Title Page
 Table of Contents
 Adsorption of carboxylic acids...
 Adsorption of chondroitin sulfate...
 Adsorption of polypeptides
 Reaction of peptides and carbohydrates...
 Adsorption of collagen
 Biographical sketch

Title: Calorimetric measurements of the adsorption of collagen and other organics onto oxide surfaces
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Permanent Link: http://ufdc.ufl.edu/UF00089830/00001
 Material Information
Title: Calorimetric measurements of the adsorption of collagen and other organics onto oxide surfaces
Series Title: Calorimetric measurements of the adsorption of collagen and other organics onto oxide surfaces
Physical Description: Book
Language: English
Creator: Buscemi, Paul John
Publisher: Paul J. Buscemi
Publication Date: 1978
 Record Information
Bibliographic ID: UF00089830
Volume ID: VID00001
Source Institution: University of Florida
Holding Location: University of Florida
Rights Management: All rights reserved by the source institution and holding location.
Resource Identifier: alephbibnum - 000208356
oclc - 04109699

Table of Contents
    Title Page
        Page i
    Table of Contents
        Page ii
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        Page v
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    Adsorption of carboxylic acids on alumina and hydroxyapatite
        Page 41
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    Adsorption of chondroitin sulfate and other carbohydrates
        Page 61
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    Adsorption of polypeptides
        Page 79
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    Reaction of peptides and carbohydrates in solution
        Page 93
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    Adsorption of collagen
        Page 117
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    Biographical sketch
        Page 164
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Full Text









ACKNOWLEDGEMENTS .. . . . . . . .ii

ABSTRACT . . . . . ... . . v


I INTRODUCTION . .. . . . . 1

Use of Protein Models . .. . . 2
Calorimetry as an Analytic Tool .. .. 4
Protein Adsorption . . .. . . 6
Surfaces. ............ . 10
Forces of Adsorption .. . . . . 15
Thermodynamics of Adsorption. . . ... 17
Heterogeneous Adsorption . . . ... 24
Experimental. . . . . . . ... 33


Introduction. . . . . . . ... 41
Experimental . . . . . . ... 43
Results . . . . . . . . 44
Discussion. . . . . ... . . 51
Conclusions . . . . . .. .... 59


Introduction. . . . . . . ... 61
Experimental . . . . . . ... 62
Results . . . . . . . . 62
Discussion. . . . . . ... 70
Conclusions . . . . . .. . . 77


Introduction. . . . . . . ... 79
Experimental. . . . . . . ... 81
Results . . . . . . . .. 81
Discussion. . . . . . ... 86
Conclusions . ... . . . . ... 92




IN SOLUTION . . . . . .

Introduction. . . .
Experimental. . . .
Results . . . .
Discussion . ..
Conclusions .. ...


Results . .
Discussion. .
Conclusions .


Experimental. .
Results . .
Discussion. .
Conclusions .





. . . . . 93
. . . . 95
. . . . . 97
. . . . 108
. . . . . 115

. . . . 117

. . . . 117
. . . .. . 117
. . . . 117
. . . . 132
. . . . 138

. . . . . . 141

. . . . . . . 141
S . . . . . . 142
. . . . . . . 142
. . . . . . 148
. . . . . . . 150

. . 152



Abstract of Dissertation Presented to the Graduate Council
of the University of Florida in Partial Fulfillment of the
Requirements for'the Degree of Doctor of Philosophy




Paul J. Buscemi

March, 1978

Chairman; R. E. Loehman
Major Department: Materials Science and Engineering

The present work is the result of the application of

solution microcalorimetry to the problem of determining the

energies of adsorption of organic molecules onto ceramic

surfaces. The systems studied were chosen to model the

attachment of collagen to ceramics and to provide some expla-

nation for the observed bonding of ceramic implants to bone.

An aqueous solution of an organic molecule such as a

polyamino acid, polysaccharide, or smaller molecules with

similar functional groups was automatically mixed in a micro-

calorimeter with a slurry of a powdered oxide such as A1203,

SiO2, or a special glass composition and the heat evolved or

adsorbed was determined. Calorimetric measurements were

performed on increasing concentrations of reacting organic

molecules for a fixed weight of powder with known surface

area. Plots of the reaction heat, Q, versus the initial con-

centration of the organic, Co, yield thermometric titration

curves which were analyzed to give the enthalpy of the

reaction AH, the free energy change AG, and by difference

the entropy change for the reaction AS.

