Potentiometrie titrations of five synthetic tetraacids as models for indigenous C80 tetraacids

Article abstract of DOI:10.1021/la902326y

The acid/base properties, critical micelle concentrations (cmcs), and. pH-dependent solubility of five synthetic tetraacids have been studied at several ionic strengths (20-600 mM NaCl) and in the pH range of 1.5-11 using high precision Potentiometrie titrations, tensiometer measurements, and. UV spectroscopy, respectively. The molecular weight of the tetraacids ranged, between 478 and. 983 g/mol. The Potentiometrie titration data was evaluated, in terms of thermodynamic equilibrium models, developed in the light of relevant solubility data, Langmuir monolayer compressions and cmc of the different tetraacids. The results indicate that for two of the tetraacids, called BP5 and BP7, two chemical forms fully dominate the speciation of the monomers; the insoluble fully protonated form, and the soluble fully deprotonated form, The partly protonated species, only play a very minor role in the speciation of these tetraacids. For the other tetraacids the results are more complicated; for the smallest tetraacid, called BPl, all species seem to play important roles, and for the most hydrophobic, BP10, the formation of micelles and aggregates severely complicates the evaluation of the speciation. For the tetraacid BP3 one of the partly deprotonated forms seems to be important, thus confirming the structure to properties relationship. In spite of the complicated micelle formation chemistry, and although not actually measured, the acid/base properties for the monomers of BPlO were interpreted by means of surface charge densities of the micellar aggregates. The modeling indicates an increase of the aggregation number of the micelle upon acidification, a result of formation of mixed micelles incorporating the fully protonated and deprotonated species. An intrinsic pK_a of 5.4 for BP5 was used, to model the monomer pk _a of BP10, and corresponded well with a monolayer acidity constant pk^s_a of 5.5 obtained from, surface collapse pressures of Langmuir monolayers as a function ofpH. 2010 American Chemical Society.

Full text of DOI:10.1021/la902326y

pubs.acs.org/Langmuir
© 2010 American Chemical Society
Potentiometric Titrations of Five Synthetic Tetraacids as Models
for Indigenous C80 Tetraacids
‡
‡
‡
Ola Sundman,*,† Erland L. Nordga
˚
rd, Brian Grimes, and Johan Sjoblom
€
†
Department of Chemistry, Umea University, SE 901 87, Umea, Sweden and ‡Ugelstad Laboratory,
Department of Chemical Engineering, NO 7491, Trondheim, Norway
˚
˚
Received June 29, 2009. Revised Manuscript Received December 3, 2009
The acid/base properties, critical micelle concentrations (cmcs), and pH-dependent solubility of five synthetic
tetraacids have been studied at several ionic strengths (20-600 mM NaCl) and in the pH range of 1.5-11 using high
precision potentiometric titrations, tensiometer measurements, and UV spectroscopy, respectively. The molecular
weight of the tetraacids ranged between 478 and 983 g/mol. The potentiometric titration data was evaluated in terms of
thermodynamic equilibrium models, developed in the light of relevant solubility data, Langmuir monolayer compres-
sions and cmc of the different tetraacids. The results indicate that for two of the tetraacids, called BP5 and BP7, two
chemical forms fully dominate the speciation of the monomers; the insoluble fully protonated form, and the soluble fully
deprotonated form. The partly protonated species, only play a very minor role in the speciation of these tetraacids. For
the other tetraacids the results are more complicated; for the smallest tetraacid, called BP1, all species seem to play
important roles, and for the most hydrophobic, BP10, the formation of micelles and aggregates severely complicates the
evaluation of the speciation. For the tetraacid BP3 one of the partly deprotonated forms seems to be important, thus
confirming the structure to properties relationship. In spite of the complicated micelle formation chemistry, and
although not actually measured, the acid/base properties for the monomers of BP10 were interpreted by means of
surface charge densities of the micellar aggregates. The modeling indicates an increase of the aggregation number of the
micelle upon acidification, a result of formation of mixed micelles incorporating the fully protonated and deprotonated
species. An intrinsic pKa of 5.4 for BP5 was used to model the monomer pKa of BP10, and corresponded well with a
monolayer acidity constant pKsa of 5.5 obtained from surface collapse pressures of Langmuir monolayers as a function
of pH.
Introduction
present paper, several synthetic tetraacids have been examined by
means of high precision potentiometric titrations in order to
understand, and to provide input to modeling of the aquatic acid/
base chemistry of these interesting, hydrophobic, compounds.
It is commonly accepted that the activity of Hþ at an interface
deviates from the aqueous bulk phase activity.5,6 The surface of a
micelle can be considered as an interface where the proton activity
will change due to the surface charge potential of the micelle,
compared to the bulk phase. Studied using indicator dyes solu-
bilized in micelles of varying surface charge have shown that the
apparent pKa values for the dyes change compared to their pKa
values in aqueous solutions. Depending on the surface charge of
the micelle, the apparent pKa value will increase or decrease. As an
example, the solubilization of hydroxycoumarin in micelles of
SDS, a negatively charged surfactant, increased the pKa value
with nearly 3.5 pH units, compared to solubilization in pure
water.5 The acid/base propertiesof dyeswill alsoreflect the acidity
constant for the surfactants comprising the micelles. Conse-
quently, titration of a negatively charged surfactant in a micelle
may give completely different apparent pKa values than its
intrinsic acidity constant at infinite dilution. Tetrameric acids,
shown to be surface and interfacially active, can thus be expected
to form micelles in aqueous solutions and apparent pKa values
different from their intrinsic pKa values. As results, potentio-
metric pH-titration of a series of tetrameric acids should be
interpreted in terms of their cmc values and if they are present
in monomeric or micellar state.
The problems associated with the occurrence of naphthenic
tetraacids in crude oil, e.g. depositions on equipment, have been
widely described in literature.1-3 The basic chemistry of these
compounds is therefore interesting and economically important,
and provides background for investigations. However, the diver-
sity of these naturally occurring molecules, hence forth referred to
as indigenous tetraacids or ARN acids, and a complicated and
time-consuming purification process, severely complicates the
studies of their chemical and physical properties. Therefore,
investigations of the chemical properties of these compounds
have, to a significant extent, lately been done by means of model
compounds. In previous papers, the synthesis3 and some of the
physiochemical properties4 of several such synthetic model com-
pounds were studied. An important finding in the latter paper was
that, in particular, one of the synthesized compounds, called
BP10, showed chemical and physical properties very similar to
those of the indigenous tetraacids. The problems associated with
ARN acids are mostly predominant at the w/o interface, and
understanding the behavior of such compounds in aquatic
environment would be advantageous. So far, no thermodynamic
modeling studies have been presented with respect to them. In the
*Corresponding author. Ola Sundman, Department of Chemistry, Umea
˚
University, SE-90187 Umea
˚
.
