The results of this work are by and large, consistent with previous studies of
benzoate adsorption at the Au@oly) electrode by Weaver and CO-workers[3,61. However,
due to the availability of the data concemhg benzoate adsorption derived fkom
chronocoulometric experiments cl] we were able to put our conciusions on a more
quantitative basis. Our work illustrates a synergistic relationship between the quantitative
studies of molecuiar adsorption by electrochemid techniques and the in sita
spectroscopie investigation of their coordination to the electrode d c e . In addition, it
provides a data base for the interpretation of IR spectra of amino z i d s adsorbed at the
Auil 11) surfàce, presented in Chapter 6 (also reference [12]).
The supporthg electrolyte, 0.1 M KCiO4, was prepared from potassium
perchlorate p d e d twice according to the procedure in ref [13]. Potassium benzoate
(99%, Aldrich Chernical Company) was used without m h e r purification A syringe type
I
IR cell with a 60° Ca& prism window was used for the in siru FIlR studies. The angle of
the incident photon beam at the electrode[soIutioninterface was about 70" - 80°, ensuring
maximum enhancement [14]. The detailed experimental procedures have been descnid in
Chapter 3.
5.3. Results and discussion
5.3.1. Electrochemical results
The e1ectrochemical measurements provide idormation concernïng the properties
and characteristics of the eleztrochernical processes of benzoate adsorption. These
measurements are necessary in order to deterrnine suitable reference and sarnple potentials
for the FTIR studies and for the interpretation of the spectroscopie data, In the present
work, the adsorption of benzoate was investigated in the potential range of -750 to 600
mV (SCE) that covers the double layer region of the Au(1 f 1). Fig. 5-1 shows (a)
differential capacity, (b) charge demity and (c) surface excess for the Au(ll1) electrode in
the 0.1 M KC104 supporthg electrolyte alone (dotted line), and in the presence of 9.04 x
1o5 M potassium benmate (solid line). The curves in Fig. 5- 1(a) and @) in the presence of
benzoate merge with the curves of the pure supporthg electrolyte at E < -400 mV. This
indicates that the benzoate ions are totally desorbed fiom the Au(ll1) surface at these
potentiais. The total desorption is also proven by the zero surfàce excess at these
.
-750 mV was chosen as the reference potential (El) in the
potentials in Fig. 5 4 ~ ) Thus,
SNlFTTR measurements, the sample potential (E2)W
~ varied
S
throughout the double layer
region ~ o -700
m mV to 600 mV. In this way, the measured potential dinerence spectra
represent the difference between the absorption s p e c t m of T (Gibbs excess) molecules
that are desorbed ftom the surface and reside in the solution at potential El and the
133
-800
-600
-400
-200
O
200
400
600
E/mV vs SCE
Figure 5-1 Differential capacity (a), charge density (b) and Gibbs excess of benzoate ( c ) at
the Au(ll1) electrode surface: dotted h e , aqueous solution of 0.1 M KC104; solid h e ,
aqueous solution of 0.1 M KC104 + 9.04 x 105M benzoate.
spectrum of the same number of molecules adsorbed at the electrode surface at potential
E2The differential capacity curve for the benzoate solution displays a tall peak at
E-350 mV and a shoulder at E-150 mV- The charge density plot shown in the middle
panel indicates that the tall peak is observed at a positively charged surface while the
shoulder is seen at negative charge densities. Consistent with the shape of the capacity
curve, the adsorption isotherm plotted in the bottom panel is quite aqmmetric with
respect to the potential corresponding to half coverage. These features suggest that
benzoate molecules may undergo a potential dependent reorientation by movuig f?om a
negatively to a positively charged electrode surface. The results of IR investigations will
be used below to discuss fùaher the charge-controlled reorientation of adsorbed
mo lecules.
5.3.2. FTIR Studies. Choice of solvent
The i'tnftared fiequency range of interest is 1300 to 2000 cm-'. Figure 5-2 compares
the t r d s i o n spectra for water and D20.A strong O-H bending deformation of water
occurs between 1600 and 1700 cm? which adversely afEects the investigation of benzoate
bands that are present close to this region. In con-
the corresponding O-D bending
mode of deuterated water is shifted to -1200 cm-'. In addition, D20 gives a very flat
background in the Gequency region of interest. For these reasons, 4 0 was chosen as a
solvent in the present studies. While working in D20one has to be aware that isotopic
exchange of hydrogen by deuterium can occur fiom the interaction of D20 with
compounds containing labile hydrogen, such as, for example, the hydrogen bonded to
nitrogen in amines and amides [151. Under normal conditions, hydrogen atoms attached to
-
Figure 5-2 Transmission IR spectra of DzO (dotted line) and &O (solid line).
the benzene ring such as benzoate are not readily exchanged- However, it is not known
how the solvent affects the adsorbed species on the eiectrode surface, and therefore it is
necessary to ensure that no deuteration or m o ~ c a t i o noccurs by the solvent. RATRS and
HREELS studies of deuterated benzoate adsorbed at a copper surface in vacuum [16]
showed ody a very s m d change of the fiequency of the symmetric stretch of the
carboxylate group, in comparison with the non-deuterated species. When the modes are
predorninately C-H or C-D vibrations, the ratio of the reduced masses
&-&-H)
produces a relative reduction in fi-equency by 1.363 compared to the non-deuterated
species. Similarly, for a mode involving predominately the carbon atom, the approximate
fàctor will be 1-038. This prediction is consistent with the experrimental data where
m/w
equal to 1.341 and 1.022 for vc-Hand vcc respectively were observed [16].
Figure 5-3 compares the transmission spectrum of potassium benzoate in D20to
selected SNlFTIR spectra of benzoate on Au(ll1) in Hz0 and D2O. Three bands could be
seen in the spectral range of interest. Foilowing reE [17] they can be assigned to COzsymmetric stretch (1390 cm-'), CO; asymmetric stretch (1547 cm-') and C-C stretch
(1595 cm-').
The assigrment of each band is given in Table 5- 1. Very small shifks in the
band position are observed on the IR spectra recorded in water and in D20. If an isotopic
substitution occurred with the benzoate ion in D20 solution, the band correspondhg to
the C-C stretch would shift to lower fiequencies by about 35 cm-'. n i e observed s W is
much srnailer (on the order of 1 cm-'), indicating that an exchange between hydrogen on
the aromatic ring and deuterium fiom the DzO solvent does not occur. The SNIFTIR
spectra measured in D20 showed a better signal-to-noise ratio, particularly in the region of
the C-C ring stretching mode (Fig. 5-3). The SNIFTIR spectra recorded using s-photon
Transmission
s-pol. in D,O
Figure 5-3 Cornparison of the transmission FTIR spectrum of benzoate in D20and the
SNIFTIR spectra of benzoate in 4 0 and in H20. Transmission s p e c t w was recorded
for a Dfl solution saturated with the benzoate. SNIFTIR spectra were recorded in O. 1 M
KC104 + 9.04 x 1O" M benzoate solution using El=-750mV and Eî
= 400 mV
(SCE).
Table 5- 1 Vibrational Frequencies and Assignments for Benzoate
Vibrational
Transmission /cm-' vibrational assignment
Iiteilture /cm‘'
il71
mode
ai
1390
CO*' sym stretch
E390 [6]
bi
1547
CO; asym stretch
1545 [6]
bi
1595
C-C ring stretch
1594 [17]
poIarisation are monopolar. En contrast, the band correspondhg to the symmetnc CO2stretch has bipolar features under the p-polarized beam.
This behaviour will be used
bdow to study the orientation of benzoate on the electrode surface.
5.3.3. Relationship between the IR intensity and the surface excess
In our experirnents, the relative change of the electrode reflectivity rneasures the
difference between the absorption spectm of T molecules that are desorbed eoom the
surface and reside in the solution at potential Eiand the spectrum of the same nurnber of
molecules adsorbed at the electrode surface at potential E2:
where
E
is the molar absorption coefficient and
r is the
surface concentration of the
adsorbed species. For quantitative analysis, it is convenient to use the iniregrated band
intensities, In that case:
(AZUR)dv = 2.3 T ~ ~ E- E(E2)ld
I ) v
where v is the fiequency in cm-'.
( 5.2
For linearIy poIarized light and a molecule adsorbed or in fiont of a reflecting
metal surface, the integrated absorption coefficient is proportional to the square of the dot
product of the transition dipole p and the electric field of the photon E [l8,l9] :
b d v a 1 p~l 2 a cos2 8 l P l 2 <>?!i
( 5-3 )
where 6 is the angle between the direction of the electric field of the photon and the
transition dipole in the rnolecule, 1 p 1 is the absolute value of the transition dipole moment
and <
J?>
is the mean square electnc field of the photon. At the potential of total
desorption El, benzoate molecules are randornly oriented inside the thin layer of the
electrolyte. In this case, the angle û has to k averaged over all possible orientations and
the result is -os2@ = 113 [18]. Consequently, at El the integrated absorption coefficient
is given by:
[LQ~I)
d v a (113) 1
1
< I?(SO)
>
( 5-4 )
for both s and p-polarized radiation. Using s-polarized light, the electnc field of the
photon at the m d k e is nearly equal to zero [19,20]and hence 4E2) is close to zero as
welL Ushg equations 5.2 and 5.4, the relationship btween the integrated IR intensity and
the Gibbs excess (0can be descn'bed by:
Fig. 5-4 shows a series of SNIFTIR spectra for benzoate under s- and ppolarization. The IR intensity at l 547 c ~under
l s-polarization has been integrated at each
sampling potential. The integrated IR intensities (med circles) are compared with the
Gibbs excess (open diamonds and solid line) determuied by chronocoulometry
[Il
in Fig.
5-5. The IR intensity agrees very weli with the Gibbs excess at all potentials, consistent
Figure 5-4 SNPTIR spectra of CO; asymmetric stretch and C-C stretch for benzoate
adsorbed on the Au(1 Il) electrode fiom 0.1 M KC1O4 + 9.04 x 10" M benzoate solution.
The spectra were acquired using (a) s-polarized and (b) p-polarized infkared light for a
series of sample potentials (EZ)whose values are indicated in the figure. The reference
potential EL=-750mV (SCE).
-800 -600 -400 -200
O
200 400
600
800
E/mV vs SCE
Figure 5-5 Cornparison of the IR band intensity at 1547 cm-' (nlled circles) with the
surface concentration of benzoate (soIid line and triangles) obtained by thermodparnic
analysis of chronocoulometric data (Chapter 4).
with equation 5.5. The positive peak measured under s-polarization represents loss of
benzoate f?om solution due to its adsorption at the electrode surface at potential Ez. We
rnay therefore assume that benzoate is not oxidied over the investigated potentials. If
benzoate had been oxidized, the most probable oxidation process at high potentials would
involve decarboxylation with the release of CO2, which gives an asymmetrk stretchuig
mode at 2340 cm-' [4]. In a SNIFTIR spectrum, the appearance of a negative band at
2340 cm-'would indicate that CO2 is formed at potential E2. Since DzO has an adsorption
band in this range of fkequencies (Fig. 5-2), it was necessary to use Hz0 as solvent to look
for the appearance of the COz band, Fig. 5-6 shows a SNIFTTR spectnim of benzoate in
Hz0 determined for Et 4 0 0 mV. Only peaks at 1390, 1547 and 1597 cm-' and no CO2
band can be seen in the spectnim Apparently benzoate is not oxidized at the Au(ll1)
when the electrode potential is equal to or lower than 600 mV.
This conclusion is
consistent with the previous work on benzoate adsorption at a polycrystalline gold
electrode
m.
5.3.4. The orîentation of the benzoate molecule on the electrode surface
The isolated benzoate rnolecuie is planar and has
spectroscopy studies in ref [8] indicate that
C2Vsymmeetry.
C2v symmetry
The wirational
is retained for the benmate
molecule adsorbed on a copper surface. It is therefore k e l y that benzoate adsorbed on
gold has Ch symmeetry as welL In that case, the a1 and br bands correspond to in-plane
deformations. The transition dipole for a, modes is parallei to the C2v axis, while the
transition dipoie for bI modes is oriented in the direction normal to the C2, axis. For ppolarized Light the electnc field of the photon at the surface is enhanced by reflection and
is nearly normal with respect to the electrode surface [19,20].In this case, the shape ofthe
1390 cm-'
I 595 cm"
Figure 5-6 The SMFTIR spectrum of 0.1 M KC1O4 + 9.04 x 1 o5 M benzoate solution ui
Hz0 for the Au(ll1) electrode recorded in the fiequency range extended to 2400 cm-';
reference potential -750 mV and sarnple potential600 mV (SCE).
spectrum for adsorbed molecules can be predicted fkom the d a c e selection rules. For a
molecule adsorbed at a metal, the swdace selection d e s state that the viiration will be
allowed if a change in the dipole moment has a non-zero component in the direction
normal to the surface [19]. When the benzoate molecule assumes a flat, ir-bonded coordination, the change of the dipole moment in the direction normal to the surface has a
zero component for both a1 and bi modes. The adsorbed molecule is optically inactive, and
the
SNLFTIR spectrum consists of oniy positive bands corresponding to the IR absorption
by molecules that were desorbed fiom the electrode surface into the thin layer cavity at the
potential E l . In this case, the spectra recorded using s- and p-po1-d
radiation should be
quite similar. In contrast, if the benzoate mdecule is CO-ordinatedto the surface through
the carboxylate group, the plane of the molecule is either vertical or tilted with respect to
the surface. For this orientation, the transition dipole of ai bands bas a strong component
normal to the surfàce, while the transition dipole of bi bands remains pardel or nearly
parailel to the surfkx. For the adsorbed species, bI bands remain opticdy inactive. In
contrast, al bands are now optically active. In the SNIFTIR speclnim, bi bands continue
to be positive (they will correspond to the absorption of IR radiation by molecules
desorbed at El). In contrast, al bands are iikely to be bipolar. The negative lobe of this
band will be due to IR absorption by moiecules adsorbed at the electrode surface at the
potential Ez while the positive lobe will be due to IR absorption by molecules desorbed
into the thin layer caviîy at El.