The systems were studied in order of increasing

structural complexity from simple carboxyls, amines, and

organic sulfates to amino acids, polyamino acids, polysaccha-

rides, and collagen. Changes in the aqueous solutions by

additions of salts or changes in pH and combinations of

organic molecules were also studied.

Results indicate that there are at least two forces

which contribute to the bonding of the collagen and other

organic to oxide surfaces, hydrogen bonding and ionic bond-

ing, the former releasing from 8 to 12 kcal/mole of functional

groups while the latter releases 4 to 6 kcal/mole of

functional groups depending on the relative polarities of

the adsorbing molecules and the surface. There are strong

indications of the denaturation of collagen at some surfaces

at which hydrogen bonding and ionic bonding act cooperatively.



The study of adsorption and interaction of proteins and

other biological molecules onto non-biological surfaces is

important because of the increasing use of prosthetic materials

[1]. It is essential to know how each of these two distinctly

different components will interact in biological media. In

this study, a further understanding of the reaction between

the connective tissue protein, collagen, and various oxide

surfaces is sought.

Two properties of collagen adsorption are of major

concern: how well does the protein adhere to the oxide surface

and does the surface change the structure of the protein?

Protein adhesion relates to the binding of tissue to a prosthetic

material and can be approached by the determination of the

enthalpy AH of the adsorption reaction [2]. Changes induced

by the surface on the protein lead to denaturation. This

increases its vulnerability to enzyme attack and eventual

rejection from the host [3]. This question can be approached

by comparison of the enthalpy of reactions of model systems

with that of reactions which have been found not to be

disruptive to the structure of protein.

Three materials serve as substrates: silica, alumina,

and hydroxyapatite. Each of these materials is well character-

ized and has potential for use as a prosthetic material [4].

The calorimetric measurements made, therefore, are not for

the purpose of further characterizing these materials but to

study their influence on the adsorption of organic molecules.

The values of the thermodynamic parameters (AG, AH, and

AS), determined from the calorimetric measurements, do depend,

however, on the structural features of the surface as well as

those of the adsorbing molecule and on their mutual environment.

A'practical approach for studying complicated systems encountered

in actual application is to study simpler model systems [5],

which provide singular features for observation.

Within a series of model compounds, the structure of

the molecules, the solvent, and the surface can be systematically

varied and correlations can be made between the thermodynamic

data and the variations in experimental conditions. In this

study, extensive use is made of several types of molecular

model compounds including those representing collagen as well

as those representing carbohydrate and other physiologically

relevant organic structures.

Use of Protein Models

Past workers have used molecular model compounds designed

to study collagen. Specifically, the polyamino acids have

been well studied in this way. Poly-l-proline has been used

in conformational studies in CaC12 solution [6] showing that

disordering of the molecule is associated with its increasing

rotational ability. X-ray diffraction studies of poly-l-proline

have demonstrated that the backbone conformations of the

molecule are similar to native collagen [7]. Poly-l-lysine

and poly-l-arginine have been used as models for collagen in

relation to the structure of amorphous ground substance [8,9].

Poly-l-lysine and poly-l-glutamic acid have been used in

conformational studies using differential capacitance techniques

[10]. In other studies, workers have used combinations of

synthetically prepared amino acids to model collagen [11,12].

Smaller molecules representing isolated residues of

the collagen molecule have also been investigated but not as

frequently as the polyamino acids. Dyes containing amino

groups have been shown to selectively adsorb chondroitin sulfate

[13], an important carbohydrate in structural tissue [14].

Surface viscosity measurements using amines, amides, and

carboxylic acids as model proteins have been studied in relation

to bilayer film formation [15] in membrane studies.

In biological studies of proteins other than collagen,

molecular models have been widely studied. Enthalpies of

aqueous solution have been determined calorimetrically for

amines and carboxylic acids [16] as part of a quantitative

description of biological systems. Heat capacity measurements

[17] on several amino acid solutions have been made to help

explain protein structure. The binding of short amino acid

chains [18] in subunit studies of immunoglobulin has also been

investigated by calorimetry. Differential scanning calorimetry

has been used to study conformation changes of many polypeptides

used as models for collagen [19,20].