€
(1) Brandal, O.; Hanneseth, A.-M. D.; Hemmingsen, P. V.; Sjoblom, J.; Kim, S.;
e
Rodgers, R. P.; Marshall, A. G. J. Dispersion. Sci. Technol. 2006, 27, 295–305.
(2) Lutnaes, B. F.; Brandal, O.; Sjoblom, J.; Krane, J. Org. Biomol. Chem. 2006,
4, 616–620.
€
˚
rd, E. L.; Sjoblom, J. J. Dispersion Sci. Technol. 2008, 29, 1–9.
(4) Nordgard, E. L.; Magnusson, H.; Hanneseth, A.-M. D.; Sjoblom, J. Colloids
(3) Nordga
€
˚
(5) Fernandez, M. S.; Fromherz, P. J. Phys. Chem. 1977, 81, 1755–1761.
(6) Mukerjee, P.; Banerjee, K. J. Phys. Chem. 1964, 68, 3567–3574.
Surf. A 2009, 340, 99–108.
Langmuir 2010, 26(3), 1619–1629
Published on Web 01/06/2010
DOI: 10.1021/la902326y 1619

Article
Sundman et al.
Figure 1. Structures of the investigated tetraacids, BP10, BP7, BP5, BP3 and BP1. Also the structure of one of the most dominating
indigenous naphthenic C80 tetraacids is shown.
The previously studied model tetraacid, BP10, have been reported
to have a very low cmc.4 Since several studies of titration of micelles
report a significant difference between the apparent acid/base
properties of monomers and micelles,7-9 it was decided that several
synthetic compounds, with different critical micelle concentrations
(cmcs), would be studied. For this task, the model substances, BP10,
BP7, BP5, BP3, and BP1 were chosen, where BP stands for
benzophenone, which is the aromatic core, and 10, 7, 5, 3, and
1 stands for the number of carbons in the attached alkyl chain. These
five tetraacids (TAs) are shown in Figure 1, where also a proposed
structure of the most abundant ingenious tetraacid2 is shown.
Potentiometric titrations are preferable performed at constant
ionic strengths >100 mM. Such conditions were therefore tested
in the present work. However, since the cmcs of the tetrameric
acids (TAs) vary with ionic strength, and since the acid/base
properties of carboxylic acids are known to depend upon the ionic
strength, it was decided that the study should involve also a
variation of the ionic strength conditions. Finally, to provide an
even more solid foundation for the evaluation of the titration
data, several solubility experiments, film compressions and cmc
studies were performed.
NORMAPUR), calibrated using dried (80 ꢀC overnight) Trizma
Base (SIGMA). The dilute NaOH solutions were prepared from a
concentrated NaOH solution, filtered through a glass funnel to
remove possible contaminants of solid carbonates, and calibrated
via titration with dilute HCl(aq) of known concentration.
Synthesis. All the synthesis work was preformed at the Ugel-
stad Laboratory, NTNU, Trondheim, Norway. BP10 was synthe-
sized as described by Nordga
˚
rd et al.,3 and the BP5 and BP7
material was synthesized according to a similar procedure, where
the methyl 11-bromoundecanoate had been replaced with ethyl
6-bromohexanoate and methyl 8-bromooctanoate, respectively.
BP1 and BP3 were synthesized via their methyl and ethyl esters,
respectively, by reacting 2,20,4,40-tetrahydroxybenzophenone with
methyl bromoacetate or ethyl 4-bromobutyrate, in K2CO3/acetone
at 65 ꢀC for 16-72 h. BP1 and BP3 were obtained by further
hydrolysis with NaOH in MeOH:H2O (4:1, 8 h) and acidification
with HCl (2M) which precipitated the undissociated TA as a white
solid. For use in cmc measurements, the tetrasodium (Na4L(s))
salts of BP5-10 were obtained from the acid form by suspension in
a small amount of methanol where the TAs dissolved upon heating.
The acids where converted to their tetrasodium salts by adding
saturated NaOH in methanol until pH > 12. The salts precipitated
upon addition of acetonitrile and were filtered and dried.
Cmc Measurements. The cmcs of the TAs were measured using
surface tension measurements with a Sigma70 Tensiometer (KSV
Instruments, Finland) using a De Nuoy ring probe, at a pH of
approximately 11. The TA was added as the Na4L salt, dissolved in
1 mM NaOH with NaCl added to the appropriate [Naþ]. These
measurements were all performed at Ugelstad Laboratory, NTNU,
Trondheim, Norway. An automatic titration setup was utilized with
two titrators (Schott, Titronic); one for addition and one for removal
of solution to change the concentration. Both titrators were compu-
ter-controlled and 15 points per decade were obtained with a
maximal deviation at equilibrium of 0.02 mN/m during three
consecutive measurements. The solution was stirred with a magnetic
stirrer for 300 s and left to stand for another 300 s before measuring
the next point.
Experimental Section
Chemicals. The chemicals for organic synthesis were all
purchased from Sigma-Aldrich, and used without further purifi-
cation. For the potentiometric titrations all aqueous solutions
were prepared from deionized (Milli-Q 185 Plus) and boiled
(to remove all gas) water. Dried (180 ꢀC overnight) NaCl(s)
(AnalaR NORMAPUR) was used to prepare all ionic media.
The HCl solutions were prepared from dilute HCl (AnalaR
(7) Kanicky, J. R.; Shah, D. O. Langmuir 2003, 19, 2034–2038.
(8) Goldsipe, A.; Blankschtein, D. Langmuir 2006, 22, 9894–9904.
(9) Dupont-Leclercq, L.; Giroux, S.; Henry, B.; Rubini, P. Langmuir 2007, 23,
10463–10470.
1620 DOI: 10.1021/la902326y
Langmuir 2010, 26(3), 1619–1629

Sundman et al.
Article
It is commonly known that the cmc decreases at increasing
salinity and the reason is that addition of salt reduces the
electrostatic repulsion between the head groups in the micelle
and lowers the energy needed for micellization. Consequently, the
cmc will appear at an earlier concentration. A relationship
between cmc and salinity has been empirically described by the
Corrin-Harkins equation10-12
Table 1. General Equilibrium Matrix Used in the Simulations and
Optimizations of the Simulation Models for the Tetraacidsa
log β
Hþ
L4-
phase
Hþ
0
0
1
0
1
2
3
4
-1
0
1
1
1
1
1
0
soluble
soluble
soluble
soluble
soluble
solid
L4-
HL3-
H2L2-
H3L-
H4L (s)
OH-
opt
opt
opt
opt
-pKw
log cmc ¼ -a þ b log½Naþꢀtot
ð1Þ
To visualize the data a plot of log(cmc) vs log [Naþ]tot was
constructed.