For p-poiarized light, bi bands shown in Fig. 5 4 b (C-C stretch and asymmetric COîbands) are positive in the whole range of potentials Ez investigated, A brief inspection of
the spectra in Figs. 5-4a and b shows that the shapes of br bands recorded
Figure 5-7 SN?FTIR spectra of C o i symmetnc stretchïng mode for benzoate adsorbed
on the Au(l11) electrode fiom 0.1 M KCi04 + 9.04 x 10;' M bernate solution using: (a)
s- and @) p-polarized h h r e d beam and a variable sample potential
figure.
(Ez)indicated in the
using s- and p-poiarized radiation are nearly identical This behaviour indiates that the
transition dipole of the bi modes is parallel or nearly parallel to the electrode surfàce in the
whole range of the electrode potentials investigated. Figure 5-7 shows the SNIFTIR
spectra of the CO; symmetric stretch (alsymmetry) for a series of variable potentials (E2).
For s-polarized Light, the band is positive at all potentials, since it corresponds to IR
absorption by the solution species. The amplitude of this band hcreases in proportion to
the surface concentration of adsorbed benzoate. In contrast, the shape of the band
recorded using p-polarized light changes with potenta At potentials lower than 200 mV
the band is positive. However, at more positive potentials, it becomes bipolar. Apparently,
the amplitude of the negative lobe of this bipolar band increases with potential. In
addition, the position of the negative lobe shifts towards higher fiequencies with potential.
The fiequency of the peak of the bipolar band is plotted against the electrode po tential in
Fig. 5-8. The relation is linear and its dope is equal to 47 cm-'v-'. These results strongly
suggest that the orientation of the adsorbed rnolecule changes with potential For this
band, the component of the transition dipole in the direction normal to the surfàce is small
at low potentials, but it increases progressive@ with potential, Such behaviour suggests
that the adsorbed benzoate mokcule undergoes a potential controlled reorientation f?om a
5 t (n-bonded)
to a vertical or tilted state. The strong dependence of the keqiiency of this
band on the electrode potential may be taken as an argument that the vertical (tilted)
orientation corresponds to the CO-ordinationof the benmate ion through the carboxylic
group to the electrode surface. For moIecules adsorbed at the metallsolution interface, the
potential dependence of the IR band i2equency may be caused by either back donatioii or
the Stark effect. When this dependence is strong, the CO; group m u t be in the proximity
siope = 47 (cm-'
-1O0
O
100
200
300
400
v")
500
600
700
E/mV vs SCE
Figure 5-8 Dependence of the fiequency of the negative lobe of the bipolar band of the
symmetric CO; stretch on the electrode potentid.
of the metal surfàce to experience either a sufncient couphg with the electronic states in
the metai or a sufEciently strong electrostatic field.
We have recently s h o w [21], that fbrther information conceming the orientation
of adsorbed molecules can be Obtained from the analysis of the integrated band intensities.
For that purpose, it is convenient to calculate the ratio of the integrated intensities
rneasured for p- and s-polarkzed radiation. Ushg Equations 5.1 to 5.3, the ratio of the
integrated intensities may be expressed as:
where (EZ,&)>
is the t h e and z averaged mean square field within the thin layer cell, and
is the t h e averaged mean square field at the electrode surfrice and z varies
between O and d, where d is the thickness of the thin layer. The angle 0 in equation 5.6, is
the angle between the direction of the transition dipole for the adsorbed molecuie and the
direction of the field of the p-polarized photon on the surface. The ppolarized photon on
the surfàce bas a predorninant component in the z-direction (direction normal to the
surface). Consequently, to a good approximation the angle 0 is a measure of the tilt angle
for adsorbed molecules. The ratio of integrated band intensities for the a, band at 1390
cm-' and the sum of br bands at 1547 and 1595 cm-' is plotted against the electrode
potential in Fig. 5-9. For bI bands, the ratio of integrated intensities changes very Little
with potential and amounts to about 1.18. This indicates that the angle between the
direction of the transition dipole of br modes and the direction of electric field of the
200
400
E/mV vs SCE
Figure 5-9 Ratio of the integrated intensities for the SNIFTIRS bands acquired using p-
and s-polarized radiation; circles, bi band at 1390 cm"; diamonds, al band at 1547 cm-'.
photon is essentially unchanged. The hi& positive value of the ratio indicates that the
component of the transition dipole in the direction of the field of the p-polarized photon is
small and that the dipole is essentidy parallei to the electrode surface.
In contrast, the ratio of integrated intensities for the al band changes with the
electrode potential. For E < 100 mV the ratio is approxirnately equd to 0.7 while for E >
300 mV it attains a constant value of -0-2-
In the potential range fiom 100 to 300 mV
the ratio varies gradually between the two lùniting values and changes sign Incidentaliy,
this is the range between potentials of the two peaks on the differential capacity cuve
shown in Figure 5-la Earlier, we assîgned these peaks to two orientations of adsorbed
benzoate ion The spectroscopie data support this interpretation In order to calculate the
exact values of the tilt angle, one has to know the field strength of the photon withh the
thin layer cavity- Mode1 calculations descriid in ref [21] indicate that the field strength
changes with the thin layer thickness in an oscillatory m e r . One has to know the exact
value of the t h h layer thickness in order to calculate the field strength. In the present case
the exact value of the thin layer thickness is unknown. Consequently, we cannot calculate
the values of the tilt angles.
Nevertheless, we can
~~ extract u s e l l information conceming the orientation of
benzoate on the surface fiom the data shown in Fig. 5-9. For E < 100 mV and the ai
band, the ratio of the întegrated întensities is positive. However, it is srnalier than the ratio
measured for the bi band. This fact indicates that even at low potentials the angle between
the surface normal and the transition dipole of the ai band is less than 90" and hence the
adsorbed benzoate ion is somewhat tilted with respect to the electrode surface. At higher
potentials, the ratio of integrated intensîties of the al band is negative. This indicates that
the component of the transition dipole in the direction normal to the d c e is signincant
in this region. However, its absolute value is small. In a recent study of pyridine
adsorption at the Au(ll1) electrode the ratio of integrated intensities of the ai band was
equal to -1
[21].
.O for a tilted (vertical) orientation of the pyridine molecule at the surface
Pyridine and benzoate molecules have the same Cz, symmetry. Hence, for the same
tilt angle, the same ratio of integrated intensities should be observed in these two cases.
Indeed, for bands of br symmetry the ratio of integrated intensities has nearly the same
value for benzoate and pyridine. The fact that for al bands the ratio of integrated
intensities is equal to -0.2 for benzoate and to -1.0 for pyridine, indicates that the angle
between the Cz, axis of the molecule and the surface normai is larger for benzoate than for
pyridine. Consequently, the potential controlled reorientation of adsorbed benzoate
involves a progressive change of the tilt angle with respect to the surface n o d .
However, the adsorbed benzoate is neither perfectly flat at the lower end of potentials nor
does it assume a vertical orientation at the most positive potentials,
5.4. Conclusions
Subtractively nomialized interfacial Fourier transform ïdhred spectroscopy and
electrochemical measurements have been applied to the study of benzoate adsorption at
the Au(ll1) electrode. In order to avoid interference f?om absorption bands of H20,
the
solvent chosen for the spectroscopic study was D20.The experiments showed that the
isotopic substitution of hydrogen atorns on the benzene ring did not occur. Benmate
undergoes a progressive orientation change on the electrode surface fiom a nearly fiat
configuration Vig. 5-10a) at potentials below 200 mV to a tilted orientation (Fig. 5- 1 Ob)
Figure 5-10 Two extreme surface coordinations of benzoate on a Au(l11) electrode.
(a) Flat orientation; (b) Vertical orientation.
at more positive potentials. The transition between these two surface orientations takes
place at potentiak close to the potential of zero charge in the benzoate solution In the
tilted orientation the benzoate ion is coordiited to the eIectrode surface through the
carboxylate group. The adsorption of benzoate has a non-dissociative character at
potentials 5 600 mV. The integrated IR band intensities are proportional to the d a c e
concentration Obtained £kom the electrochemicd measurements- The Overali agreement
between IR and eIectrochemical measurements is excellent-
Refeaences
1. H. Li, S.G. Roscoe and J. Lipkowski, J, Sol. Chem accepted.
2. J. Lipkowski and L. Stolberg in J. Lipkowski and P. Ross (Eds.), Adsorption of
Molecuies at Metal Electrodes, VCH, New York, (1992) 171.
3.
P. Gao and M.J. Weaver, J. Phys. Chem 89 (1985) 5040.
4. D.S. Corrigan, E.K. Krauskopc L.M. Rice, A. Wïeckowski, M.J. Weaver,
J. Phys. Chem 92 (1988) 1596.
5. A. Chen, J. Richer, S.G. Roscoe and J. Lipkowski, Langmuir, 13 (1997) 4737.
6 . D.S. Corrigan and M. J. Weaver, Langmuir 4 (1988) 599.
7. B.G. Frederick, Q. Chen, F.M. Leibsle, S.S. Dhesi and N.V. Richardson,
Surf. Sci 394 (1997) 26.
8. B.G. Fredenck, P.D.A. Pudney and N.V. Richardson, J. Electron Spectroscopyand
Related Phenornena, 64/65 (1993) 115.
9. T.S. Jones, M.R Ashton and N.V. Richardson, I. Chem Phys. 90 (1989) 7564.
10. S. Pons, T. Davidson and A. Bewick. J. Electruanal. Chem. 160 (1984) 63.
11. S. Pons, J. Electroanal. Chem 150 (1983) 495.
12. H. Li, S.G. Roscoe and J. Lipkowskï, J. Electroanal. Chem. submitted.
13. J- Richer and J, Lipkowski, J. Electrochem Soc. 133 (1986) 121-
14. P.W. Faguy, W.R. Fawcett, Appl. Spectroscopy, 44 (1990) 1309.
15. E. Blout, in O.V.Whilelock, F.N. Furness and H.P. Cameron (Eds), Annals of the
New York Academy of Sciences, New York Publishg Academy, 69(Art. 1) (1957 84.
16. P.D.A. Pudney, B.G. Frederick and N.V. Richardson. SurE Sci 307 (1994) 46.
17. J.H.S. Green, W. KKynaston and AS. Lindsey, Spectrochim.. Acta, 17 (1961) 486.
18. W. Liptay, Angew. Chem 8 l(l969) 195.
19. M. Moskovits, J. C h e n Phys. 77 (1982) 4408.
20. KG. Greenler, J. Chem. Phys. 44 (1966) 3 2021. M. Hoon-Khoshla, W-R Fawcett, A Chen, J. Lipkowski and B. Pettinger,
Electrochim-Acta, 45 (2000) 6 11.
Chapter 6 Electrochemical and FTIR Shidies of LPhenylalanine Adsorption at the Au(ll1) Electrode
6.1. Introduction
The interest in the interfacial behavior of proteins at solid surfaces originates fiom
the need to better understand the mechanisms of processes associated with their use in
advanced technical applications and industriai problems. In order to understand the
behavior of proteins, it is important first to understand the adsorption behavior of amino
acids- L-PhenyIalanine was chosen for this study as it contains the hydrophobie aromatic
group and therefore follows a series of small organic compounds (i.e., cyanopyridine Cl],
benzonitrile [2], and benzoate [3,4], pyridïne [553 and 2,2' bipyridine
[q)that
we have
been studying using electrochemical and in situ FTIR to examine their orientation and
coordination at the Au(l11) electrode swhce.
The adsorption of phenyiahhe (Phe) at metal surfàces has been investigated by a
number of techniques- DifEerential capcitance and radioactive indicaiors were used to
investigate the adsorption of phenylalanine and tyrosine at bismuth and mercury electrodes
[7,8]. Other electrochemicai studies of amino acids at mercury electrodes were made with
glycine [9], glycyl-glycine [10],and methionine [Z Il. Studies have also k e n carried out on
glycine [12,13], a- and p-alanine 113,141, a-,P- and y-aminobutyic acid [13,15], and
tyrosine and tryptophan [16] at po lycrystalline Pt. A mechanism involving adsorption
through the carboxylate group foliowed by decarboxylation was proposed for the
oxidation of the amho acids.
In situ FTZR spectroscopy is a technique that has been used very successflllly in
the study of electrified interfàces [17-201. This technique d o w s the identification of
intermediates and products of an electrode reaction, and it has been widely used to study
electrocatalytic oxidation mechanisms of smali organic molecules at noble metai electrodes
[21-231. Because of the surface selection rules, the technique provides information on the
orientation and coordination of molecules adsorbed at metal surfaces. Applications of this
technique have been made to the adsorption of CO to platinum group metals E3.3-261, as
weii as a number of organic molecules on polycrystalluie platinurn [27, 283, mercury
[29,30] and gold [3 1, 32,331. The in siCu FTIR studies were concemed either with the
identification of species generated at the electrode surface or the determination of the
characteristic of their s d a c e coordination in most of these applications. For example, in
situ FTIR and electrochemicai measurements on the adsorption of glycinate on Pt(ll1)
showed the molecule to be two-fold coordinated to the Pt through the carboxylate group
and with COz formation at the higher potentiak 1341. Less work has been done to develop
FîIR speçtroscopy as a tool for quantitative analysis of adsorbed species tbat could
provide information about the composition of the interlEicid region Only a few papers
report quantitative studies [20,35,361.