Molecular models in non-biological studies have also

been used. The adsorption of molecules containing the same

functional groups which proteins possess, amines [21], sulfates

[22], and organic acids [23], have been examined by various

methods. Calorimetry has not been extensively used for this

purpose. In general, the use of molecular models in biological

and non-biological systems appears widespread for the determina-

tion of various properties.

Calorimetry As An Analytic Tool

Calorimetry has long been used to measure the enthalpy

of adsorption (heat of wetting) of various liquids onto dry

oxide surfaces. Such measurements are made by immersing a

clean, dry solid powder into a liquid. The heat change is

measured in a calorimeter. For our purposes, the most

relevant liquid used in previous studies was water. Heats of

wetting of silica [24-26], of alumina [27-29], and of hydroxy-

apatite [29] have each been measured. The results of these

works show several consistent features. First, there are

differences in the heats of wetting with variation in the out-

gassing pressure and with temperature of evacuation, indicating

surface heterogeneities. Also there are differences in heats

of wetting with variation in particle size. Finally, the heats

of adsorption of water range from -10 to -20 kcal/mole of water

adsorbed and are attributed to hydrogen bonding of the water

to the surface [30].


Calorimetry has also been used in a similar manner

to measure the heat of adsorption of a second component from

an aqueous solvent [31]. The heat of adsorption of sodium

dodecyl sulfonate (SDS) was found by subtracting the heat of

wetting of alumina in pure solvent from that found with SDS

present. In this case the initial adsorption attraction was

attributed to coulombic forces between the surface and ion

and was calculated to be -12 kcal/mole.

Results for similar experiments have been confirmed by

other methods [32,33] using water and other polar and ionic

liquids as adsorbents. For example, the differential heat of

adsorption for the adsorption of octadecyl alcohol on alumina

from a benzene solution, calculated from the temperature

dependence of the adsorption coefficient was found to be -8.6

kcal/mole while that found directly from calorimetric measure-

ments was -8.68 kcal/mole [34].

The heat of adsorption of water on quartz determined

by adsorption measurements at several temperatures was found

to be between -11 and -14 kcal/mole of water adsorbed [25].

There were differences found when the quartz was ground and

exposed to water vapor prior to drying and evacuation. This

was attributed to the formation of an amorphous layer of silica

on the surface. These differences disappeared when the rough-

ened surfaces were annealed at 7000C. The heats of wetting

agree well with those found from calorimetric measurements [24].

Calorimetry has been extensively used in biochemical

applications [35]. However, there are few calorimetric data

on the adsorption of proteins onto oxide surfaces and there

is apparently none for the adsorption of collagen. The

calorimetric data most relevant to adsorption primarily

involve such globular serum proteins as serum albumin [36].

The remainder of this section will therefore be devoted to

previous studies of protein adsorption use with particular

emphasis on oxide substrates.

Protein Adsorption

The demonstration of molecular attachment of cell

proteins on foreign substances has been accomplished by various

methods. Multiple internal reflection spectroscopy has

been used to measure protein interaction using germanium [37]

as a substrate. A KRS-5 prism pressed against protein on a

hydroxyapatite substrate allowed protein-hydroxyapatite inter-

action to be studied [38]. Although energy calculations were

not carried out in these studies, the changes in adsorption

frequencies indicate chemical interaction with the surface of

the substrate.

Film compression studies [39] using collagen, gelatin,

and poly-l-alanine with silica gel showed adsorption hysterisis

indicating an irreversible process. The maximum interaction

of the silica gel and collagen occurred at pH 5.2. The iso-

electric point,where there is charge neutralization of the

protein, is 5.5. There was also interaction between alanine,

which has no ionic side groups, and the gel. The interaction

appeared to be of the same type as that of collagen. The

primary binding force was assumed to be hydrogen bonding [40].

Adsorption of bovine serum albumin (BSA) on hydro-

philic silica [41] exhibited a maximum surface density at

pH 5.5. The isoelectric point (IP) of this protein is 4.9.

The surface of the silica is negatively charged at this pH.

Desorption occurred readily at pH's away from the IP indicating,

as suggested by the author, that binding was due to hydrophobic

interaction. Other studies [42,43] showed that even after

extensive washing with water and EDTA that not all of the BSA

adsorbed onto pyrex glass could be recovered. Maximum

adsorption was near the IP of the protein and the free energy

change was estimated to be -2.5 kcal/mole of protein. The

enthalpy was not calculated.