Potentiometric Titrations. All titrations were preformed at
soluble
a The β-values indicated by “opt” were optimized for each of the
models evaluated.
the Department of Chemistry, Umea University, on the high-
˚
filled with 0.6 M NaCl and calibrated using a solution of known
[Hþ] and by using eq 2. Following this, the solid material was
removed by centrifugation and the remaining clear solutions were
filtered through 0.20 μm filters before analysis. The amount of
remaining TA(aq) in the samples were then finally analyzed by UV
measurements (Shimadazu UV-2100) at 322 nm.
precision potentiometric titration equipment, developed at loca-
tion, and used in several previous papers.13-16 All titrations were
performed in an oil bath thermostated at 25 ( 0.1 ꢀC and in a
thermostated room at 25 ( 1 ꢀC. Before and during the titrations
all solutions were protected from CO2 contamination using inert
N2-gas which was washed through 10% H2SO4(aq) and 10%
NaOH(aq). The cell used for the determination of [Hþ] was
Langmuir Monolayer Compressions. Surface pressure-area
isotherms were recorded using a KSV Langmuir trough (Helsinki,
Finland) of effective film area 364 ꢁ 75 mm at ambient temperature.
The trough was made of Teflon and the barriers of Delrin. The
trough and barriers were thoroughly cleaned with acetone or
ethanol, tap water and ultrapure water before use, and by aspirating
the surface with a Pasteur pipet connected to a vacuum-aspirator.
The surface was assumed pure when the surface pressure of
subphase only did not exceed 0.5 mN/m upon full compression.
Isotherms were recorded at different pH of the aqueous subphase;
pH 2.2 adjusted with 1 M HCl, pH 5.6-5.8 using ultrapure water
with a resistivity of at least 18.2 Ω and pH 7 and 8 buffers from
Merck. When necessary, additional NaCl was added in order
to obtain equal Naþ concentrations for all buffers. BP10 was
dissolved in spectrophotometric grade chloroform at a concen-
tration of 0.3 mM. An amount of 40 μL was spread on the subphase
with a 50 μL Hamilton syringe. The solvent was allowed to
evaporate over 5 min before compression at a barrier speed of
5 mm/min.
-Ag/AgCl|ionic media (NaCl) solution/suspension in NaCl(aq)|glass
electrodeþ
The potential (E, in mV) of this cell is given by eq 2:
RT ln 10
ð2Þ
E ¼ E0 þ
log½Hþꢀ þ Ej
F
where E0 isanapparatus constant, determinedin the initial part
of each titration by reacting a dilute HCl-solution with OH-, in 4
to10steps, until theacidwasnearlyneutralized.R ineq2 isthe gas
constant, T is the absolute temperature and F is the Faraday
constant. Ej is a function of [Hþ], the two junction potentials
jac and jalk, and the ionic product of water, Kw, according to eq 3:
-1
Ej ¼ jac½Hþꢀ þ jalkKw½Hþꢀ
In eq 3, relevant values for Kw, jac, and jalk were adopted from
ð3Þ
the literature.17 The equilibrium condition was set to a drift in
E less than 0.2 mV/h.
Premicellar Equilibrium Modeling. When no micelles are
formed, the tetraacids can be treated as an ideal tetrameric acid.
To model the acid/base properties of these ideal tetrameric acids,
the following principle was applied; the system was limited to two
components; Hþ and the fully deprotonated tetraacid studied,
L4-. Only monomers were allowed to form, and the insoluble
tetraacid was considered to be a solid product. All other mono-
mers were considered to be soluble aquatic species, cf. Table 1.
A general equation describing the formation of each of these
species of a tetraacid can then be written as
After calibration, the tetraacids were added in solid form to a
total concentration of approximately 1 mM. Dilute HCl(aq) and
NaOH(aq) solutions with corrected ionic strength (NaCl), were
used in the acidic direction and the alkaline direction, respectively.
Solubility Experiments. The aqueous solubility of the tetra-
acids, as a function of pH, and at constant ionic strength of
600 mM Na(Cl) was analyzed via batch experiments at the
Department of Chemistry, Umea University. The 10 mM solu-
˚
tions of the tetraacids were prepared by dissolving a weighed
sample of the tetraacid in an above stoichiometric amount of
dilute NaOH(aq) in a volumetric flask, and adding solid NaCl to
the appropriate total sodium concentration. The suspended solid
material were then dissolved using an ultrasonic bath (Bandelin,
Sonorex Digitech). Different amounts of dilute HCl(aq) were
added to the samples, which were left on an end-overend rotation
test tube holder for 24 h or more. Following this equilibration, the
pH of the samples were carefully measured using a combination pH
electrode (Orion Ross 8103SC, Thermo Electron Corporation),
pHþ þ L4 - h HpLp -4
ð4Þ
where p is the stoichiometric number of Hþ for each molecule
of TA.
The general thermodynamic formation constant, β0 for a
particular species may then be written as
p,1
fHpLp -4
g
½HpLp -4
ꢀ
γH Lp-4
p
β0
p,1
¼
¼
ð5Þ
p
p
p
fHþg fL4 -
g
½Hþꢀ ½L4 -
ꢀ
þ
ðγ Þ γL4 -
H
where γ represents the activity coefficients as described by eqs 6
and 7.
(10) Corrin, M. L.; Harkins, W. D. J. Am. Chem. Soc. 1947, 69, 683–688.
(11) Jacquier, J. C.; Desbene, P. L. J. Chromatogr. A 1996, 743, 307–314.
ꢀ
(12) Matsuoka, K.; Suzuki, M.; Honda, C.; Endo, K.; Moroi, Y. Chem. Phys.
fAg ¼ ½AꢀγA
ð6Þ
Lipids 2006, 139, 1–10.
€
€
(13) Sjoberg, S.; Hagglund, Y.; Nordin, A.; Ingri, N. Mar. Chem. 1983, 13, 35–
!
pffiffi
44.
pffiffi
I
€
log γ ¼ -0:511 z2
-0:2
I
ð7Þ
(14) Ohman, L.-O.; Sjoeberg, S. Acta Chem. Scand. 1981, A35, 201–12.
pffiffi
3
3
(15) Laine, J.; Løvgren, L.; Stenius, P.; Sjøberg, S. Colloids Surf., A 1994, 88,
1 þ
I
277–87.
(16) Sundman, O.; Persson, P.; Ohman, L.-O. J. Colloid Interface Sci. 2008, 328,
248–256.
€
In eq 7, z denotes the charge of the ion considered, and I is the
ionic strength. The activity of solid phases is either 1, if it exists,
or 0, if it does not exist.
€
€
(17) Sjoberg, S.; Hagglund, Y.; Nordin, A.; Ingri, N. Mar. Chem. 1983, 13,
35–44.
Langmuir 2010, 26(3), 1619–1629
DOI: 10.1021/la902326y 1621

Article
Ifa constantionicstrength and temperatureisapplied, and thus
constant activity factors can be assumed, the following expres-
sion, eq 8, valid only at a specific ionic strength and temperature
can be applied for the soluble species.