The objectives of our present work are the-fold; (3 ta employ the surface
selection rules of IR spectroscopy to investigate the potential induced reorientation of the
adsorbed molecules, (2) to use IR spectroscopy to determine the potentïal range within
which the molecules are stable at the electrode surface and to detennine the potential for
the onset of Phe oxidation, and (iiii to correlate the h c e concentrations of Phe
molecules detemhed fiom electrochemicai studies with the intensities of selected
IR
bands for adsorbed molecules.
6.2. Experimental
L-Phenylalanine is a zwitterionic molecule at the pH of its isoelectric point which
corresponds to 5.75 [37]. AU experiments descnid in this work were performed in a
neutrd unbuffered electrolyte at pH values close to the isoelectric point. The experimental
procedures for both electrochemical and SNIFTIRS studies have k e n descnid in
Chapter 3. The transmission spectra were recorded using a thin layer of a solution of
phenylalanine in D20between two flat ZnSe windows or solid phenylalanine immobilized
in a KBr pellet.
6.3. Results and Discussion
6.3.1. Eiectrochemical results
Cyclic voltammetry (CV), differential capacity O C ) and chronocoulometric
measurements provide information conceming the properties and characteristics of the
electrochemical processes of phenylalanine adsorption ai the Au(ll1) d c e . This is
particdarly important in order to determine suitable reference and sample poteztials for
the FTIR studies and for the interpretation of the spectroscopie data. Figure 6- la shows
the cyclic voltammetry curves recorded in the double layer region of Au(ll1) in 0.1 M
KClQ and in the presence o f 1.06 x 1 0 - ~M Phe. Three pairs of quasi-reversble peaks
appear between -300 to 600 rnV resdting f?om the adsorption of Phe. The peaks
recorded in the positive voltage scan are quite broad relative to those associated with the
I
I
I
-600
-400
I
I
I
-200
O
200
E / mV vs SCE
I
1
400
600
Figure 6-1 For a Au(ll1) electrode in 0.1 M KC104 (dashed lines) and 0.1 M KCIO4 +
1 .O6 x 1
M Phe solutions (solid ihe), a) cyclic voltamrnetry curves recorded using the
sweep rate 20 mV
Ç'. b)
dBerential capacity recorded using a 5 mV
rnodulated at 25 Hz and a 5 mV 5' sweep rate.
(mis)
sine wave
adsorption of inorganic ions such as Cl- [38]. Figure 6-1 b shows differential capacity
curves recorded for the supporthg electrolyte alone and in the presence of 1.06 x 10"
M
Phe. Consistent with the CV curves, the capacities display three peaks suggesting that Phe
adsorption is a three-state process The peaks on the differentkd capacity curves are
observed at potentials;
- -100 , - 300 and - 500 mV.
The curves recorded in the
presence of Phe merge with that of the pure electrolyte, at 4 0 0 mV, indicating total
desorption of Phe at this potential. Thus,-600 mV was chosen as the initial potential for
capacitive chronocoulometry and iII situ
FïïR measurements. In addition, the CV and
differential capacity curves were helpfül in establishing the upper potentials lùnit for the in
situ FTIR studies.
Capacitive chronocoulometry was used to measure the ciifference between the
charge density at the potential of adsorption ( Ei ) and that of total desorption ( Ef ) [39411:
AOM(E~)
= CM&)
-
(6-1)
Solutions of phenyhlanhe were prepared over the concentration range of IO-' to 10" M in
0.1 M KCl04. When Phe concentration was lower than 104 M, solutions were stirred in
order to e h c e mass transport rate and to ensure that the state of adsorption equili'brium
was establisned. The potential of zero charge @zc) was determined fiom the position of
the difEUse layer minimum of a differential capacity c w e measured separately with a dilute
solution of the supporting electrolyte. It was equal to 280 mV versus SCE. The absolute
charge density at El was calculated fiom
AGM@ZC)= CM(&) - ~ ~ ( p z=cSM(&)
)
(6-2)
Since there is no adsorption at Efithe charge, CM@),
molecdes in solution. T'usAQ&)
is independent of the presence of the
and -(El) were calculated using the above equations
for the complete set of potential range and solutions investigated. The charge density (w)
versus potential plots, determined for the pure electrolyte and selected concentrations of
Phe are chosen in Figure 6-2. PLU charge density curves merge at -500 mV, indicating that
Phe molecdes are already desorbed at this potential. The charge density curves also
dispiay multiple-step characteristics of a rnulti-state adsorption process of Phe. The
potentiais of these steps correspond to the position of peaks observed in both the CV
(Figure 6- 1a) and DC (Figure 6- 1b) cuves. The pzc shifts to more negative values in the
presence of Phe, This feature indicates that the adsorbed molecules assume an orientation
in which the negative pole of the permanent dipole moment is directed towards the
Au(ll1) sufface and the positive pole towards the solution
6.3.2. Gibbs excess and Gibbs energy of adsorption
The i i h pressures p)of adsorbed phenyhbine were calculated using
where 8 and 8 = O represent the presence and absence of the amino acid in the bdk
solution, respectively. The relative Gibbs excesses were then calculated fiom
Figure 6-3 shows the Gibbs excess (T) versus potential curves for various concentrations
of Phe in O. 1 M KCIOs solution. The Gibbs excess increases with potential until it reaches
a maximum at 300 mV followed by a decrease at higher potentials. The maximum Gibbs
6.52x lo4 M
2.27 x IO-^ M
6.36 x lo9 M
0.1 M KCIO,
-3v
1
I
I
1
I
I
I
-600
-400
-200
O
200
400
600
E / m V v s SCE
Figure 6-2 Charge density curves determined fiom chronocoulometry in 0.1 M
KCQ(dotted h e ) and, for clanty, a selected set of concentrations of Phe are shown
mging fiom 1.83 x IO-' M to 1.06 x 10" M.
-600
-400
-200
O
200
400
E / mV vs SCE
Figure 6-3 Gibbs excess vs. potentid plots for Phe on Au(ll1) at 1.26 x 1 0 M
~ ~(circle);
2.27 x 105 M (square); 3.52 x IO-' M (up-triangle); 6.36 x 10'~ M (down-triangle); 1.06 x
1 0" M (diamond). Inset, for 1 -06 x10" M Phe solution, cornparison of the Gibbs excess
and the integrated intensities of the 1 574 cm-' (open triangle) and 1410 cm-' (open circle)
SN?FTRS bands in the spectrum of Phe in D20 recorded using s-polarized photons.
K
excess of Phe is about 2 x 10-'O mol cm-2. This number is somewhat lower
than the
maxirnumexcess of Phe determined at the Bi(lI1) and Hg electrodes [7,8] .Qualitatively,
adsorption of Phe on Au(ll1) and Bi(l1l) as well as on Hg electrodes displays many
common features. Somewhat lower values of T Observed for the Au(1ll) electrode may
be explained by the eight times lower concentration of Phe used in our work.
The Gibbs energy of adsorption AGads is uually detennined fiom a fit of the
experimental data to an equation of a particdm adsorption isothem The adsorption of
Phe has apparently a multiple state character and no simple isotherm can be used to
describe the change of its surface concentration. However, in the limit of zero coverage,
all isothenns simple to the Henry isotherm [39]:
where .x: is the mrface pressure, T,, is the lunithg Surface concentration, c is the bulk Phe
concentration, and
p is the equili'brium constant which is related to the G%bs energy of
adsorption through the equation AGds
=
-RT Inp. The e q u i l i i constant can be
evaluated fiom the initial slopes of the d a c e pressure versus the bulk concentration
plots. The value of ,
T
equai to 2 x 10-'O mol
was taken to calcuIate
B. The Giibs
energies of adsorption determined in this way are plotted againsi the potential in Figure 64. The standard state corresponds to unit mole fiaction of the organic species in the buk
of the solution and to monolayer coverage by the noninteracting adsorbates. One can also
use the value of RTln( c e - ~ s/55.5) as an additional measure of the energetics of Phe
-300 -200 -100
O
100 200
E /mV vs SCE
300
400
Figure 6-4 For Phe on Au(1 1 1), Gibbs energy o f adsorption and RTln ( ~ ~ = ~ . s / 5 )5 . 5
plotted versus the electrode potentid.
4
adsorption When adsorption is descnid by the Langmuir isothenn, then RTln(cs0.s
l55.5) = AGnds. Therefore, the difference between RTln(cer5 /55.5) and AGad,detennined
f?om the Henry's Law can be used as a rneaçuce of the deviation of the behavior of a given
system fiom that descriid by the Langrnui isotherm. The values of RTln(cea5 i55.5) are
plotted dong the AGadsdetennined fiom the Henry's Law in Figure 6-4. The dflerences
between the hvo sets of data are quite srnall (less than 2 kJ mo 1-') and this feature indicates
that the sudàce properties of Phe do not deviate too much fkorn those of the system
descriid by the Langmuir isothexm. The Gibbs energy values for Phe have comparable
magnitude to the values of Gibbs energies of other weakly chemisorbed aromatic
molecules [3 91.
6.3.3. FTIR studies
(i) Choice of solvent
For îhe reasons descnkd in the section 5.3 -2 in Chapter 5,
a0 was chosen as a
solvent in the present shidies. However, isotopic exchange of hydrogen by deuteriurn can
occur fiom the interaction of DzO wÏth compounds containing labile hydrogen, such as,
for example, the hydrogen bonded to nitrogen in amines and amides [42]. Under this
condition, the deformation bands of ND3' move to the fiequencies below 1200 cm-'.
(ri) General properties of IR spectra
Amino acids are amphoteric. At low pH values, both functiond groups are M y
protonated so that the amino acid molecule assumes the cationic fonq i-e.,
RCH(NH3')COOK, cationic form (c):
~t high pH values, al the acidic protons have been
I
removed so that Ï t assumes the anionic form (A); i-e., RCH(NH2)COO-. At pH near the
isoelectric point (PI), the amino acid exkts as the nvitterionic forrn (Z), RCH(N&')COO-.
The pKi and pK2 for Phe is 2.20 and 9.31 respectively 1371 and the p1 is 5.75. The
Lwitterionic fom will therefore be present in varying amounts in the pH range of
- 4 to
1o.
To facilitate interpretation of the electroreflectance spectra, we wiU first discus the
transmission spectra of Phe shown in Figure 6-5. Curve a shows s p e c t m of pure solid
Phe pressed in a KBr pellet. The molecule of Phe in the KBr pellet is in the zwîttenonic
form - Z [43]. Table 6.1 gives the assignment of major bands, based on the band
assignment for aknine C43-451. The band at 1625 cm-' is assigned to -NHSfasymmetrk
deformation (Amino Acid 1 band). The band at 1562 cm-' corresponds to asymmetric COO' stretch. Its broad envelope contauis also the symmetric -NI&' deformation (Amino
Acid II band) . Bands at 1495 and 1458 cm-' correspond to CH2 scissoring deformations
and the band at 1410 cm-' is the symmetric 400stretch band. The curve b in Figure 6-5
shows the spectrum of Phe in a neutd solution of 9 0 (pH = 7).At this pH the molecule
is predomuiant1y in the zrvitterionic form There are signiscant differences between the
spectra of form Z in the KBr pellet and in the DzO solution. Spectnim b contains only two
strong bands at 1616 and 1410 cm-' that rnay be assigned to the asymmetric and
symmetric 400-stretch, respectively [43,44]. The weak bands are lost in a relatively
high noise observed for this spectrum Curve c in Figure 6-5 shows spectnim of Phe in
alkaline (pH = I 1) solution of D20,where the molecule is in the anionic forrn - A [43].
This spectnim is &O dominated by the two -COO- stretch bands.
2000
1900
1800
1700
1600
1500
1400
1300
Wavenumber / cm-'
Figure 6-5 Transmission spectra of Phe and related compounds; (a) spectnim of Phe in a
KBr pellet, (b) spectnun o f Phe in neutrd solution of DzO, pH=7, (c) spectrum of Phe in
alkaline solution of D20,pH-Il ,(d) spectnim of phenylacetate in D20.
8
For the anionic form, the asymmetric 400- stretch appears at 1574 cm-' and is
red shifted with respect to the position of the same band in the zwitterionic form
[Ml.
Characteristically, the 1625 cm-' band seen in spectmm a is absent in spectra b and c. The
absence of this band suggests that the use of DzO as a solvent results in deuteration of the
amino group due to the presence of labile hydrogens attached to the nitrogen. Hence, the
form Z assumes a structure RCH(ND1')COO- and form A is RCHO\TD2)COO'. Because of
the change in the reduced mass fiom N-H to N-D, the deformational modes shift to lower
fkequencies by an amount v K b K H= 0.72 [43]. This corresponds to fiequencies below
1200 cm-', outside the spectral range examined in this work. Consequently, in D 2 0 the IR
spectra of forms Z and A consist chiefly of two strong bands correspondhg to syrnmetric
and aqmmetric 400- stretches and two small CH2 scissoring deformations. To ver*
this assignment, transmission s p e c t m for phenylacetate in D2O (spectnim d) is also
plotted in Figure 6-5. Indeed, the IR spectrum of the A f o m of Phe resembles closely the
specnim of phenylacetate.
Figure 6-6 compares the transmission spectra of Phe in DzO recorded in the an<aüne
solution (pH
=
11) - spectnim a and in the acidic solution @H = 2)
ciifferences between the two spectra are very si@cant.
- spectnim b.
The
In scidic solution, the IR
spectrum of Phe consists of two broad bands at 1732 and 1458 cm-' that may be assigned
to the C=O stretch and the symmetric -COOD stretch of the deuterated carbxylic group.