Serum globulins have been shown to be preferentially

adsorbed by silicic acid [44] and silica [45] and by other

minerals [46]. Maximum adsorption took place at the isoelectric

point on these surfaces as well as on calcium phosphate gel

[47]. In none of these studies were determinations of the

enthalpy of the various adsorption reactions made.

The adsorption of albumin, fibrinogin, and globulin on

polyethylene has been determined by internal reflection spectro-

scopy [48]. The adsorption isotherms followed a Langmuir isotherm,

a common finding in which the quantity of solute adsorbed, X,

at the equilibrium concentration, C, is given by X = aC/(l + bC)

where a and.b are constants. The adsorption was assumed to be

due to hydrogen bonding because of the shift in the amide I

(C=O stretching) band at 1640 cm-1

Calorimetric measurements of the adsorption of human

serum albumin on negatively charged polystyrene (PS) [49]

were shown to be pH dependent. Maximum adsorption of the

protein occurred near the IP (4.9) where the enthalpy of the

reaction was positive. At pH values removed from the IP the

reaction was exothermic. It was suggested that at pH values

away from the IP the conformation of the adsorbed protein

changes for energetic reasons. Denaturation of the protein

is not surprising since it is known that the internal bonding

of serum albumin is weak [50]. The' enthalpy of the adsorption

reaction varied between zero and 8 kcal/mole as the surface

charge varied between -1.0 mpCcm-2 and -7.5 mpCcm-2. The

most negative enthalpy values were recorded for pH 3.8 and 9

and were near -8.4 kcal/mole.

Enthalpy values for the adsorption of albumin, gamma

globulin, and fibrinogin, were calculated from adsorption

measurements at several temperatures [51]. The results indicated

that the adsorption took place in two distinct ways. Both

types were apparently Langmuir and took place on separate

membrane sites. One type of adsorption was easily reversible

with a heat of adsorption in the neighborhood of -10 kcal/mole.

The other type of adsorption reaction was hydrophobic, endothermic,

and with heats of adsorption in the range of 5 to 20 kcal-


In another study [52], the forces involved in the

adsorption reactions between several globular proteins and

glass surfaces were determined to be primarily ionic amine-

silanol bonding and hydrogen bonding. Two rates of adsorption

were noted. The first appeared to be related to the number

of amines present on the surface of the protein. The second

was slower and seemed to be dependent on the molecular weight

of the protein. Hydrogen bonding was suspected since the

proteins could not be completely washed from the surface with


Ionic bonding of ribonuclease to glass was indicated

to be strong [53] since very little of the protein could be

removed by rinsing in several solvents. No enthalpy determina-

tions were made.' There was a decrease in adsorption with

an increase in ionic strength.

Further review of the literature reveals that the

various adsorption studies cannot be readily compared due to

the large number of experimental variables and to the random

manner in which they are controlled in each experiment. A

few common features in the study of protein adsorption do

emerge, however.

There is usually more than one type of interaction

present for any particular system and one of these is usually

hydrogen bonding. The observed enthalpy values are in the

range of -10 to +10 kcal/mole of protein. Finally, maximum

adsorption density appears to take place near the IP of the

protein. There are many exceptions to these general results,


The effect of changes in ionic strength on adsorption

is also not well understood [54]. Dissolved salts disrupt

hydrogen bonds which proteins depend on for conformational

stability [55]. Changes in the ionic strength of the solvent

will also have effects on the adsorbing surface. For example,

phosphate, a common component in buffers, will change the surface

charge of alumina [56]. Generalizations are difficult to

make about the action of specific ions on adsorption unless the

specific system understudy is clearly defined.


Silica, alumina, and hydroxyapatite, the three materials

used in this study, are oxides. Hydroxyapatite is sparingly

soluble at neutral pH whereas silica and alumina are virtually

insoluble [57]. All are hydrophilic and each displays a

surface charge which varies with pH.

The surface charge results from exposed surface atoms

attempting to complete their coordination of nearest neighbors

[58]. Exposed cations do this by pulling an OH" ion or H20 from

solution and anions by attracting a proton from the aqueous

phase. The result is adsorbed H+ or OH- ions which assume their

respective charges on the surface.

Any other ion which can pass between the solid and

liquid phases may also help to establish the surface charge.