Sundman et al.
a configurational variable, h.20 Consequently, the electrostatic
potential in each sub volume associated with a micelle can be
described by the Poisson-Boltzmann equation.20,21 Assuming
that the micelle can be described as a charged sphere, the
Poisson-Boltzmann equation can be expressed as follows:
½HpLp -4
ꢀ
ꢀ
ꢁ
N
d2Φ 2dΦ
F
zie
kT
X
βp
,
¼
ð8Þ
1
p
½Hþꢀ ½L4 -
ꢀ
þ
-
ziCi¥ exp
-
Φ
ð11Þ
dr2
r dr εεo
i ¼1
The formation constants, βp,1, can then be used to model the
acid/base properties of the TA at the specified conditions. Also,
the acid/base model can be used to simulate the speciation and
solubility of the TAs.
These activity-to-concentration calculations are performed
applying the WinSGW software, and the resulting equilibrium
model is an activity-scale equilibrium model. By using the prin-
ciples described above, the general equilibrium matrix presented
in Table 1, was considered the simplest model suitable to simulate
the systems.
For the modeling of the acid/base properties and optimization
ofthe log β valuesWinSGW18,19 was also applied. To optimize the
log β values, rough estimations were tested in the simulation
software until a decent fit of the theoretical titration curve to the
potentiometric titration data was achieved. When a decent model
was achieved, the values were mathematically optimized to
describe the data. In the program, the following relation is
utilized:
In eq 11, r is the radial direction, Φ denotes the electrostatic
potential, ε represents the relative static dielectric constant, εo is
the permittivity of free space, F denotes the Faraday constant,
N represents the number of charged species in the liquid phase
and zi (i = 1, 2, ..., N) is the charge number of ionic species i, Ci¥
(i = 1, 2, ..., N) denotes the concentration of species i at a distance
far away from the charged surface where dΦ/dr = Φ = 0 and
P
i=1,N zi Ci¥ = 0, e represents the charge of an electron, k is the
Boltzmann constant, and T denotes the absolute temperature.
Assuming that the protonated and dissociated groups are
uniformly distributed across the surface of the spherical model
micelle, and that the solution is electrically neutral at the bound-
ary of the volume element, the boundary conditions of eq 11 can
be expressed as follows:
dΦ
dr
σs
1 Remnag
εεo 4πR2m
at r ¼ Rm,
¼ -
¼ -
ð12Þ
εεo
X
2
dΦ
dr
SSQ ¼
ðlog½Hþꢀexp -log½Hþꢀcalc
Þ
ð9Þ
at r ¼ Rv,
¼ 0
ð13Þ
In this equation -log[Hþ]exp is the measured -log[Hþ], i.e., in
principle the pH, in the experimental point from the potentio-
metrictitration, and -log[Hþ]calc is the respective calculated value
from the ideal model. The log β values in the models are optimized
by iteration until the smallest sum of error squares possible is
achieved. Each system, at the specified ionic strength, is then
optimized to the data.
In eqs 12 and 13, Rm denotes the radius of the model spherical
micelle, Rv is the radius of the volume of solution associated with
each micelle (for all practical purposes involved in this work,
Rv can be taken as any arbitrarily large value such that dΦ/dr =
Φ = 0 at r = Rv due to the fact that the micelles can be considered
to be near infinite dilution at the cmc), σs is the surface charge
density of the model spherical micelle, R represents the degree of
dissociation of the acidic surface groups, nag is the aggregation
number of the micelle, and m denotes the total number of acidic
groups on each surfactant in the micelle. For a given degree of
dissociation, eqs 11-13 are solved numerically by the method of
orthogonal collocation onspectralelements22 toobtainthe spatial
profile of the electrostatic potential, Φ.
Z Plots. The data from the potentiometric titrations are
illustrated inthe form ofZ-plots, whereZ is defined as the average
number of protons adsorbed to each molecule of acid, (TA). This
definition is shown in eq 10.
½COO-ꢀ H -h þ Kwh-1
Z ¼ -
¼
ð10Þ
½TAꢀ
B
Once the functional form of the electrostatic potential, Φ, is
know, thetitration curveofapolyelectrolyte titrated with HClcan
be determined from the following equation:20
where H and B are the total concentration (mol/dm3)
of protons and tetraacid, respectively. h is the concentration
(mol/dm3) of free protons, [Hþ], and Kw is the autoionization
product of water at the ionic strength in question.
R
Δμ0
2:3kT
Φjr ¼R σs dA
pKa
¼
þ
þ pKa0
ð14Þ
m
If Z is plotted as a function of -log[Hþ], the degree of
deprotonation over the whole pH-interval can easily be over-
viewed. Although not entirely correct, apparent pKa values, for
each of the species, i.e., H4L(s), H3L-, H2L2-, and HL3- can be
evaluated as Z = -0.5, -1.5, -2.5, and -3.5 respectively.
Postmicellar Modeling. For the experimental conditions
described in this work, the concentration of the aggregate
(micelle) is taken to be very dilute, such that the interaction
between micellar aggregates can be neglected and the concentra-
tion of the micelle only determines the size of the electroneutral
volume associated with each micelle. Additionally, it is assumed
that the electrostatic free energy of the charged micelles does not
vary significantly through various important configurations and
the configurational entropy of the micelle as well as the average
local concentration of acidic groups should depend only on
2:3mRkT
In eq 14, Δμ0 represents the difference of the chemical potential
of the neutralized and un-neutralized acidic groups, and pKa0
denotes theintrinsicacidconstantwhich isdefinedasthe apparent
pKa at the limit of zero degree of dissociation (pK0a = limRf0
pKa(R)).23 The space charge density, Fcd, in the ionic solution
surrounding the spherical model micelle is given by eq 15.
ꢀ
ꢁ
N
X
zie
kT
Fcd
¼
ziCi¥F exp
-
Φ
ð15Þ
i ¼1
Equations14and 15can beemployedalongwith the solutionto
eqs 11-13 in order to determine the appropriate value of the
€
€
(21) Gunnarsson, G.; Jonsson, B.; Wennerstrom, H. J. Phys. Chem. 1980, 84,
3114–21.
(22) Grimes, B. A. Ph.D. Thesis, University of Missouri-Rolla: Rolla, MO,
2002.
(23) Sonnefeld, J.; Vogelsberger, W.; Rudakoff, G. Chem. Phys. Lett. 1991, 176,
309–14.
(18) Eriksson, G. Anal. Chim. Acta 1979, 112, 375–383.
(19) Karlsson, M.; Lindgren, J. 2000; p The WINSGW software was developed at
Department of Inorganic Chemistry, University of Umeå. See http://www.dagger.mine.
nu/MAJO/WinSGW_eng.htm for more details.
(20) Marcus, R. A. J. Chem. Phys. 1955, 23, 1057–68.
1622 DOI: 10.1021/la902326y
Langmuir 2010, 26(3), 1619–1629

Sundman et al.