Apparently, the use of 4 0 as the solvent in acidic electrolyte results in deuteration of
-NH3'
group as well. n i e structure of the phenylalanine molecule becomes
RCH(Nû33COO~(form C) [45]. The net result of these changes is that the IR spectnim
Figure 6-6 Transmission spectra of; a) Phe in alkaline solution of D20,
pH=l 1. b) Phe in
acidic solution ofDzO (pH = 2), c) phenylacetic acid in DzO.
Table 6.1. Vibrational Frequencies and Assignments for Transmission and SNIFTIR
Measurements
Vibrational frequency in cm-'
Zwitter- Zwitter Anionic Cationic
ionic
form in form in
ionic
acidic
form in form in alkaline
neutral
KBr
D20
D20
1
1
1341
1410
1458
1495
1562
1610
1625
-
1
1
D20
14i0
1616
1
1
1341
1412
1458
1495
1
1
-1458
-
1 CH bend
1 COO-symmetric stretch or COOD
CH2 sciçsoring deformation
NH3+symmetrïc deformation viiratiow
COO- asymmeb5c stretch
1574
-
Assignment of transmission spectra
I
1732
NH3' asymmetric deformation vibrations
C=O asymrnetric stretch of COOD group
2 - zwittenonic form, RCH(NK3')C00- (or RCH@lD3')COO' in D20)
C - cationic form, RCH(NH3')COOH (or RCH(ND3')COOD in 9 0 )
A - anionic form, RCH(NH2)COO- (or RCH(ND2)COO- in D20)
of phenylalanine m acidic D20 resembles spectnim of the phenylacetic acid shown as
spectnim c in Figure 6-6. The SNIFTIRS spectra described below were recorded m an
unbuffered solution of O. 1 M KCI04. In this solution, reduction of trace owgen or water
at negative potentiais may cause an ïncrease in pH in the thin layer cell.
Therefore, we
have examined changes in the character of the transmission IR spectra of Phe in the pH
range fiom 7 to 1 1. These spectra are plotted in Figure 6-7.The ratio of the zwitterïonic
to anionic form changes sigdicantly in this region of pH. In the IR spectrum of Phe, this
could be seen as a change of the amplitude of the band correspondmg to the asymmetrk
4.200-stretch for the Z and A forxns of the molecule. At pH = 7 the zwitterionic form
3
1I
2000
1900
1800
1700
1600
1500
1400
1300
Wavenumber / cm-'
Figure 6-7 Transmission spectra of Phe in solutions of DzO with pH ranging fiom 7 to 1 1.
The pH vaiues are indicated at the corresponding spectra.
2000
1800
1600
1400
Wavenumber /cm"
Figure 6-8 SNIFTIR spectra over the range of fiequencies, 1300 to 1800 c ~ acquired
'
using the s-polarized i.d?ared beam, for the Au(ll1) electrode in 0.1 M KCIOI + 1-06 x
1 o ' ~M Phe solution in DzO.For each spectrum, the reference potential Ei was equd to
-0.60V/SCE, and the value of the E 2 is indicated in the figure.
#
dominates and only the 400stretch at 1616 cm-' characteristic for form Z is seen in the
spectrum. In contrast7the anionic form predorninates in a solution of pH = 11 and here
only the band at 1574 cm-' is observed. At intermediate pH values, the concentrations of
zwitterïonic and anionic form are of comparable magnitude and here the I R spectra
contain the two 400' stretchuig bands in the -1 600 cm-' region.
(iii) SNIFTIRS spectrafor s-polmized photons
For s-polarized light, Figure 6-8 shows the evolution of SNIFTIRS spectra of Phe
as a fùnction of the sample potential E2in unbuffered O.lMKC104 in D20. Since for spolarized light with SNIFTIRS the spectrum is that of solution species, and therefore in
principle, their shape shouid not Vary with E2.Ody the amplitude of bands in these spectra
should change with Ez,due to the change of T with the p o t e n a [3,5,6,20]. This is indeed
the case for the spectral region below 1500 cm-'. Integrated intensities of the 1410 cm-'
bands are plotted versus E2,dong the d a c e concentrations in the inset to Figure 6-3.
Clearly, the specttroscopic data track weil the surface concentration plot.
However, the dative intensities of the 1616 and 1576 cm-' bands change quite
dramaticaily with the electrode potential. Using Figure 6-7, these changes may well be
explained in terms of a variable pH in the thin layer cavity. At the most negative potentials
the amplitudes of the 1616 and 1576 cm-'bands have comparable magnitude. This feature
suggests that pH w i t h the thin layer cavity is about 8.5. By moving potential in the
positive direction, the amplitude of the 1616 band decreases and that of the 1576 cm-'
band increases, suggesting that the pH of the solution increases and that Z is progressively
converted into the A form When E2 > 300mV,the 1616 c d starts to grow again at the
r
expense of the 1576 cd
band, suggesting that the pH in the thin layer cavity begins to
decrease. We have already mentioned eariier that the initial increase of pH may be
explained by the reduction of residual oxygen or reduction of water at the reference
potential. Since the spectra were acquired sequentidy, moving fkom the negative Limit in
the positive direction, the increase of pH rnay actually be caused by the increasing the time
of electrolysis rather than by changing the potentiaI. We will show later that the decrease
of pH observed at E > 300 mV correlates welI with the onset of Phe oxidation and
formation of a weak acid such as CO2 in the oxidation reaction. The changes of the
spectral features seen in Figure 6-8 ilIustrate difliculties in studying surface properties of
amino acids adsorbed at the solid surfàce fiom unbuffered solutions. Unfortunately, anions
present in suitable buffers adsorb more strongly than Phe, preventing their use in this
study.
(iv) Orientation of adrorbed rnoZecuZes
The orientation of adsorbed molecules can be conveniently studied using ppolarized radiation The electric field of ppolarized light has non-zero strength both at
the metal surfàce and in the solution For this photon polarization, the SNIFTIRS
spectrum is a dxerence between the spectra of rnolecuies adsorbed at the electrode
surface at potential E2 and spectra of molecules desorbed h t o the thin layer cavity at El.
The measured (ARlR)can be expressed as:
a
where R is the eiectrode reflectivity, 8 is the angle between directions of the electric field
of the photon and the direction in which the dipole moment of the molecule changes, e is
the rnolar absorption coefficient, r is the surface concentration of the adsorbed molecules.
At potential Ei, Phe molecules are desorbed fiom the electrode surface and are randomly
oriented. The fùnction cos0 has to be averaged over al1 possible orientations and the result
is
= 1/3. At potentiai
Et, the magnitude of cos€@)
depends on the
direction of the transition dipole with respect to the normal to the surface because the
electric field of the p-polarized photons is nearly normal to the surface.
Figure 6-9 shows the famiy of SNIFTIR spectra of 0.0 1 M Phe in 0.1 M KC104 in
D20witb p-polarized Iight. Consistent with equation 6.6, the SNIFTIR spectnim for ppolarized photons consists of negative and positive bands. The negative bands correspond
to adsorbed and positive bands to desorbed molecules. The strength of the electnc field of
the p-polarized photon on the electrode sudiace is enhanced ( cm)> &(El)) 1471. Hence,
bands of adsorbed molecules may be stronger than bands of the solution species.
However, the amplitude of the band for adsorbed molecules a h depends on the
magnitude of angle O@).
Vkational modes that have large wmponent of the transition
dipole in the direction perpendicular to the surface are strong, while w'brational modes
with the transition dipole paralle1 to the surface are inactive. In contrast, for molecules
desorbed into the thin layer cavity at El, all IR modes are active.
In the region of 1725 cm-', a bipolar peak appears n o d y assigned to the C=O
stretch in protonated (deuterated) carboxylic acids [45]. In addition, a positive peak at
1574 cm-' assigned to an asymmetric carboxylate stretch and a bipolar peak at
- 1410 due
Figure 6-9 SlVFïIR spectra for Phe adsorbed at the Au(ll1) electrode fkom 0.1 M
KCIOd
+ 1.06 x
1 0 - ~M Mhe solution in DzO using p-polarized uifrared beam For each
spectnim, the reference potential El was equal to -0.60 VBCE, and the value of the E2 is
indicated in the figure.
to the symmetric carboxylate stretch also are observed in the spectra. The simuitaneous
presence of the band due to C=O stretch of a protonated (deuterated) carboxylic group
and the asymmetric and symmetric COO- bands of the dissociated carboxylic group
indicates that the molecule of Phe is undissociated when it is adsorbed at the d a c e (i-e.
as the neutral moIecule) and is dissociated into the solution (A form) when it is desorbed
into the solution In fàct, it is well-documented [48] that, in the adsorbed state, pK, of a
weak acid may dBer fiom the pK, in the bulk- It is therefore reasonable to expect that an
adsorbed moiecuie is protonated at the surface but dissociates when desorbed into the
bulk. We note that the symmetric COOD stretch, seen at
- 1450 cm.' in the transmission
spectnun of the protonated moIecule in Figure 6-6, is absent in the SNIFTIRS spectnim.
This suggests that in the adsorbed molecule the transition dipole of the symmetric COOD
stretch is oriented in the direction parallel to the d a c e and consistent with the surface
selection niles is IR inactive. We also note some differences between SNIFTIRS spectra
recorded using s- and p polarized light. In the spectra acquired ushg s-polarized photons
(Figure 6-8) the 1616 ceL band was present when E c O mV. In the spectra for p-
polarized light Figure 6-9) this band is absent a .E c 300 mV. These ciifferences suggest
that in addition to the species adsorbed in the neutral fom, as discussed previously, a
certain amount of Phe molecule reside on the surfàce in the zwitterionic f o m with the
ND,' and COO' on the surfêce and the plane of the COO- group k i n g nomial to the
surface. In this orientation the permanent dipole of the zwitterion and the transition dipole
of the symmetric COO- stretch are parailel to the surface while the transition dipole of the
asymmetnc COO' stretch has a signifiant component in the direction n o d to the
surfâce. When p-polarized light is used, the bands due to the solution (at El) and
adsorkd species (at E2) cancel each other and this band does not apF
in the SNlFTIRS
specA close inspection of the spectra in Figure 6-9 reveals that amplitudes of the C=O
stretch band and the two COO- bands display different dependence on the electrode
potential. In Figure 640% the peak-to-peak amplitudes of bipolar bands at
- 1725 cm-'(
C=O stretch) and - 1410 cm-' (symrnetric COO- stretch) are plotted against the electrode
potential. It is useful to recall at that point that the pzc in the 1.O6 x 10" M Phe solution is
equal to 150mV. Clearly, the amplitude of the C=O stretch is large at negatively c h g e d
sud'ace, attains maximum at E = -1 00 rnV and decreases at E > - 100 mV. In contrast, the
amplitude of the symmetric COO- stretch is very s m d at negative potentials and begins to
increase to a maximum at
- 200 mV,just positive to the pzc. At these potentials there is a
progressive change in the orientation of the X00-group towards the surface, with the ND,' rotating off the Surface. In this orientation, the transition dipole of the symmetric
COQ- stretch and the permanent dipole of the molecule become aligned in the direction
n o d to the surface while the transition dipole of the asymmetric 400-stretch becomes
paralle1 to the surfàce. This behavior suggests that at negatively charged surface, Phe
molecules display tendency for preferential adsorption in the undissociated - neutrd form
as &H5CH2CH(ND2)COODthat is mixed with the zwitterionic form. By moving Som the
negative end of potentials in the positive direction, the neutrai form progressively
transforms into the zwitterion and the zwitterion changes orientation so that the direction
of its permanent dipole rotates fiorn behg pardel to king normal with respect to the
surface. For cornparison with spectroscopie data, the
CO"
ND,'
1
-600
-400
-200
O
L
200
400
600
E /mV vs SCE
Figure 6-10 For a Au(l I l ) electrode in 0.1 M KC104 + 1.06 x 10-~M Phe solutions, a)
integrated intensities of 1 725 cm-' band (square, Ce
stretch of the deuterated carboqdic
group), 1410 cm" band (circle, symmetnc X O O - stretch) and 2343 cm-' (triangle, COz
band).b) Gibbs excess versus electrode potential plot.
surface concentration of Phe determined fiom chronocoulometrïc data is plotîed in Figure
6-Lob. One can clearly see that the first step o n the suroice concentration plot correlates
weii with the maximum of the band intensity for the neutrd species and the second step on
the r versus E plot correlates well with the s t e p on the band intensiv for the zwittenonic
form. Apparently, the spectroscopic data explaim the complex character of the adsorption
isothenn determined fiom electrochemical measurements.
The bipoIar character of SNIFTIRS bands indicates that the IR bands for the
adsorbed species are shifted to lower frequencies with respect to the IR bands of
molecules in the solution. This shift depends o n the electrode po tential. In Figure 6- 11,
fiequencies corresponding to the minimum of the negative lobe o f the C=O stretch and the
syrnmetric COO- stretch are plotted against the electrode potential.
Apparently,
fiequencies of these two bands move in opposZte direction with the electrode potential.
We do not know whether this shifi is caused by the Stark effect or is due to the miang of
the molecular orbitals o f the adsorbate with the electronic states in the metai.
We suspected that the decrease of ï seen in electrochemical measurements at E >
300mV may be due to the oxidation of Phe at these positive potentials. To
CO-
this
hypothesis, a series of SNIFTIRS measurements were made to determine ifC02 is formed
as a product. CO2has an asymmetric stretching mode [36] at 2343 cm-'. Tnis band shoukl
appear as a negative band in SNIFTIRS, because CO2 is produced only at the potential
(E2).The
broad absorption band of D20centered at 2500 c d [42] interferes with the
2343 cm-' band of C02. Therefore, it was necessary to use &O a s a solvent for these
measurements. Fig. 6-12 shows the SNIFTIR spectra of Phe in the range fiom 2000 to
2600 cm-'. A srnd negative peak at 2343 cm-' appears at EI = 300 mV indicating the
O
200
400
600
800
E / m V v s SCE
Figure 6-11 Dependence of the Eequency of the minimum on the negative lob of the
bipolar bands (closed points) at 1410 cm-' and (open points) at 1725 cm-' on the electrode
potential.