Such ions are called potential determining ions. Thus,
OH and H are potential determining ions for eachof the
OH and H are potential determining ions for each of the

+2 3
three surfaces used here. In addition Ca and P04- are

potential determining ions for hydroxyapatite. Certain

ions added as impurities may alter the surface charge such as

aluminum ions [59] or cobalt [60] on silica or phosphate on

alumina [61]. The wide variety of buffering systems used in

biological studies involving adsorption can thus lead to

differing results for proteins even if the same material is

used as a substrate.

The surface potential can be altered by a change in

pH. For each of the three oxide substrates in this study, there

exists a pH at which the surface charge is zero. This pH is

called the point of zero charge (pzc) and is listed in Table

1 [62,63] for the substances used as adsorbents.

As the pH varies from one side of the pzc to the

other, the sign of the surface charge will change as will the

adsorption properties of the protein.

Ions in solution which do not pass through both phases

but are attracted near the surface by electrostatic forces

are called counter ions. They will form a diffuse layer of

ions in solution near the surface and will tend to neutralize

the surface charge. The concentration of the ions generally

decreases exponentially with distance from the surface [64].

The.higher the concentration, the more compact the diffuse

layer will be. The thickness of the diffuse layer ranges from

about 10 A at .1 M solutions to about a few hundred A in .001 M

salt solutions [65].


Table I

PZC of Substances Used as Adsorbents

alumina 9

silica 2-4

hydroxyapatite 7.5

Due to the shape of the diffuse layer, the influence

of dissolved salts on proteins will be greater near a surface

than in bulk solution. Some proteins, because of their

large size, may extend entirely through a double layer. The

effect of dissolved salts on adsorbed proteins would then be

difficult to explain in detail.

The adsorption properties of the substrates are due

as much to adsorbed water as to their intrinsic stucture.

The -surface of silica in aqueous solution has been shown by

infrared spectroscopy to possess three types of surface ions

[36] as represented below:

Si 0-

Si 0-H

Si 0-H2

Unless the adsorbed water and associated ions are driven off

by heating, silicon ions cannot chemically react with organic

molecules arriving from solution. It has been shown that

ammonia will not react with hydroylated silica, although

chemisorption will occur if the silica has been subjected to
a prior vacuum degassing at temperatures in excess of 400 C

[66,67]. Silica treated with ammonium fluoride solution

showed evidence for Si-F bonding instead of silanol [68] but

this reaction occurred after heating the substrate to 4000C

in vacumm.* Trimethylsiloxane can be covalently bonded to

silica by refluxing them in acetone for 24 hours at 500C [69].

This gives some idea of the difficulty of penetrating the

adsorbed layers on the silica surface. It can be seen that

the reactions of aqueods solutions of proteins with silica may

occur with either surface oxygen or hydrogen, depending on

the compositional purity of the surface.

The same effect can be seen with alumina. Steric

acid was adsorbed from CC14 solvent after the alumina had been

evacuated at 8000C for one hour. Without the pretreatment,

steric acid would not covalently bind to the alumina surface

[70]. Methanol has been shown to adsorb on alumina [71] after

successive evacuation and heating at 4000C, heating in oxygen

to rid any hydrocarbons present and then heating again at

10-5 torr at not less than 3500C for 1/2 hour. A methoxide

surface is formed when the clean dry surface is exposed to

methanol vapor. From studies of adsorbed acetylene on alumina

it was concluded that the surface contains electron poor

and electron rich [71] sites (oxide ions, hydroxyl groups, and

aluminum ions) after the sample had been heated to 8000C.

Even at these temperatures not all the hydroxyl groups were

removed from the surface [72]. Hydrogen and hydroxyl ions on

alumina are exposed to the solution interface, yielding a

surface structure similar to that of silica.

Hydroxyapatite is assigned the formula Cal0(PO4)6(0H)2.

In solution the surface undergoes hydrolysis, yielding a surface

having the formula Ca2(HP04) (OH)2 [57]. It is different from

silica and alumina in that, in addition to surface OH- and H

ions, there are also Ca+2 ions which are capable of binding

adsorbing anions [73]. The multiple internal infrared spectra

of several synthetic and naturally occurring calcium phosphates

exposed to organic acids show shifts in the P-0 stretching

frequency [39] which were attributed to hydrogen bonding. Other

workers [30] have shown that H30+ ions are hydrogen bonded to

the calcium. There have been suggestions that there is some

covalent bonding between organic constituents and hydroxyapatite

in bone [74]. Binding of calcium ions by collagen has been

demonstrated by solution analysis [75]. It has not been shown

that collagen can attach to the surface of the hydroxyapatite

without an intervening water molecule or hydroxyl ion.