Article
Figure 2. (a)Surface pressure-areaisotherms of LangmuirmonolayersofBP10atdifferent aqueouspHwith[Naþ] =0.1 M, and(b) surface
collapsepressure versusthe finalaqueouspH. Thecollapsepressures aretakenwhere thehorizontalbreaksoccur andthe lines havebeenfitted
2
to eq 16 with the following data for BP10: ω = 300 A , T = 298 K, πZH = 11.7, πZNa = 17, and pKsa= 5.5.
˚
aggregation number, nag, required to fit an experimental titration
curve. In this work, this is carried out for BP10.
Results
Monolayer Dissociation Constant. A convenient way to
obtain an estimate of the dissociation of a surface active naphthe-
nic acid when present at an interface is by use of the Langmuir
trough technique. A monolayer of the surfactant to be studied is
spread on an aqueous phase, and the film is compressed by two
movingbarriers. The surface pressure, π, is recorded and when the
pressure suddenly drops or remains constant after a steady
increase, the film has most likely collapsed. Havre et al. have
shown that it is possible to obtain a monolayer acidity constant,
Ksa, for a naphthenic model monoacid, 5β(H)-cholanic acid, by
correlating the surface collapse pressure to the aqueous pH under
formation of sodium naphthenates.24 Assuming the same partial
molecular area of the acid and the soap, (ωZH = ωZNa = ω), the
collapse pressure of the monolayer, πc,m, can be found in eq 16,
Figure 3. Solubility of the studied TAs as a function of -log [Hþ]
at 600 mM Na(Cl).The symbols represent data recorded for (2)
BP10, (1) BP7, ([) F BP5, (f) BP3, and (ꢁ) BP1. The lines
represent the solubility as predicted by the acid/base models
presented in the Supporting Information, for BP7 (solid line)
BP5 (dashed-dotted line), BP3 (dashed line) BP1 (dotted line)
respectively. The total concentration of the respective tetraacid
was 3.33 mM in all experiments and simulations.
2
!
À
Á
kT
ω
½Hþꢀ
-π
ω
c
,
ZH
4
kT
πc ¼ -
ln
e
,
m
Kas þ ½Hþꢀ
!
#
À
Á
Kas
Kas þ ½Hþꢀ
-π
ω
c,
ZNa
of the pressure increase may be a better approach in this case.
kT
þ
e
ð16Þ
2
˚
For BP10 these areas were found to be located around 300 A at
pH < 7, and this value has been used in the fitting process.
The dissociation data for BP10 resembles what was found for
5β(H)-cholanic acid, i.e., pKsa= 5.65.24 However, the methodol-
ogy of surface collapse pressure will not take into account the
aqueous bulk phase behavior of surface active compounds, such
as the influence of the solubility of all species and micellization. A
thermodynamic evaluation through the whole sequence, from
dissolution of solid material to monomers in solution, possible via
micellization, is necessary to understand the acid/base chemistry
of these fairly complex molecules. The next sections will conse-
quently deal with the acid/base behavior from a bulk phase
perspective including factors like solubility and micellization. It
will later be shown however that the monolayer acidity constant
obtained for BP10 is comparable to the intrinsic acidity constant
for BP5, a less hydrophobic tetraacid.
It was investigated if this relationship could also be used to find
apparent monolayer acidity constant for BP10. Figure 2a shows
the isotherms obtained for BP10 as a function of pH and [Na]þ =
0.1 M. Surface collapse pressures obtained from Figure 2a
together with a plot of the fitted curve according to eq 16 is given
in Figure 2b.
The fitting of the theoretical curve to the experimental values
gives pKsa = 5.5 for BP10. The highest collapse point was found
at pH 8 due to enhanced water solubility at higher pH and,
assuming that the collapse point at pH 8 is close to the maximum
collapse pressure, the data fits quite well to eq 16. However, the
equation is dependent on similar partial molecular area of acid
and soap. This can normally be found from where the increase in
pressure sets in, but for BP10 this was difficult to obtain. There-
fore, the limiting area, obtained by extrapolation of the linear part
Solubility of Synthetic Tetraacids. The solubility of the
tetraacids was studied in 600 mM NaCl ionic medium. This
medium was chosen since the problems caused by the indi-
genous tetraacids often occur in quite saline waters, i.e., seawater.
€
(24) Havre, T. E.; Ese, M.-H.; Sjoblom, J.; Blokhus, A. M. Colloid Polym. Sci.
2002, 280, 647–652.
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Sundman et al.
Figure 4. Surface tension versus concentration obtained by tensiometer measurements at pH ≈ 11. The data for (a) BP10, (b) BP7, and
(c) BP5 at different ionic strengths. The cmc values have been obtained by visual consideration where the slope intersects the constant surface
tension value after cmc. (d) Corrin-Harkins plot of log cmc vs log [Na]tot for BP10, BP7, and BP5. The solid lines were fitted with linear
regression and the parameters are given in Table 2.
Table 2. Cmc Data Obtained for BP10-5 as a Function of NaCl Added and Parameters in the Corrin-Harkins Equation Fitted to the Experimental
Cmc Data, at pH ≈ 11, as a Function of [Naþ]tot
cmc [M]
compound
20 mM NaCl
100 mM NaCl
200 mM NaCl
600 mM NaCl
a
b
R2
BP10
BP7
2.2 ꢁ 10-5
5.1 ꢁ 10-6
1.5 ꢁ 10-3
2.0 ꢁ 10-2
3.8 ꢁ 10-6
7.0 ꢁ 10-4
1.6 ꢁ 10-2
5.9 ꢁ 10-7
1.4 ꢁ 10-4
8.4 ꢁ 10-3
6.33
4.07
2.21
-1.03
-1.23
-0.69
0.9559
0.9905
1.0000
4.1 ꢁ 10-3
BP5a
a The cmc of BP5 at 20 mM NaCl was too high to be measured with the available amount of compound.
The results, shown in Figure 3, show that the solubility of the
different TAs, as a function of -log [Hþ], show similar type of
behavior, but at different -log[Hþ]. Interesting to see in Figure 3
is that the acid/base models evaluated from the potentiometric
titrations given later, and presented in the Supporting Informa-
tion, accurately predict the solubility.
also given in Table 2. The cmcs of BP1 and BP3 were found to be
too high to be interesting. From the data the cmc can be estimated
at a given salinity. These results confirm that the alkyl chain
length have a very significant impact on the cmc of the tetraacids,
as can be expected.
From parts a and d of Figure 4, it was concluded that the cmc
for BP10 precluded titrations of BP10 monomers in any Na(Cl)
medium. Titrating on a micellar system will complicate the
thermodynamic behavior and interpretation of the system. The
use of the other tetraacids was necessary for studying the acid/
base properties of tetraacid monomers. It was thus considered a
good approach to use the titration data for BP5 and BP7 to
understand the properties of monomeric BP10.