AWR = 1 x lo9 a.u.
500 mV
2343 c\
400 mV
- - -
2600
2500
2400
2300
2200
21 O0
2000
Wavenurnber / cm-'
Figure 6-12 SNIFTIR spectra in the CO2 asymmetric stretch region resuiting fiom the
olridation o f Phe adsorbed at the Au(l11) electrode in 0.1 M KC1O4 + 1.06 x 1 0 - ~M Phe
solution in H20, acquired using s-polarized bfiared beam. For each spectnim, the
reference potential El was equal to -0.60 V/SCE, and the value of the Ez is indicatecl in
the figure.
onset of the oxidation process. The intensity of this peak is plotted against potenta in
Figure 6-1 0. It increases rapidiy with E indicating substantid oxïdation of Phe at the
electrode surface. The increase of the CO2 band intensity correlates very well with the
decrease of the surface concentration of Phe seen in Figure 6-lob. The appearance of the
1616 cm-' band observed at E > 300mV in the SNlFTiRS spectra for both s- and ppolarized light, may be explained by the decrease of pH in the thin layer cavity, due to the
production of CO2, that shifts the equilibrium between the A and Z for-
towards the
Zuritterions.
6.4. Summary and Conclusions
The adsorption of Phe at the Au(ll1) electrode surface has been descn'bed in
terms of Gibbs excesses and Gibbs energies of adsorption. In addition, S
m
S has been
employed to acquire molecular level information concerning the nature of the adsorbed
species and the orientation of the adsorbed molecule at the surface. Figures 6-10a and b
illustrate the complementarity of the eiectrochemical and spectroscopic measmements.
The surfkce concentration of Phe increases with potential up to 200mV and then drops
down. The rising section on the surface concentration plot displays two i ~ e c t i o n s(two
overlapping steps) characteristic of a two state character of adsorption The spectroscopic
data exptain very well this shape of the surface concentration pIot. They show that the fïrst
(most negative) step corresponds to adsorption of the neutral f o m of the molecule. They
also show that the second step is due to the transformation of the neutral into the
zwitterionic form and reorientation of the zwitterionic f o m The decrease of the surface
concentration at E > 300 mV correlates welI with the intensity of the CO2 band and is well
explained by the decarboxyiation of Phe molecules at these positive potentials.
On the bask of information extracted fkom combined thermodynamic and
spectroscopic studies we cari propose the mode1 of adsorbed molecule shown in Figure 613. At negative potentiak (negatively charged surface) Phe is adsorbed Ui the neutral form
with phenyl ring oriented parallel to the surface and with both the -NDî and the 4 O O D
groups interacting with the surface. The plane of the -COOD group is oriented in the
direction normal to the surface. This surface geometry is very shdar to the structure
proposed for gIycine and a-alanine adsorbed at a silver particle by Suh and Moskovits
[49]. In thk orientation the transition dipole of the C=O stretch should have strong
component in the direction normal to the surface while the transition dipole of the
symmetric COOD stretch should be nearly parallel to the surface, consistent with the
SNIFTIRS data. We do not have spectroscopic evidence to support flat orientation of the
, the phenyl ring bands were either weak or hidden within the envelope of
phenyl ~ g since
the asymmeetric COO- stretch bands. However, at the negatkely charged Au(L11) s w k e
aromatic molecules are known to assume a £kit x-bonded d c e geometry [2,3,5] and
these data suggest th& such orientation is energeticaily favorable.
At positive potentiak,
the neutral molecuie is converted into the zwitterion. The molecule is rotated so that the
amino group and the phenyl ring are directed towards the solution and the carboxylic
group towards the metal surface. The plane of the 400- group is nearly normal with
respect to the surface- In this orientation, the transition dipole of the symmetric 400stretch has strong component in the direction normal to the surface. Simultaneously, the
transition dipole of the asyrnmetric COO- is paraiiel to the surface, consistent with the IR
Figure 6-13 Models descniing orientation o f Phe at the Au(ll1) electrode surface; a) at
negative potentials, b) at positive potentials.
data The present results illutrate a need for a concerted use of thermodynarnic and
spectroscopie technique to give a complete description of the adsorption of molecules at
the metal solution interface.
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Chapter 7 Characterization of Phospholipid Bilayers at the
Metal-solution Interface
7.1. Introduction
Bio logical membranes control complicated functions of living ceIls. However,
owing to the complexity of biological membranes, it is desirable to investigate artificialfy
constituted membranes. Artsacial bilayer membranes share many similarities with the
natural membranes [l]. The study of these much simplified models provides valuable
information concerning properties of real bioiogical membranes.
Over the last thirty years, the bilayer lipid membrane has been used to elucidate the
mo lecular mechanisms of biomembrane fiinctions such as ion sensing, material transport,
excitability, gated channels, antigen-antibody binding, signal transduction, and energy
conversion [2-51. A self-assembled bilayer lipid membrane @LM) on a solid/gel
substrate was introduced in 1976 as an attempt to develop a mode1 solar energy
conversion device [6-101. The supported BLM (s-BLh4) was then formed on metallic
wires [Il], conducting Snû2 gIass [12], gel substrates [13-151 and microchips [16]. The
dropping Hg electrode was subsecpently used to deposit phospholipids at electrode
surfàces. The h c e concentration of lipids [17] on the electrode was detennined using
radioactively labeled Zipid rnolecules. The results showed that the fipids concentration on
the electrode was very similar to that at the G-S interface around the pzc. The lipid
moIecules could t u over with the polar head towards the electrode surface, followed by
the formation of a lipid bilayer between the e!ectrode and the electrolyte solution at very
negative potentials [la-201. The eIectrochemica1 properties of the s-BLM were consistent
with those of conventional BLMs [2 1,221.
In the very early studies, the s-BLM was f o d by two consecutive self-assembly
steps [1l,23-27]. Firstl the iipid molecules were placed in contact with a hydrophiIic
rnetal surface. A monolayer of lipid molecuIes was irreversibiy bound to the metal
surface with the nonpolar tail towards the solution Second, the metal covered by the lipid
monolayer was immersed in an aqueous solution. The molecules in solution interact with
the hydrophobie chahs of lipid molecules in the monolayer and assemble to form a
bilayer- This method is no longer attractive because it produces more defects than more
recent methods and a long waiting time is needed. Currently, the Langmuir-Blodgett @B)
technique [28-331 is very popular. One can rnake a Iipid monoIayer on a substrate at
known film pressures by use of a Langmuir trougk The transfer ratio of the LB technique
is close to unity. Other advantages of LB technique [34] include: 1) continuous coatings,
2) very precisely controlled thickness and very good unifonnity of the film, 3) coating of
curved surfâces with q u a i uniformity, 4) amending pores up to the micrometer range on
certain types of substrates. However there are some inherent weaknesses of LB method.
1) The coating involves dipping of all objects into an aqueous sub-phase. 2) In many
cases the substrate
has to be modified to enhiince its adhesion properties. 3) Some
proteins do not toleraîe the transfer process and their activities may be lost. The vesicle
fusion approach allows for easy incorporation of proteins into the bïIayer of the s-BLM
via incorporation of proteins into the bilayer of the vesicles 1351. It is a very promising
technique for biomimetic studies.
Many papers on vesicle h i o n have been published in the last few years. Sackmann
et al. [35-381 studied the mechanism of vesicle spreading on a solid substrate. They found
that four processes must occur in order to form a lipid bilayer: vesicle approach to the
substrate, vesicle adsorption, vesicle rupture and the fusion of adjacent patches of bilayer
on the substrate. They also proposed two dBerent spreading models involving by sliding
and by rolling, illustrated in Figure 2-10. Lipid sliding will occur on substrates that are
very hydrophilic and do not dehydrate the membrane. Kalb et al. [39] compared the lipid
bilayer ftom LB and that fiom indirect vesicle fusion. The fluorescence recovery
experiments showed that the lateral diffusion coefficient was somewhat independent of
the method by which the bilayers were prepared. Groves and Boxer 1401 developed a
simple method to generate the electric field-uiduced concentration gradients in s-BLMs.
The gradient of fluorescently lakled probes was obsewed at field strengths as Iow as 1 V
cm*'.
The objective of our research was to h d an efficient procedure for the deposition
of s-BLMs on Au(li1) surfaces. The DMPC bilayers fiom LB deposition and direct
vesicle fusion d l be characterized by using both electrochemical and in situ FTIR
techniques. Our experimental results will demonstrate how to overcome the problems of
stabliity and reproducibility of s-BLMs. These results are relevant for the development of
biosensors.
7.2. Experirnental
7.2.1 Preparation of SUVs and determination of the vesicle size
SUVs (small unilamellar vesicles) were prepared according to the Barenholz
procedure [41] and described in Chapter 3. Determination of the distribution of vesicle
sizes was performed by the dynamic light scatterhg @LS) as per Hallett er al. [42,43].
Figure 7-1 shows the size distriibution of the DMPC vesicles. The range is narrow. The
0.005
0.01 O
0.015
0.020
0.025
0.030
Particle radius /pm
Figure 7-1 Size distribution of DMPC vesicies measured by dynarnic light scattering
analysis.
mean radius for the vesicles is 14-7 nm which is near the low iimit of srnaII unilamellar
vesicles
w]. Those vesicles are suitable for spontaneous fusion to
an appropriate
substrate.
7.2.2 Measurement of the film pressure by the LB microbalance
The film pressure of DMPC monolayer at the gas-water interface was measured
using a Wilhelmy plate attached to a LB microbalance. The cleanliness of solution and
glassware is critical for the measurement. Even traces of impurities will innuence the
film pressure significantly. Our measurement was made inside the electrochemical ce11
and under a gentle flow of argon. The same measurement was repeated in a closed
chamber without argon flow. The film pressures determined under both cases were
plotted in Figure 7-2. Their values are very close, indicating that the argon flow did not
affect our measurernents.
7.2.3 Electrochemical and SNIFTIRS studies of DMPC Films
The DMPC monolayers and bilayers were deposited on a Au(l11) electrode surface
as described in Chapter 3. The instrumentation for electrochemical and SNIFTIRS
techniques was similar to that used in the previous chapters. However, a jacket ceil was
used here in the SNIFTIRS experiments in order to keep the temperatures inside the cell
constant. Meanwhile, a hemispherical ZnSe window was used to collimate the beam and
to enhance the sensitivity. The spectra repcrted in this chapter were the results of
averaging at least 6000 interferograms. The film pressures of both monolayers and
bilayers were calculated using the charge density data.
7.2.4 AFM studies of DMPC films
In terms of analysing biornaterials in general, the AFM has four major attibutes [64]:
- . - - - -under
Argon environment
under Air environment
38 mN/m
20
Time /Min
Figure 7-2 Film pressures o f DMPC monolayer at the gas-solution (0.05 M NaF)
interface measured by Langmuir-Bbdgett microbalance.
2 . The AFM can achieve very high resolution of sample features in three
dimensions,
2. The AFM can obtain sample topographies without surface treatment or coating.
The surface can be irnaged in its natural state.
3. The AFM can acquire images within a liquid medium Therefore, the AFM
provides a method of imaging surfaces at solid-so lution interfaces.
4. The AFM can be run under a force mode to rneasure the force curves. From the
force vs. separation curve, the thickness of the DMPC bilayers will be measured.
Nanoscope III Atomic Force Microscope was employed in this thesis. The central
component of the AFM is the micro fabricated tip and cantilever. The tip is made of
Si3N4 crystal. The piezoscanner tube is positioned beneath the sample. Scanning is
achieved by moving the sample relative to the statiomry tip. Before obtaining an image,
it is necessary to set up the desired repulsive force between the probe and the surface. By
precisely controllhg the height of the sample relative to the tip, it is possible to maintain
repulsive forces lower than 1 nN. This repulsive force is recorded by the deflection of the
cantiiever which, in tuni, is monitored by the angle of reflection of the laser as measured
by the four quadrant photodiode. If the AFM tip scans over a raised area, the repulsive
force increases. A feedback loop fion the photodiode to the piezoscanner tube then acts
to lower the sample and thereby restores the pre-set repulsive force. 3-D pictures of
constant force are obtained by plotting the sample heigbt against the lateral tip position.
The AFM study in this chapter is preliminary. More measurements have to be done when
we can control the temperature of the solution inside the AFM cell.
7.3 Results and Discussion
73.1 Cbaracterization of the DMPC bilayer by AFM
Initially, we investigated the deposition of DMPC bilayers onto a d a c e of a glass
slide via vesicle fiision We used AFM to image the morphology of the bilayer deposited
on a glass siide and to measure its thickness i?om the force curve analysis. Figure 7-3
shows the AFM images of uncoated and DMPC coated giass slides. The surface became
more flat and unifonn after the deposition of DMPC vesicles. However, some pinholes
could be seen in the film of phospholipids. The discontuiuities in the bilayer were
probably produced during the cooiing of the sample. The fusion of vesicles on the
substrate occmed at 30 OC and the image was acquired at room temperature (approx
23 OC). The change in temperature causes the membrane to shrink and this results in the
holes (defects) in the membrane [45]. For this reason, we did not continue AFM imaging.
It was decided that these measurements should be made at a Iater stage when it is possible
to control the temperature in the AFM ceii. However, the discontinuities in the lipid
bilayer allowed the analysis of thickness to be made [46].