Forces of Adsorption

The relationship between the enthalpy of a reaction and

the total energy is H = E + PV. Most biological processes occur

in liquids rather than in the gas phase [76]. In this case

the changes in pressure and volume are small. To a good

approximation then, dE ~ dH.

The total energy of adsorption is affected by the type of

interaction between the surface and the molecule. This energy

is comprised of several components. They may be classified as

non-polar, ionic, hydrogen, and covalent bonding [77-79].

Non-polar dispersivee) forces are always present between

molecules. They arise because the time-averaged electron cloud

.interaction between uncharged atoms is attractive [80]. They

are moderately strong, producing energies in the range of 1 to

10 kcal/mole. Hydrophobic bonding is a result of dispersive

forces. This is a consequence of a decrease in displace-

ment of a polar medium when less polar components coalesce,

thus creating a lower energy state.

Ionic or electrostatic attraction can occur between

oxides and proteins both of which are normally charged in

aqueous solutions. Ionic bonding strength is decreased by

an increase in the ionic strength of the solvent because of

shielding, whereas dispersion forces are not affected. Ion

interaction is essentially independent of temperature [80,813.

Hydrogen bonding is partially ionic and partially

covalent [82,83]. It arises from the electrostatic force

acting between hydrogen and a lone-pair of electrons of

nitrogen, oxygen, or fluorine. The small size and the close

approach of the hydrogen atom accounts for the partial (20%)

covalent character [82]. Typical bond strengths are of the

order 1 to 10 kcal/mole. Hydrogen bonds are also weakened

by an increase in ionic strength of the solvent. Completely

covalent bonding rarely occurs in the adsorption phenomena

in which we are interested [62].

The classification of physical or chemical adsorption

is somewhat arbitrary [79,83]. If the adsorption is found

to be readily reversible and has an energy of the same order

of magnitude as the liquefaction of gas, it usually is classified

as physical [77 ]. Osipow states that Van der Waals forces

are responsible for physical adsorption whereas Fuerstenau

also includes coulombic attraction. Chemical adsorption is

irreversible and the magnitude of the energy change is of the

order of chemical reactions [78].

Thermodynamics of Adsorption

Complete reviews on the thermodynamics of adsorption

are given elsewhere [76-78]. In this section, only those

points needed to explain the following data will be presented.

Limited explanation of the ideas of earlier workers would

be in order, however.

Gibbs gave the first rigorous thermodynamic explanation

for why a given material should either adsorb or desorb at a

surface. He was able to predict the functional relationship

between surface tension and surface concentration and the bulk

concentrations of the surface-active solutes. His derivations

assume that substances tend to minimize the free energy of

the surface region by either becoming concentrated or depleted

there. As a result, it has been the free energy which has been

traditionally determined, and surface concentration measurements

are the most common method of doing this. Other methods, such

as the thermometric titration method [84], may be useful for

obtaining thermodynamic data.

The thermometric titration technique is an analytical

method in which the heat effect of a titration reaction is

used to measure the titer of a sample. It is applicable to

reactions of the type

RI + R2 P (1.1)

which entails a heat of reaction Q. 'In equation (1.1) Ri

refers to reactants i and P refers to the product. For single

step reactions an equilibrium constant, K, may be written


K = [P]
[R1] [R2]

where the brackets denote concentrations. Through knowledge

of equation (1.1) and the reaction heat, Q, the enthalpy change

of the reaction and the equilibrium constant may be determined.

The method by which this is done will be detailed for adsorption


For an adsorption reaction equation (1.1) can be written


Su + R So

where R1 in equation 1.1 has been replaced by Su which is an

unoccupied site on a solid surface capable of adsorbing from

solution a reacting component R to produce an occupied site So.

The equilibrium constant is then written

K = [S ,

[Su] [R]

The concentration of occupied sites [Sol is equal to

the number of adsorbed molecules per unit area, Na/A, while

the concentration of unoccupied sites is equal to the total

number of sites N minus the number of occupied sites per unit

area, (N0 N )/A. The equilibrium constant is then

K = N

(No N ) ER]
s a

Dividing;- through by Ns gives
K = N /N
a (1.3)

(1-N /Ns) [R]
a s

K= 6

(1-0) [R]

where 0 = Na/N0 is the fraction of occupied sites. Equation
(1.4) is the Langmuir adsorption isotherm [ ] and will be

used later. Its use requires that each site is occupied by no

more than a single molecule and that no two sites interact.