A rapid increase in solubility with -log[Hþ] can, for all
tetraacids, be seen in a narrow -log[Hþ] interval. This interval
is shifted upward with the length of the alkyl chain, and this goes
in good agreement with all the data in this study. It is likely that
the solubility of BP10 does follow the same simple pattern as the
solubility of BP1-BP7, i.e., as the sum of deprotonated molecules,
but we have not been able to model this solubility.
Critical Micelle Concentrations. Critical micelle concen-
trations of the different acids were studied at different ionic
strengths. The results, according to the tensiometer technique
used, are presented in Figure 4a-d and Table 2. The parameters
in eq 1 fitted to the experimental data for BP10, BP7, and BP5 are
Acid/Base Properties of TA Monomers. To understand the
acid/base chemistry of the tetraacids, it is important to differ-
entiate between the acid/base properties of the TA monomers
from that of aggregated TA. The cmc of BP10 was early on
determined to be too low to allow for this. The cmcs of BP5, BP3,
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Sundman et al.
Article
Figure 7. Z-plot for BP1. The symbols represent data recorded at
(9) 20 mM, (f)100 mM, and ([) 600 mM ionic strengths,
respectively. The dotted, dashed and solid lines represent the
respective models, as presented in the Supporting Information.
Figure 5. Z-plot for BP5. The symbols represent data recorded at
(9) 20 mM, (2) H 50 mM, (f)100 mM, and (() 600 mM Na(Cl)
ionic strengths, respectively. The dotted, dashed-dotted, dashed
and solid lines represent the respective acid/base models, as pre-
sented in the Supporting Information.
Figure 8. Z-plot for BP7. The symbols represent data recorded at
(9) 20 mM ionic strength. The dotted line represents the evaluated
model, as presented in the Supporting Information.
Figure 6. Z-plot for BP3. The symbols represent data recorded at
(9) 20 mM, (f)100 mM, and ([) 600 mM ionic strengths,
respectively. The dotted, dashed and solid lines represent the
respective acid/base models, as presented in the Supporting
Information.
presented in Figure 6 illustrates this; also here the deprotonation
gives rise to a steep curve, indicative on a solid in equilibrium with
the deprotonated forms. Nevertheless, a difference in the shape of
the titration curves can be seen at Z j -2, and this is believed to
originate from a relatively stable HL3- ion, illustrated in Figure 9b.
For BP1, quite different properties were observed; when the data
in Figure 5 - 6 are compared with the respective data in Figure 7 the
most obvious differences are the slope of the titration curve and the
position in -log [Hþ] of the deprotonation (“pKa values”). The
steep curves in Figures 4 and 5 are due to the existence of a solid
phase, in equilibrium with the fully deprotonated aqueous species.
The solid phase of BP1 is not stable -log[Hþ] above 2, and
therefore other species must be responsible for the buffering effect
above this -log[Hþ] value. The interpretation of the data is that
stepwise release of protons, making all intermediate species im-
portant, describing the acid/base properties of this molecule.
As a result, the distribution diagram, cf. Figure 9, for BP1 is quite
different from that of the other TAs. The equilibrium model predicts
that all species give an important contribution to the speciation of
BP1, while for BP5 only the solid form, H4L(s), and the fully
deprotonated form, L4-, contribute significantly to the speciation.
Presented in Figure 8 are the data recorded, and the optimized
acid/base model for BP7 at 20 mM Na(Cl) (below the cmc). In this
and BP1, however, were high enough to enable potentiometric
titrations of the monomers at all ionic strengths of interest.
The cmc of BP7 was found to be ∼4 mM in 20 mM Na(Cl) and
∼0.14 mM in 600 mM Na(Cl). By using a total concentration of
∼1 mM, it was thus possible to titrate BP7 both above and below
the cmc depending on the salinity.
In Figure 5 all the potentiometric titration data recorded for
BP5 is shown. From this figure it can be seen that the acid/base
properties of BP5 show a dependence of ionic strength (20-
600 mM Na(Cl)), but that the shape of the titration curve is
unchanged. The steep and constant slope of the titration curve is
indicative of a solid phase in equilibrium with only one soluble
species. The slope itself can not predict the identity of this species,
but combined with the observation that the experimental points
and the theoretical model levels out at Z ≈-4, the soluble species
can be identified as the fully deprotonated L4- species. As an
example, the proposed speciation, in 600 mM Na(Cl), as a
function of -log [Hþ] is shown in Figure 9a and illustrates this.
For the tetraacid with slightly shorter alkyl chains, BP3, similar,
but not exactly the same, properties could be observed. The data
Langmuir 2010, 26(3), 1619–1629
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Sundman et al.
Figure 9. Distribution diagram of (a) BP5, (b) BP3, (c) BP1, and (d) BP7. The lines are constructed according to the equilibrium models
evaluated at 600 mM NaCl medium. The model for BP7 is extrapolated from 20 mM, using eqs 7 and 8. The equilibrium calculations were
preformed with a total tetraacid concentration of 1 mM.
figure also a model extrapolated, using eqs 7 and 8, to 600 mM
(where the acid/base properties of the monomers could not
be measured) is presented. The shape of the titration curve
of BP7 below the cmc is similar to BP5, but with a higher
apparent pKa.
The model lines in Figures 5-9 are the result of optimized
equilibrium models. The formation constants and the log β values
used can be found in the Supporting Information.
Titration of TAs above the Cmc. In contrast to the titration
curves shown in Figures 5-8, all the results from titrations of
BP10, and the titrations of BP7 at 600 mM, where the total
concentration is above the cmcs, are divided into two different
type of curves; virgin titration curves, i.e., curves resulting from
the first dissolution of the TA, and precipitation titration curves.
The latter is resulting from titration of dissolved tetraacid at high
-log [Hþ] in the acidic direction. The “splitting” phenomenon,
the hysteresis, is visualized in Figure 10, where data from such
Figure 10. Z-plot for BP10 and BP7 above their respective cmcs.
Comparison between titration data recorded for (9) dissolution
titrations are shown.
The dissolution curves showhow the solid consumes OH- ions,
and (b) precipitation of BP10 and (2) dissolution and (ꢁ) pre-
cipitationofBP7The totalconcentrationofBP10 was 0.5-1.5mM
and the data were recorded at 600 mM Na(Cl) ionic medium. The
data for BP7 was also recorded in 600 mM Na(Cl) and the total
concentration of BP7 was 0.5-1 mM.
while the precipitation curves show the stepwise precipitation of
solid material from micelles uponloweringthe pH. The latterkind
of data was used for the evaluations of micelle formation
In Figure 11, dissolution data for BP10 recorded at several
ionic strengths are illustrated simultaneously, these curves illus-
trate how the solid material buffer the -log[Hþ] as it consumes
OH- ions and establishes an equilibrium with micelles. The
corresponding illustration of the precipitation of BP10 is shown
in Figure 12.