The thickness of the membrane can be determineci by operating AFM in the force
mode. The tip of the AFM was engaged while the lateral scan parameters were tumed offTip deflection versus 2-piezo position c w e s were captured. The tip deflection values
were then converted to forces according to Hookes law (F = -kz). The thickness of the
membrane could be observed in the force versus distance curve. Figure 7-4 shows the
force-distance curve recorded for a DMPC coated glass slide. The DMPC bilayer was
deposited using the LB technique. The major feature is the pronounced "jump on" point.
The presence of a repulsive force at large separation between the tip and the film
Figure 7-3 AFM images (1 p m by 1pm) of a glass slide before (a) and after (b) coating
with 1 mg d-'
DMPC vesicle solution for 1 hour at 30 OC.
Figure 7-4 Extending force as a function of separation distance between the M M tip and
a glas siide coated by a DMPC bilayer using LB deposition.
followed by a sudden attraction is a strong evidence for the presence of an adsorbed
layer. The distance that the AFM tip "jumped over" corresponds to the thickness of the
DMPC bilayer. A similar measurement has k e n made on the DMPC bilayer formed by
the direct vesicle spreading. Figure 7-5 demonstrates a similar force-distance curve. A
uniform thickness for the DMPC bilayer of 6 to 7 nm is consistently observed. This value
agrees very well with the thickness of a DPPC (dipalmitoyphosphatidylchoLine) bilayer
deposited on a silicon wafer [47]. Based on the neutron reflectivity rneasurements for the
DPPC bilayer, Koenig et al. found that the total thickness of the film between the singlecrystalline silicon substrate and the aqueous subphase was 7
and 6.5
* 0.5 n m in the gel phase
* 0.5 nrn in the liquid-crystalline phase 1471. It seems that the thickness of the
supported phospholipid bilayer is weakly dependent on the types of phospholipids and
the methods by which the bilayers have k e n deposited ont0 a solid support. To c o n h n
this point we also deposited an egg-PC bilayer onto the glass substrate by direct vesicle
fusion in silu in the AFM ceLl. The force curve-distance curve for egg-PC in figure 7-6
displays similar features to that observed earlier for the bilayer of DMPC. The thickness
of the bilayer formed by egg-PC is similar to the thickness of the bilayer f o d by
DMPC. It is important to note that there is always a water layer between the solid
substrate and the phospholipid bilayer due to the hydration of the hydrophilic substrate
&or
the polar heads on lipid molecules tumed to the solid substrate. n e thickoess of
the water layer between the DMPC bilayer and the glass was estimated to be 1.7
[48]. The water layer was thicker (3
&
* 0.5 nm
1 nm) when the bilayer was formed on a quartz
substrate [49]. Quartz is less polar than glass. On quartz, more water is needed to b e e r
the polarity of the head of phospholipid molecules in order for the phospholipid to form a
Separation (nm)
Figure 7-5 Extending force as a function of separation distance between the AFM tip and
a glass slide coated by a DMPC bilayer using a direct vesicle fùsion technique. The
concentration of DMPC vesicles is 1 mg m ~ - DMPC
'
in 0.15 M KCl solution-
O
10
20
30
40
50
60
Separation (nm)
Figure 7-6 Extending forces as a function of separation distance between the AFM tip
and a glass slide in an AFM liquid celL The cell was filled up with 1 mg r n ~ "Egg-PC
vesicles in 0.15 M KCI solution for 3 0 minutes, followed by gently flushhg the cell with
90 mL of 0.15 M KCI solution.
stable bilayer. The thickness of the water Iayer has been included in the force curve
measurements. Therefore, the real thickness of the phospholipid bilayer shodd be less
-
than 6 ML The thickness of the bilayer is usually in the range of 4 5 nm [46,47,49] with
a thickness of hydrocarbon tails between 3 to 4 nm [47,49].
7.3.2 Electrochernical studies of DMPC bilayers at the electrode-solution interface
As discussed in section 1-2.4, the eIectrica1 field signif icantly affects the properties
of the biological membrane. For these reasons it is interesting to study artficial
membranes at the electrode surface- Gold electrodes display ideally polarizable behaviour
over a broad range of electrode potentials. Under these conditions, the metal solution
intefiace behaves as a capacitor. The capacity of the electrode decreases when organic
molecules are present at the interface [SOI. The measurement of electrode capacity offers
a convenient tool to study the phospholipid bilayer at the electrode-solution interface.
The capacity can be determined by cyclic voltammetry experiments where a linear
voltage sweep is applied to the electrode- The voltammetric current is equal to the
product of the sweep rate and the differential capacity of the electrode. Figure 7-7shows
the cyclic voltammograms for Au(l1l) with the DMPC bilayer at the electrode-solution
interface. The DMPC bilayers fiom the vesicle fusion and fiom the LB deposition display
similar features. The voltammetric currents in the presence of lipid bilayers are lower
than those in the pure supporting ekctrolyte within a large potential range ( h m -400
mV to 500 mV, SCE)- The two pairs of peaks between -400 and -600 mV overlap with
each other. There is not enough evidence to determine these processes. AFM and
spectroscopy techniques may give more detailed information Tentatively, it may be
assumed that the adsorptioddesorption and/or phase transitions of DMPC bilayers occur
- . - - - -0.05
- M NaF
----
-800
-
BiIayer frorn veside fusion
Bilayerfrom LE deposition
-600
-400
-200
-.-*-
..
O
200
400
600
E/mV vs SCE
Figu re 7-7 Cyclic vo ltammograms of DMPC bilayers at Au(l11 )-solution interface.
in this range of potentials. The peak potentials o f the anodic and cathodic directions dBer
somewhat. The hysteresis indicates slow kinetics o f the inteficial pheno mena
The curves in Figure 7-8 are the daerential capacities rneasured directiy from the
ac voltammetric experhnents. Again, the DMFC bilayers fiom LB deposition and fiom
vesicle fusion show s-imilar patterns and values of the differential capacities at al1
potentials. At extreme negative potentials (less than -700 mV), the differential capacity
curves in the presence of the bilayers merge with the corresponding plot detennined fiom
the phospholipid-free Interfàce. This means h t the phospholipids desorbed fiom the
electrode surface. The deep jump fkom -400 to -500 mV represents the phase transition
of the lipid bilayer. Ho-wever, it is not clear to which state the bilayer changed and what
mechanism fiom bilayer to vesicle is assurned. Firstiy, the DMPC is a double chah
amphiphile. Compared with the head group, its hydrocarbon chains are too buky to form
a srna11 aggregate such as micelle. Secondly, the formation of vesicles eliminates the
energeticauy ~ v o u r a b l eedges at a flinite aggregation number, which is enîropically
favoured. The sIow kinetics of the adsorption-desorption phenomena are responsible for
the absence of the absorption/desorption peaks in the differential capacity curves.
The minimum values of the dif5erential capacity on both curves in Figure 7-8 are
close to 5 pF cm-2which is similar to the value for a nIm of SDS (dodecyl sulfate) at the
Auil 11)-solution interface [SI]. Actually, the capacitance of lipid bilayers at the Iiquidliquid interface is only about 1 pF cme2 [Cl. This value generally depends on the
composition of the bilayer, bathing solution, and Iipid so lvent used, but is independent of
the fiequency (105 to 1,06Hz) and applied voltage (up to 50 mV) during the
-
..
'
- - - - - - 0.05 M NaF
.
.
@
-
--
a
Bilayer from LB deposition
Bilayer from vesicle fusion
---
b
*
I
#
I
,
.
.
-
O
8
.
*
8
-
.
O
8
8
-
I
1
I
I
1
1
E / mV vs SCE
Figure 7-8 Forward differential capacity curves of DMPC bilayer at the Au(ll1)-
solution interface.
rneasurement, There are at least three reasons to account for the difference between our
measurements and the fiterature. 1) The gold single crystal surface is a mosaic of well
ordered (1 11) domains separated by steps, domain boufldaries, etc., and is energetically
inhornogeneous. The defects in the supported bilayer are inevitable. 2) There must be a
water layer between the lipid heads and the gold substrate, where the charged electrolyte
species c m be trapped. 3) Some solvated electrolyte ions are embedded within the lipid
bilayer or attached to the polar heads of the phospholipids. The desolvation may take
place in the presence of electrical fields, which increases the dielectric constant of the
bilayer membrane.
The adsorption of phospholipids not only changes the differential capacity, but &O
the charge density on the electrode surface. The charge density at the electrode surface
may be conveniently measured by using the chronocouiometry technique. Figure 7-9
compares the charge density data for the bilayer fiom LB deposition with that fiom direct
vesicle fusion on Au(ll1). The difference between the two sets of data is extremely
small. Apparently, the direct vesicle fusion on gold is a reliable rnethod for the
substitution of LB deposition. The step in the poten?ial range of -600 to 400 mV shows
the formation of a DMPC bilayer on the electrode surfàce. The bilayers are somewhat
stable in a broad potentiai range (400 to 400 mV) where the charge densities increase
slightly with increasing electrode potentials- At more positive potentials, a change in
conformation of the adsorbed phospholipid molecules may occur. A deep increase in the
charge ciensities is observed. In situ FTIR spectroscopy will be used to probe such
conformation changes in the next section. The phospholipid biIayer has a hydrophobic
interior which does not allow the passage of charged species, However, the supported
+0.05 M NaF
+Bilayer from LB
+ Bilayer from vesicle
-800
-600
-400
-200
O
200
E / m V vs SCE
Figure 7-9 Charge density versus electrode potential plots of DMPC biiayers at a
Au(1 11)-solution interface.
bilayer is always irnperfect because a perfectly smooth substrate is hard to achieve.
Defects (pinholes and pits) in the supported bilayer always exist and this results in a
leakage of the charged species. If an electroactive species penetrates the bilayer to reach
the electrode surface, a certain electrochemical reaction may occur at appropriate
electrode potentials. TI+ has k e n successfülly used to probe the state of adenosine layer
on a Hg electrode surface [52]. BiPotto et 02- 1531 studied the barrier properties of 4pentadecylpyridine coated gold surfaces. They found that the ferricyanide ion (F~(cN)~~-)
was an ideal probe to examine the quality of the organic films on gold electrodes. Figure
7-10 shows some cyclic voltamrnograms used to explain the barrier properties of the
DMPC bilayer at the Au(l 11)-solution interfàce. The ferricyanide ion is chosen because
its reductiodoxidation potential is within the potential range where the DMPC bilayer is
stable. The dashed line in Figure 7-10 shows that the reduction and oxidation potentials
are 145 and 215 mV respectively. The small shoulder peak at about 70 mV in the anodic
half is due to the diffusion of femcyanide ions fiom the buk to the electrode. When the
DMPC vesicles were added to the electrolyte solution, a bilayer formed at the electrodesolution intefice and isoiated the electrode fiom the electrolyte. The redox reaction of
ferricyanide was blocked, and the corresponding peaks disappeared. The CV curve
becarne featureless (solid he). The DMPC bilayer displays a good barrïer propew
although it may contain some defects.
7.3.3 In situ FTIR studies of DMPC bilayers at the electrode-solution interface
The application of i d k e d spectroscopy to the study of the Iipid membrane was
pioneered by Chaprnan and CO-workersin the late 1960s [54]. For a long tirne, the
absorption of water in the IR range prevented the use of the IR spectroscopy to study the
- - - - - - 0.05 M NaF
--Added ferricyanide
Presence of DMPC bilayer
---
-800
-600
-400
-200
O
200
400
600
E /mV vs SCE
Figure 7-10 Cyclic voltmograrns of Au(ll1) in 0.05 M NaF, 0.1 rnM Femcyanide
and 0.1 mM Femcyanide with a DMPC bilayer formed by a vesicle fusion technique.
Scan rate was 20 mV s-' .
properties of phospholipids. The combination of the attenuated total r e k t i o n and
Fourier transform techniques proved to be a way to overcome this problem [55,56].
Duevel et al. [57]studied the orientation and organization of phospholipid mono layers at
gold surfaces by using polarization modulation FTIR spctroscopy. Here we will use
SNIFTIRS to characterize the DMPC bilayer on the Au(ll1) surface at different
electrode potentials.
Figure 7-11 shows the SNLFTIRS spectra of DMPC at different sample poteniials.
The spectra were recorded by stepping the potential between the sample potential
indicated in the figure and a reference potential, -700 mV at which DMPC molecules are
desorbed fiom the electrode surface into the thin layer cavity of the IR ceiI. The SNIFTIR
spectra effectively plot a merence between the spectnun of DMPC desorbed to form
vesicles at the reference potentid and the spectnim of a bilayer of DMPC spread at the
goid d a c e at the sample potential. There are no absorption bands at E = -600 mV. This
behaviour indicates that no DMPC bilayer is formed on the goId surface at this potential.
The characteristic bands of IR spectra for phospholipid molecules appear when the
electrode potential is higher than 4 0 0 mV. Table 7-1gives the vibrational assignments
for the surfàce-active modes of DMPC in a bilayer. T'ose fiequencies are very close to
the correspondhg vibrational modes in DPPC.
The asymmetric (-2965 cm?) and
symmetric (-2880 cm-') stretching modes for CH3 are misshg in the SNIFTIRS spectra
The reason is either that the absorption is weak or that the vibrational modes are surfaceinactive. The intensities of CH2 asymmetric (2922 cm-') and symmetric (2852 cm-')
stretching increase as the electrode potential is made more positive. However, their peak
positions change little with appiied potentials. This indicates that the interactions between
AWR = 4 x lU3 a.u.
2500
2000
1500
Wavenumber /cm"
Figure 7-11SNlTXR spectra of DMPC bilayers at Au(1 1 1)-solution interface made with
the vesicle fusion technique. Potentials are shown in the graph
Table 7-1 Vibrational assignments for the surfâce active modes o f DMPC bilayer
vibrational modes
assignments
DPPC,cm-'
2922
va(CH2)
asym CH2 stretch
2925 [57]
2852
vdCH2)
sym CH2 stretch
2855 [57]
1742
vc=o(co;?R)
sn-1 C=O stretch
1742 [58]
1730
vc=o(CQzR)
sn-2 C=O stretch
1727 [58]
1466
WH2)
CH2 bend
1468 [5Tj
1230
va(PO23
asym PO2-stretch
1231 [58j
1086
v~(pOt?
sym PO2-stretch
1085 [59]
DMPC, cm-'
the CH2 and the electrode surfàce are weakly affected by the electrode potential.
Actudy, the akyl chahs are relatively far fiom the electrode sudace, as expected for a
bilayer modeL Both asymmetnc (1230
cm1)and symmetric (1086 cm-')stretching of
PO2- groups show bipolar features in the spectra in figure 7-1 1. The bipolar shape usually
suggests a strong (short range) interaction between the adsorbed molecules and the
electrode d c e . Those results are consistent with our assumption that the polar heads of
the f k t DMPC Ieaflet are attached to the electrode surface. The C=O stretching band of
DMPC comprises of two components which c m be assigned to sn-1 (1742 cm") and sn-2
(1730 cm-') C=O groups [58]. The split of the two peaks is clear at 400 mV in Figure 71 1. As the potential increases, the relative peak heights of these two bands change, while
there are no fkequency shifis of the individual components. The changes observed in the
C=O stretching band indicates that the re-arrangement of the bilayer affects the glycero1
mo iety.
The increase in band intensity with electrode potentials can be attributed to two
factors. 1) More molecules adsorb ont0 the electrode at more positive electrode
potentials. 2) A certain type of vibrational modes is favoured by the orientation change of
the adsorbed molecules. Figure 7-12 correlates the charge densities for DMPC bilayer to
the IR band intensities at 2852 cm-' (CH2 symmetric stretch). A good agreement has been
achieved at E < 500 mV- When the electrode is more positive than 500 mV, the
electroporation (see section 1.2.4 in Chapter 1) or breakdown of the bilayer probably
occurs. Leaking of charge results in a steep increase in the measured charge densities on
the metal side. Figure 7-13 shows that the IR intensities of both symmetnc and
asymmetric stretching modes increase with the hcreasing electrode potential. If the
increase of the band intensity results o d y fiom the increase of the number of adsorbed
molecules, the ratio of the intensity at 2922 cm-' to the intensity at 2852 cm-' should be
constant. Actuaily the ratio decreases with increasing electrode potential as shown in the
inset of Figure 7-13. This behaviour suggests that the allcyl tails of adsorbed DMPC
molecules change codormation as a function of applied potentials.
7.3.4 Cornparison of DMPC bilayers and monolayers at electrode-solution interface
The DMPC monolayer has been made by using the horizontal touch technique (see
section 2.4.2 in Chapter 2). Here the horizontally orientated substrate is lowered
vertically and is brought into horizontal contact with the compressed monolayer at the G-
S interface (Figure 2-8a). The meniscus was made thereafter. Al1 the electrochemical
measurements were performed without tramferring the electrode to another cell. The
horizontal touch method has proven to be an efficient way to transfer a monolayer of
amphiphiles fkom the G-S interface to the metai-solution interface. The DMPC
IR band intensity at 2852 cm"
-800 -600
-400
-200
O
200
400
600
E /mV vs SCE
Figure 7-12 Correlation of the charge densities for DMPC bilayer on Au(ll1) to the IR
band intensity at 2852 cm".
A
V
peak at 2922 cm''
peak at 2852 cm-'
E /mV vs SCE
Figure 7-13 Dependence o f CHz stretchuig modes on the electrode potentials for DMPC
bilayers at the Au(1 l 1)-solution interface. Inset: ratio o f IR intensity at 2922 cm-' to the
intensity at 2852 cm-' versus the electrode potential.
mono layer fiom horizontal touch has a mo lecular orientation with their hydrophobic tails
towards the electrode d c e . However, the electrode potential may change the
orientation of the monolayer. Nelson and coworkers [60-621 proposed a flipping
mechanism for the adsorbed phospfiolpid rnonolayer on a mercury electrode. Figure 7-14
compares the differential capacities of DMPC monolayers and that of bilayers. They
show completely different features although al1 the curves merge at -750 mV denoting a
total desorptioa The minimum value for monolayers is much higher than that for
bilayers, indicating that the bilayer has a lower perrneability. There are two peaks on the
dserential capacity curve for the monolayer. The quick orientation change rnay be
responsible for the large peak in the range of -250 to -1 00 mV. The slow adsorption
/desarption process results in the s
d peak in the range of -350 to -250 mV.
Figure 7- 15 is the plot of charge density data measured fiom the chronocoulometry
experiments for DMPC monolayers and bilayers. Those charge densities represent the
total charge on the metal side of the double Iayer. The total charge (Q) mntajns two
contributions as shown in equation 7.1.
Q = CoE + L.FCD-
(7-1)
Where C is the infinite fkequency capacity, E is electrode potential, F is the Faraday
constant,
L is the electrosorption valency and
r-
is the SUfface concentration of
adsorbed DMPC molecules. The hrst term (CoE) is responsible for the charging of the
double Iayer (charging of the infinite fiequency capacity). And the second term
@*ForD-) represents the charge due to DMPC adsorption. The charge densities for the
mondayer increase g r a d d y up to about 400 mV. They increase drastically when the
potential is greater than 400 mV. The steep change could be the results of the oxidation
*....
- - - - - - 0.05 M NaF
-
----
DMPC monolayer
DMPC bilayer
--
1
l
.
.
1
s
I
I
e
-
I
,
-
.
8
.,
1
8
.
\;
I \
I \
I l
-
..
.
.
r
r
I !
I I
1
-
-800
t
I
-600
-400
I
-200
1
1
O
200
I
E / mV vs SCE
Figure 7-14 Cornparison of the differential capacities for DMPC bilayers with that for
DMPC monolayers at the Au(1 l 1)-soIution interface.
-
+0.05 M NaF
DMPC bilayer from vesicle
b DMPC monolayer
E / mV vs SCE
Figure 7-15 Cornparison of charge densities of DMPC bilayers with that of DMPC
monolayers at Au(1 11)-solution interface.
of the electrode surfàce because it follows the trend of the pure supportkg electrolyte.
Another explanation may be the break-down of the double iayer capacitor at high
potentials as described earlier. The oxidation of DMPC molecules is unlikely as shown
by the FTIR spectra at those potentiats. There is a step O bserved for -600 <E< 4 0 0 mV
on the charge-potential curve recorded for the bilayer. The step is followed by a linearly
increasing section up to 300 mV. The second term in equation 7.1 is responsible for the
presence of the step on the charge-potential curve. The step on the charge-potential curve
for the monolayer is more graduai and the linear section where charge changes slowly
with the potential is restricted to a narrow potential range, O to 400 mV, where the
dserential capacities change little.
The product of charge and potential is equal to energy. Therefore, the area
contained between the charge density curves correspondhg to the film-fkee and the filmcovered electrode is equal to the change of surfàce energies due to the adsorption of
DMPC molecules. In fact, this energy change is equal to the pressure of the film at the
electrode-solution interfàce [63]. As descriid by equation 2.5, the film pressures
calculated for the monohyer and the bilayer are plotted in Figure 7-16. AU curves are
parabolic in shape and are strongly dependent on the electrode potential. The maximum
film pressure is about 60 mN m-'at 300 mV for a DMPC bilayer, and 39 mN rn-' at 200
rnV for a rnonolayer. Theoretically, the maximum film pressure of monolayers at the
electrode-solution interface should be the same as the film pressure of a compressed
monolayer at the gas-solution interface. In order to measure the film pressure of a
monolayer of DMPC at the gas-solution interface, the LB microbalance was used. Figure
7-16 shows the film pressures at the gas-solution interflice measured over a period of 70
-800
-600
-400
-200
O
200
400
600
E / m V v s SCE
Figure 7-16 Cornparison o f the film pressure for DMPC bilayers with that for DMPC
monolayers at Au(1 1 1)-solution interfaces. Dot: bilayers by LB deposition; Square:
bilayers by vesicle fusion; Triangle: monolayers by horizontal touch.
-800 -600 -400 -200 O
E /mV vs SCE
10
20
30
40
50
200 400
60
70
80
Time /Min
Figure 7-17 Cornparison of the film pressures for DMPC monolayers at the gas-solution
with that at the rnetal-solution (inset) interface.
minutes. A constant value o f 37 mN II?
was observed. This result is quite close to the
maximum film pressure at the electrode-solution interfiace (as shown in the inset of
Figure 7-17). The consistency of the film pressures measured at the gas-solution and the
electrode-solution interfaces indicates that the film was transferred fiom the G-S to the
electrode-solution interface with the transfer ratio of 1:l. It is usefùl to note that the film
pressure for the bilayer is not exactly the double of the film pressure of the monolayer.
This behaviour indicates that the bilayer does not simply duplicate the monolayer; Le. the
physical properties of one leaflet in the bilayer d s e r from that of the rnonolayer. In the
case of a monolayer, the hydrophobie tail of the DMPC is facing the electrode sudace. In
contirist, the polar head o f the DMPC is facing the electrode surfàce when a bilayer is
formed at the electrode-solution interface.
7.4 Conclusions
Both LB deposition and direct vesicle fusion techniques have been employed to
deposit the DMPC bilayers at the gold electrode-solution interfàce. The thickness and the
electrochernicd properties o f DMPC bilayers are independent of the methods used for the
deposition However, the direct vesicle fûsion technique facilitates the application of
SNIFTIRS which can provide detailed information concerning the orientation of
phospholipid molecules. Another advantage of the use of vesicle fiision is the ease in
which proteins can be incorporated in the phospholipid bilayers. Some proteins may not
tolerate the transfer process during the LB deposition while others need to be embedded
into the bilayer at the tirne of vesicle preparation. The success of the direct vesicle fusion
technique opens a way to initiate new directions of fundamental research In the hture,
this technique may be used for development o f biosensors,
The bilayer o f DMEC is formed in the potential range between -600 to 4 0 0 mV.
The measured differential capacity of DMPC bilayer at the gold-solution intertàce is
approximately 5 pF cm-2. This value is much higher than the capacity of the lipid bilayer
at the solution-solution interface (4 pF
The higher value of the capacity may be
due to the trapped electrolyte, defects in the DMPC bilayer and the imperfection of the
electrode surface. The bilayer is fâirly stable up to 400 mV. Its barrier property was
tested using ferricyanide as a probe. The formation of the DMPC bilayer prevents the
diffusion of femcyanide ions to the electrode surface. Therefore, the reduction/oxidation
peaks in the cyclic voltammograms of this probe are absent when DMPC vesicles are
present in the solution.
The maximum values of the film pressure for the DMPC bilayer and monolayer are
60 and 39 mN m-' respectively. The film pressure for the monolayer at the electrode
surface agrees weU with the value of the film pressure for a compressed monolayer at the
gas-solution interface. The film pressure of the bilayer is not equal to double that of the
monolayer. This is because the bilayer does not simpiy duplicate properties of the
monolayer. The bilayer has a tail-to-tail cofigmation with the polar heads towards the
electrode surface.
The in situ J?ïR technique complements the electrochemical studies. It provides
additional information c o n c e d g the molecular organization of the DMPC bilayer. This
technique is particularly suitable for monitoring subtle changes in the bilayer structure
induced by changing the electrode potentiai. The SNIFTIRS spectra demonstrate that the
re-arrangement of the DMPC biiayer occurs when the electrode potential is increased. It
aiso shows that the polar heads of phospholipid molecules are attached to the electrode
d a c e at alI adsorption potentials.
The formation of the phospholipid bilayers at the
electrode surface is, therefore, confirmed.
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Chapter 8 Summary and Future Research Directions
8.1 Summary of the Thesis
This thesis extended the research on ionic and molecular adsorption at solid
electrodes carried out in our laboratory to a class of important biological molecules.
Previously in our laboratory, the adsorption and CO-adsorptionof inorganic ions had been
systematically investigated on gold [1-31 and platinum [4,5] electrodes. The results
showed that the adsorption of both anions and cations was highly dependent on the
crystallography of the electrode surface, the bu& concentration of ions of interest, the
c o ~ s i t i o of
n the electrolyte solution and the electrode potential applied. On the other
han&
the interfacial behaviour of srnall organic molecules [3,6,7] and insoluble
surfactants [8,9] at the gold-solution interface have d s o been studied by using
electrochemical and spectroscopie techniques- These studies were tailored to investigate
surface behavior of molecules that are used as polishing and leveling agents in
electroplating and e l e c t r o r e ~ g My
. research extended this broad research program
onto studies of s d c e behaviour of amino acids and phospholipids. This constitutes a
new thrust in the direction of biomimetics. Its long-term goal is to d
c the properties of
biologicat membranes deposited at the metal-sohtion interfàce.
In this thesis, the direct vesicle fusion technique has been developed for the
deposition of phospholipid bilayers at the metal-solution interface. This technique was
then employed to study the interfacial behavior of DMPC by using a combination of
electrochemical, AFM imaging and in situ =IR
techniques. The direct fiision technique
was compared to the conventional Langmuir-Blodgett deposition Both AFM and
electrochemical studies showed that these two methods produced Lipid bilayers of similar
properties. However, the direct vesicle fùsion provides a convenient means to insert other
membrane components such as cholesterol and proteins that rnay be explored in fùture
studies.
1 aIso investigated the adsorption of amino acids such as phenylalanine.