Since the cmc of BP10 is very low, almost all deprotonated
tetraacid will immediately form micelles. A hypothesis is that
these micelles, by means of the hydrophobic effect, interact with
the remaining solid, thereby stabilizing the solid, fully protonated,
1626 DOI: 10.1021/la902326y
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Sundman et al.
Article
Figure 13. Comparison of the experimentally measured back-
titration curve for BP10 to the theoretical prediction made with
eq 14. The values of nag and Δμ0 were adjusted to fit the experi-
mental data.
Figure 11. Z-plot for the dissolution of BP10, above the cmc
(approx 1 mM). The symbols represent data recorded at ([)
20 mM, (2) 50 mM, (f) 100 mM, and (9) 600 mM ionic strengths,
respectively.
of the micelle radius, Rm, was taken to be the length of an
extended 10 carbon chain. The total length from the carboxyl
carbon to the phenolic oxygen was estimated by geometrical
2
˚
considerations and using the length of a C-C bond of 1.54 A .
Furthermore, the value of the intrinsic acid constant, pK0a, was
estimated to be approximately 5.4 (from the titration curve of
monomeric BP5). This corresponds well to the monolayer acididy
constant (pKsa = 5.5) for BP10 obtained from eq 16 and
Figure 2b. The value of Δμ0 that was found to best represent
the experimental data was 8.0. It can be clearly seen from
Figure 13 that a micelle having an aggregation number, nag, with
a value of 60 provides a good fit with the experimental data from
Z = -4 to Z = -2. Once the value of Z ≈ -2, the value of the
aggregation number, nag, that provides a good fit to the experi-
mental data is about 102. This sudden change in the value of the
aggregation number, nag, seems to indicate that once the value of
Z ≈ -2, the acid may be starting to precipitate or form mixed
micelles composed of protonated and deprotonated species. The
total electrostatic repulsion in the micelle would decrease and
promote micellar growth. It should be noted that due to the small
radius of the micelle, the behavior of eq 14 is extremely sensitive to
changes in the value of nag and hence changes in the micelle radius.
In addition, it is assumed that the TA is formed as a spherical
micelle with all four chains directed to the micellar surface. An
arrangement with two of the four chains sticking into the micelle
interior would yield a higher micellar radius, affecting the
modeling. The possibility that the molecules aggregate into a
cylindrical shaped micelle with the aromatic cores lying flat on
each other can neither be excluded. Consequently, the results in
Figure 13 should be interpreted cautiously and suggests that the
micelle size, shape and value of nag should be experimentally
determined in future work.
Figure 12. Z-plot for the precipitation of BP10, above the cmc
(approx 1 mM). The symbols represent data recorded at ([)
20 mM, (2) 50 mM, (f)100 mM, and (9) 600 mM ionic strength,
respectively.
species. At Z r 3, it can be hypothesized that most of the material
is found in micelles form, i.e., micelles in acid/base equilibrium
with other micelles. In this region the titration curve is not only
reversible but also the data from four different ionic strengths
(20-600 mM) are overlapping, cf. Figures 11 and 12.
When the system then is back-titrated the deprotonated form is
stabilized due to formation of mixed micelles incorporating all
species including solid material.
It is clear that there is a significant difference between the acid/
base properties of BP10 at 20 and 600 mM Na(Cl). The data
recorded in 50 and 100 mM Na(Cl) are, however, overlapping and
almost indistinguishable. The phenomenon responsible for this is,
however, not possible to model using any simple equilibrium
model, and is explored using micelle formation theory.
Modeling of Experimental Data by Formation of
Micelles. By utilizing the methodology described in the Experi-
mental Section, an approach using surface charge densities and
aggregation numbers to model the acid/base behavior in a
micellar system has been carried out for BP10. In Figure 13, the
theoretical prediction of the precipitation-titration curve for BP10
in a 20 mM NaCl solution at 20 ꢀC is plotted versus the
experimental data. As a first level of approximation, the value
Discussion
Of the studied materials, it is clear that BP10, previously shown
to have physiochemical properties similar to indigenous tetra-
acids, is the most interesting model tetraacid from an industrial
perspective. It is nonetheless interesting to study a series of similar
compounds in order to fully unfold the complexity of the
chemistry involved. A major problem with potentiometric titra-
tions on BP10 is its low cmc, and the resulting low monomer
concentration, even at quite low ionic strengths. This must be
overcome, and the use of a series of tetraacids, with a range of
cmcs over several orders of magnitude, was considered a good
Langmuir 2010, 26(3), 1619–1629
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Sundman et al.
These TAs both act closely to the behavior of four different
separate ideal monoprotic carboxylic acids. For the intermediate,
BP3, the species with 3 negativecharges attracts the last remaining
Hþ enough to stabilize this species over a significant -log[Hþ]
interval, which also is what could be expected from interpolating
the knowledge from BP1 and BP5. Furthermore, the deprotona-
tion of the solid BP1 occurs at a very much lower -log [Hþ] than
what is the case of BP3-10. This is likely partly due to the
inductive effect from the neighboring oxygen atoms, pulling
the negative charge toward the ketone group of the benzo-
phenone core and stabilizing the negatively charged anions. Also,
possible formation of intermolecular hydrogen bonding, for the
partly deprotonated species, could contribute to this signifi-
cantly increased acidity. A similar situation is seen if benzene
1,2-dioxydiacetic acid, with a first pKa ≈ 2.6,25,26 is compared to
phenoxyacetic acid, pKa ≈ 3.0.27,28
Dupont-Leclercq et al.9 illustrated that the apparent pKa of a
fatty acid depend highly on the amount of micelles available. This
could thus possible explain the peculiar slope of the Z-plots for the
micellar systems. And, since the tendency to form micelles are
greatly dependent on the ionic medium present, cf. Figures 11 and
12, the ionic strength-dependent shape of the Z-plots could
possibly also be explained, because added salt reduces the
electrostatic repulsion and the surface charge of the micelle.
Moreover, the formation of micelles is most likely responsible
for solubilizingthepartly deprotonatedspecies, thus also effecting
the shape of the titration curve.