Pheny [alanine is an amino acid that has hydrophobic propert ies. Its hydropho bic
properties play an important role in the tertiary structures of proteins. The investigation of
the adsorption of phenylalanine is the fïrst step towards the understanding of the tertiary
structure of proteins in the adsorbed state at solid surfaces. The adsorption of an amino
acid is quite cornplex, Ln order to facilitate the interpretation of the spectroscopic data for
phenylalanine, additional studies of benzoate and phenylacetic acid adsorption at the gold
electrode surface have been performed.
The major conclusions in this thesis can be sumuiarized as follows. For the
deposition of phospholipids at the gold-solution interface:
ï) DMPC bilayers cari be deposited onto the gold electrode surfàçe by using both
LB deposition and direct vesicle fision techniques. The bilayers deposited using different
methods showed similar thickness and electrochemical properties. However, the vesicle
fiision approach faciiitates the application of SNIFTIRS in studies of iipid bilayers.
Another advantage of the vesicle fùsion over the LB deposition is that it provides a
convenient rneans to insert other membrane components into the bilayer. This is the key
issue for the fùture researches in this area
I
ii) The DMPC bilayer showed very good b e r properties. The reduction-oxidat ion
peaks of ferricyanide were blocked by the vesicle fùsion onto the electrode surface.
However, the daerential capacity of the DMPC bilayer at the gold-so Iution interface was
about 5 pF cme2, which was much higher than the capacity of a lipid bilayer at the
solution-solution intefice (-1
crK2) [15]. Defects in the bilayer a d o r on the
electrode surface were responsible for the high value of dserent ial capacities.
For the adsorption of phenylalanine and benzoate:
i) Both benzoate and phenylalanine favored a flat orientation (x-bonding) at
negat ively charged surfaces, and a vertical (tilted) orientation at more positively charged
surfaces with the carboxylic groups attached to the eIectrode surface.
ii) A change in the electrode potential promoted the transformation of phenylalanine
fiom the neutral state at negatively charged surface to the zwitterionic state at a positively
charged s d a c e .
This transformation is coupled with the potential-controlled reorientation of the adsorbed
molecuIe. At a positively charged d a c e , phenylalanine is turned with the negative
carboxylate group towards the surface and the positive protonated amine group towards
the solution.
iii) The benzoate ion is chemically stable in a broad range of electrode potentials. In
contrast, decarboxylation of phenylalanine was Observed the electrode potentials greater
than 300 mV (SCE). CO2 was detected as the oxidation product of phenylalanine at these
positive potentids.
vi) intensities of IR bands for adsorbed benzoate and phenylalanine correlated very
well with surface concentrations of adsorbed pheny lalanine and benzoate ions. The
combination of electrochemica1 and in situ FTIR techniques provides complementary
information conceming the energetics and mechanism of molecular adsorption at the
metal-so lution interface.
8.2 Future Work Directions
As mentioned earlier, the technique of direct fusion of vesicles that was developed
in this thesis opens a possibility of new biomimetic research. Below 1 outline three
directions in which this project may develop in the fùture.
i) Mechanistic studies of the formation and phase transitions of the supported
phosphoüpid bilayer: We w w know tbat the phospholipid bilayer can be deposited at a
well-defined gold single crystal electrode using the direct vesicle fusion technique. The
electrode potential effectively influences the spreading process of vesicles. It should also
affect the structure of the bilayer. AFM and STM were used recently to study the
potential-controlled phase transition in a film of sodium dodecylsulfate (SDS) at the gold
electrode surface [16]. Those techniques can dso be applied to study the structure and
potential controlled phase transitions in phospholipid bilayers.
In addition, further
mo lecular level infiormation concerning the properties of the supported bilayer may be
O btained
using in situ FTIR spectroscopy. Our specidy designed ce11 for FTIR allows
programming of temperatures. It is, therefore, possible to examine the effect of
temperature on the phase transitions in the supported bilayer and the orientation of the
phospholipid molecules in the bilayer at dBerent electrode potentials. These studies can
provide information concerning the choice o f the optimum temperature and potential for
the deposition of a lipid bilayer at the electrode surfàce.
ii) Modification of the phospholipid bilayers with cholesterol: Biomembranes
are generally in the Buid Iiquid crystalline phase and the membrane fluidity is critical to
preserve its hctions. The fluidity of the naturat bio logical membrane is preserved due to
the presence of cholesterol.
Cholesterol constitutes about 30% of the mass of the
membrane lipids of muiy animal ce11 plasma membranes. The incorporation of
cholesterol into phospholipid bilayers is expected to increase van der Waals attraction
[17], decrease electrostatic repulsion due to an increase the area per lipid molecule
[18,19] and mod@ the short-rang repulsive force [19,20]. Therefore, the incorporation of
cholesterol into the mode1 membranes (vesicles or planer bilayers) should affect the
balance of intermolecular forces and alter the stabïlity of the bilayers and the fusion of
vesicles ont0 a solid support. It would be of great interest to study the properties of
phospholipid biiayers containing cholester01 mo lecules by using both electrochemical
and spectroscopie techniques.
iii) Ion transfer across the phospholipid bilayers: An important fiinction of the
bio logical membrane is to provide selective permeability to ions. The hydrophobic core
of the biIayer is a very effective barrier against the passage of inorganic ions although it
allows nonpo lar so lutes to pass across the membrane. The permeability to inorganic ions
is mainiy due to the presence of ion channels in the bilayers. The DMPC bilayer
deposited by the vesicle fiision technique had good barrier properties probed using the
d
ferricyanide reduction reaction (Chapter 7, section 7.3.2). This bilayer may, therefore, be
used as a m a t e for the incorporation of ion channels. Three typical ion channels that
may be used for that purpose are (1) nystatin, a polyene antibiotic which forms a complex
with cholesterol (Figure 8-l), (2) gramicidin A, a peptide which forms monovalent
cation-selective channels (Figure 8-2), (3) alanethicin, a peptide which forms voltagegated channels (Figure 8-3). Electrochemical techniques may be used to study ionic
transport across the bilayer with those ion channels. These studies may provide valuable
information conceming the mechanism of the ionic transport across the charme1 and the
effect of the electric field at the interface on the rate and mechanism o f the ion transport
through the channel,
Figure 8-1 Chemical structures of nystatin and a ''barre1 stave" mode1 of the channel
formed with nystatins. The protuberance on the bottom represents the amino sugar and
the shaded interior represents the hydrophilic polyhydroxyl portion of the molecule. The
exterior surface of the channel is completely nonpolar. (Adapted from ref. [2 11.)
Figure 8-2 Illustration o f the structure of the channel formed by a dimer (right)
gramicidin A (lefi) in the bilayer. (Adapted fiom ref. [2 11.)
Ac- Aib-Pro-Aib-Ala-Aib-Ma-GIn-Aib-Val-
Aib-Gly-Leu-Aib-Pro-Vai-Aib-Glu-Glu-Glu-PH
(A ibr aminobufyrïc acid. Phl: L-phenylalaninol.)
Figure 8-3 Proposed structures formed by an oligomer (bottom) of alanethicin (top) in a
Iipid bilayer. Le#: The conformation in the absence o f an applied tram-membrane
voltage. Righl: The open channel conformation in the presence o f an applied voltage.
M m e : An intermediate conformation. (Adapted f?om ref [21].)
References
Zhichao Shi, PkD. Thesis, University of Guelph, 1996.
Shijie Wu, Ph.D.Thesis, University of Guelph, 1996.
Aicheng Chen, PhD.Thesis, University of Guelph, 1998.
Walter Savich, PbD. Thesis, University of Guelph, 1997.
Nanhai Li, PhD.Thesis, University of Guelph, 2000.
Lome Stolberg, Ph.D. Thesis, University of Guelph, 1990.
Dongfang Yang, Ph-D. Thesis, University of Guelph, 1995.
Dan Bizzotto, Ph.D. Thesis, University of Guelph, 1996.
Vlad Zamlynny, M-Sc. Thesis, University of Guelph, 1998.
10. K-Q. Li, S. G. Roscoe, and J. Lipkowski, J. EIectroanal.Chem. 478 (1999) 67.
I l - Z. Shi, and J. Lipkowski, J. Electroanai- C h e a 403 (1996) 225.
12.2. Shi, J. Lipkowski, S. MïrwaId, and B. Pettinger, J. Chem. Soc. Faraday Trans. 92
(1996) 3737.
13. D.G. Marangoni, RS. Smith and S-G. Roscoe, Cam J- Chern 67 (1989) 921.
14-D.G. Marangmi, I.G.N. WyLie and S.G- Roscoe, Bioelectrochem and Bioenerg.
25 (1991) 269.
and Practice, 1974, Marcel
15. H.T. Tienw.) Bilayer Lipid M e m b r a n e s @ L ~ T h e o r y
Dekker, Inc. pp8.
16.1. Burgess, C . JeBey, G. Szymanski, 2. Galu and J. Lipkowski, Langmuir, 15
(1999) 2607.
17. S. Nu, Prof S d . Sci. 8 (1 977) 1.
18. P.L. Yeagle, Biocbim Biophys. Acta, 822 (1985) 267.
19- T.J. Mchtosh, AD. Magid, S A Simon, Biochemistry, 28 (1989) 17.
20. L.J. Lis, M. McAlister, N. Fuller, RP. Rand, V.A Parsegian, Biophys. J. 37 (1982)
667,
2 1. R.8. Gennis(Ed.), Biomembranes-Mo lecular Structue and Function, 1989, SpringerVerIag, New York-
Appendix
Error analysis
The thermodynarnic data analysis presented in this thesis involves at Ieast one
integration and one differentiation step. It is important to estirnate the error propagation in
the data processing in order to mess the reliability of the calculated quantities. The detailed
error analysis has k e n discussed extensively by L. Stolberg [l], J. Richer [2] and D. Yang
[3]. Shce similar techniques and data processing procedure were used in this study, the error
analysis d e m i h d by those authors is applicable to this study. In the following section, a
brief sumrnary of this topic wilf be given.
Error in charge density measurement
The value of relative charge density (AcM)is determined by integration of the
current transient. The accuracy of the current integration procedure can be determined with
the help of a dummy ceU It is found to be equal or less than 0.05 pC cmm2.
The accuracy in
the charge density is not chiefly determined by the arors associated with the current
integration, but by the errors associatecl with reproducing the contact area between the single
crystal d c e and the electrolyte solution and reproducing the surface structure using the
flame annealing procedure. The reproducibility of the contact area is approximately 1%.
However the precision can be improved if the whole series of measurernents is made on an
electrode that is brought in contact with the electrolyte solution only once [4]. In such a case
precision on the order of 0.1 pC c ~ can
' be acachied. The accuracy of AoMwas found to be
e . 7 % for organic molecules [l]. The accuracy of absolute charge density not only depends
on the accuracy of the relative charge density but also on the accuracy of the value of pzc.
The value of p z is determined fiom the position of the diaise hyer minimum of the
dinerential capacity versus electrode potential curves. The accuracy in the value of the pzc
was estirnated to be M.005 V.
Error in the calculation of film pressure
The film pressure at constant potential was calculated by htegration of the charge
density with respect to potentials using equation 2.13. The error propagation in the data
processing has k e n discussed by Oldham [SI. The error of integration can be estimated by
where S, is the standard deviation of the charge density, AE is the potential range over wirh
the integmtion was performed, and N is the number of charge density points taken in the
integration. The equation predicts an increase of S, when the electrode potential becornes
more positive. If the standard deviation of the charge density was assumed to equal to 0.1
pC
and the AE of 0.5 V was used, the error in nImpressure amounts to 0.14 mN m-'.
Emor in the calcuIation of reIative Gibbs excess
The relative Gibbs excess was determined either by graphical differentiation of the
x (or a)versus h(c) curves (Eqs. 2.21 and 2.22). The standard deviation of relative Gibbs
excess determined by the fkst method is given by the express [SI:
where S, is the standard deviation of the film pressure and A is the increment of RTln(c).
r,),
The Sr depends on the magnitude of b For the values of T
c
detennïned by this method, are estirnated to have an error of &IO%.
The error associated
with
r,,
which are
which is determined by linear regression of the linear section of the x versm
potential curve, has an error approximately *5% [l].
Error in the calculation of Gibbs energy
The Gibbs energies of adsorption is related to the adsorption coefficient P. The
P
was determined fiom the initial dope of the ir versus X (molar fkction of organic molecule)
plot using Henry's law or fiom a fit of T verjus h(c) curves with the Fnimkin Wthenm For
this procedure, the error of Gibbs energy can be calculated fkom the formula 121:
where Mn:and 6r-/ïnrax
are the relative errors for the film pressure and the maximum
SUTfàce concentraiion of the adsorbate, respectively. For n l e s than 1 mN m-', the relative
e m r is on the order of 3Ph. Since the relative error in,
,
T
is 10%, the un&@,
~AGO~~S
is equal to 1 kJ mol-'.
Errors in the calculation of other Damneters
The shifi of p z , EN,is equal to the difference between the pu: for an electrode
which is covered by a monolayer of organic molecule @ m l ) and the pzc observed in the
absence of organic molecule (pz-).
uncertainty of pzc+l and pz-.
Therefore, the error associateci with EN depends on the
The relative e m r of EN is estirnated to be 3.8%.
The electrosorptionvalencies (y3 were calcuiated f?om the sbpes of the m versus r
plots. The uncertauity in y' depends on the standard deviation of the dope of CM
- r plot.
The relative error is approximately 5% for gold single crystd electrode in the presence of
organic molecules [2].
References
1. L. Stolberg, PbD thesis, University of Guelph, 1990.
2. R J. Richer, PbD thesis, University of Guelph, 1990.
3. D. Yang, PbD, thesis, University of Guelph, 1995.
4. k Hamelin, S. Morin, J. Richer and J. Lipkowski, Electroanai. Chem 304 (1 99 1) 195.
5. KB. Oldham, JElectrodChem, 208 (1 986) 1.