When first discovering this unique class of surfactants which is
responsible for calcium naphthenate deposits, potentiometric
titrations of ARN in an organic solvent mixture was carried out
together with commercially available mono- and diacids.29 The
titration showed in fact only one dissociation step comparable to
the observed aqueous behavior of BP7 and BP5 and the mono-
Figure 14. Apparent pKa of BP3-10 at 0.5-1 mM concentration,
defined as the -log[Hþ] at Z = -2, in 20-600 mM NaCl as a
function of the number of carbons in the alkyl chain length. The
“missing” data point for BP7 at 100 mM was due to insufficient
material for further potentiometric titrations. The lines illustrate
the linear relationship.
approach to this problem. The five different molecules investi-
gated showed upon quite different apparent acid/base properties,
and it is from the data obvious that the formation of micelles
significantly affects the apparent acid/base properties of the
molecules. Thus, it is clear that the most interesting area is the
micelle formation area, and that the focus of further studies
shouldbeinthe areawhere micelles are formed. It is also clear that
an attempt to understand the physiochemical properties with
respect to the formation of micelles, e.g., aggregation numbers
and type of micelles, would be of interest, and therefore such
calculations were preformed.
layer measurement for BP10. From oilfield observations it is
ARN
estimated that pKa
∼ 6,30 again resembling BP5. However,
An apparent pKa value can be defined as the -log[Hþ] at which
50% of the carboxylic groups are deprotonated. Thus, an obvious
observation is that the apparent pKa values of the carboxylate
groups increase with the alkyl chain length. This is in good
agreement with previous observations.7,9 In the work by Kanicky
and Shah,7 it is observed that the apparent pKa for long chain
carboxylic acids, are significantly higher than for short chain
carboxylic acids. A linear relationship between the chain length of
the carboxylic acids and the apparent pKa was postulated, and the
authors assigned this effect to premicellar aggregation and to the
formation of micelles. As is seen in Figure 14, this is also the case
for the tetraacids studied here. Although the linear relationship is
not obvious for the data collected at the lowest ionic strenght, it is
very clear at higher ionic strength. This behavior is therefore
confirmed also for a series of tetraacids, thus providing new
knowledge. Since the TAs with longer alkyl chains tend to have a
much stronger driving force for the aggregation, Figure 14 clearly
illustrates the significant influence from the formation of aggre-
gates on the apparent acid/base properties.
the concentration of ARN in crude oil is in most cases around 1
ppm (∼10-6 mol/dm3) which may very possible be below the cmc
of ARN. And, since the cmc varies with the salinity of the
produced water, which may differ from oilfield to oilfield,
this can give varying estimations of pKaARN, and, to our know-
ledge, limited data of direct determinations of the pKaARN can be
found in the literature. Furthermore, to be able to perform a
potentiometric titration, the concentration of the acid studied has
to be around or above approximately 0.5 mM in order to
maintain sufficient sensitivity for the cell. Thus, performing
potentiometric titration experiments using isolated indigenous
tetraacids, around or above the cmc, would yield apparent pKa
values significantly deviating from field observations. Combined,
this could be the reason for limited data for this compound class.
A higher apparent pKa value can be seen for BP7 compared to
BP5. As can be seen in Figure 4, the increase of equilibrium
surface tension below the cmc for these tetraacids is quite slow.
This might indicate premicellar aggregation when approaching
the cmc. This aggregation affects the deprotonation of BP7 in
The difference in speciation behavior between monomers of
BP7, BP5, BP3, and BP1, shown in Figure 9, is probably due to
several factors. For BP1, all species tend to give an important
contribution to the speciation, while for BP7 and BP5 only the
solid, H4L(s), and the fully deprotonated species, L4-, contribute
significantly. For BP3 the HL3-(aq) ion also seems important. The
observed phenomenon is likely due to the alkyl chain length and
molecular size; the short inner molecular distances in BP1 makes
the charge from one of the carboxylic acids affects the depro-
tonation of the others, which is not the case with BP7 and BP5.
(25) Suzuki, K.; Hattori, T.; Yamasaki, K. J. Inorg. Nucl. Chem. 1968, 30, 161–
66.
(26) Hasegawa, Y.; Choppin, G. R. Langmuir 1977, 16, 2931–2934.
(27) Suzuki, K.; Yamasaki, K. J. Inorg. Nucl. Chem. 1962, 24, 1093–1103.
(28) Pettit, L. D.; Royston, A.; Sherring, C.; Whewell, R. J. J. Chem. Soc. B.
1968, 588–90.
(29) Baugh, T. D.; Grande, K. V.; Mediaas, H.; Vinstad, J. E.; Wolf, N. O. SPE/
IADC 93011. The discovery of high molecular weight naphthenic acids (ARN
acids) responsible for calcium naphthenate deposits. In SPE 7th International
Symposium on Oilfield Scale; Aberdeen, 2004; SPE, Richardson.
(30) Brocart, B.; Bourrel, M.; Hurtevent, C.; Volle, J.-L.; Escoffier, B. J.
Dispersion. Sci. Technol. 2007, 28, 331–337.
1628 DOI: 10.1021/la902326y
Langmuir 2010, 26(3), 1619–1629

Sundman et al.
Article
this concentration range, increasing the apparent pKa also below
the cmc.
apparent pKa values at 20 mM Na(Cl) show a linear dependence
on the alkyl chain length, which must be correlated to premicellar
aggregation. Finally, we show that an accurate estimation of the
acid/base properties of the BP10 micelles could be achieved via
modeling based on the electrostatic potential of the micelles
and assumed geometries of the micelle. The modeling indicated
an apparent increase of the micelle aggregation number when
protonating the tetraacid in the micelle. This increase could be
due to solubilization of the protonated species in the micellar
environment.
Synthesis of ARN itself or shorter-chained ARN molecules is a
very difficult task. In addition, the purity ofthe isolatedARN acid
may vary.4 Both these problems are however circumvented
by using the model tetraacids in this study, and for BP5 it is a
good correlation between the potentiometric titrations and the
apparent pKaARN from field observations.
Conclusions
It was early concluded that the cmc of BP10 did not allow
potentiometric titrations of the monomers of this tetraacid.
Therefore, several other tetraacids were developed and investi-
gated. The model tetraacids with long alkyl chain length basically
show only one protonation step, while the model tetraacids with
short alkyl chain length show more than one protonation step.
The strange titration curves for the tetraacids, when titrated
above their cmcs, strongly indicate the importance of micelle
formation for the acid/base properties of large tetraacids, that
estimated intrinsic pKa values in some cases must be used for
modeling, and that model compounds sometimes can be used for
this purpose. For BP10 it is interesting to see that the pKa
estimated from Langmuir monolayer films coincide with the
intrinsic pKa for the BP5 model compound, which was shown
to be present in a monomeric state. Moreover we show that the
Acknowledgment. The authors thank the JIP consortium,
consisting of AkzoNobel, Baker Petrolite, BP, Champion
Technologies, Chevron, Clariant Oil Services, ConcoPhillips,
Shell Global Soulutions, Statoil ASA, Talisman Energy and
Total, for major financial support of the present work. Further
the Research Council of Norway is acknowledged for partly
funding the work preformed at NTNU. Professors Staffan
€
€
Sjoberg and Lars-Olof Ohman (Umea University) are very
˚
greatly acknowledged for valuable scientific discussions through-
out the whole course of the present work.
Supporting Information Available: Table giving the thermo-
dynamic data used in Figure 7. This material is available free of
charge via the Internet at http://pubs.acs.org.
Langmuir 2010, 26(3), 1619–1629
DOI: 10.1021/la902326y 1629