These conditions are satisfied when the concentration of R is


The concentration of reacting molecules is equal to the

number of moles of R in solution divided by the total volume

Nr/V. If No is the total number of moles of R on the surface

and in solution then N No Na. The equilibrium constant
r r a
can now be written as

K= Na V

(N Na) (N Na)
j r a

The enthalpy change for the adsorption reaction, AH, is

related to the reaction heat Q by

AH = -Q/Na (1.6)

where the minus sign denotes, by convention, that an exothermic

reaction (positive Q) will yield a negative enthalpy change.

Substituting equation (1.6) into (1.5), one obtains

-K = -QAHV (1.7)

(N AH + Q)(AHNo + Q)
s r

To determine K, at least two adsorption experiments are completed

in which Q is measured but a set of values is usually completed

to determine the endpoint of the titration. The original

number of moles of reacting molecules, No, is varied, and No

is held constant by keeping the surface area of the adsorbing

substrate constant. Values for Ql, Q2, N1, N, and sare

recorded where N1 and N are values for Nr. The value of K is
1 2 r
assumed to be nearly constant if N1 is not too different from

N2 Equation (1.7) can then be solved for AH by using the

quadratic equation (1.8)

AH2 Ns (Q1o 2 Q2N) + AH Q1Q2 (N2-N1) +Q1Q2(Q2-Ql)=O

for the two sets of values. The total number of surface sites

can be estimated by dividing the total surface area by the

known cross-sectional area of the adsorbing molecule. This

method is valid as long as the surface concentration of adsorbed

molecules is low and lateral interaction does not occur. Alterna-

tively three sets of values can be used to eliminate N from

equation 1.7. Both methods give results = 10% of each other,.

The volume V is held constant. The value of AH is substituted

into equation (1.7) to find K, and therefore AG by using the


AGo = -RT In K


In order to complete the thermodynamic data (AGO =

AHo TASo) Ho must first be determined or approximated.

The number AH used in equations (1.6) to (1.8) is the total

enthalpy change for the reaction under experimental conditions

while AHo is the standard enthalpy change. Under suitable

conditions, AH can be shown to be = AH so that ASo can be

determined. Those conditions will now be explained.

The chemical potential /i or partial molar free energy

Gi of component i in a chemical reaction is defined by

'i = Gi = ('G/aNi)tp (1.10)
where G is the free energy change for all components in the

reaction. The chemical potential can be expressed as a function

of the activity ai of component i and of the chemical potential

in some reference state

^i = /iref + RTlnai (1.11)

The term RTlnai takes into account the energies of interaction

of component i with other components at a given concentration in

the mixture. The choice of reference state is quite arbitrary

and varies for experimental convenience. Generally, no matter

what reference state is chosen, the activity is expressed as a

function of the mole fraction component i, Xi, and a parameter

known as the activity coefficient fi

ai = xi fi
The activity coefficient approaches 1 in pure solutions for

the solvent while in dilute solutions of component i, fi

approaches a constant which may be greater or lesser than one.

The partial molar enthalpy is found to be

Hi = -RT2 ( lni/ T)pn (1.12)

or in view of equation (1.11)

H-i = -RT2(ln i ref/DT) RT2 ( n ai/DT)pn (1.13)

or = -RT2(OIn iref/T)n RT2 (~in fi/DT)p

Under the constant composition Xi does not vary with temperature.

The first term of equation (1.13) is the enthalpy

change for component i which would occur if the reaction was

held under reference conditions. The standard enthalpies, HO,

of pure compounds are the enthalpies of reaction of building

up those compounds from their elements under standard conditions

(P = 1 atm T = 298K). The chemical elements themselves have

zero standard enthalpies of formation. If the reaction is

carried out under standard conditions, the enthalpy change for

the reference state is equal to H?. Then
o 2
H. = H. RT (~ln f./ T) (1.14)

Knowledge of fi for the systems under study is lacking.

We therefore make the approximation that AHo is significantly

larger than the natural logarithm of the temperature variation

of fi. The validity of this approximation relies upon the non-

interaction of solute molecules. This condition is assumed to

hold for dilute solutions. Therefore

Hi = Hi (1.15)

To relate HO to AH it is noted that AH can be written

as the difference of the sums of the partial molar enthalpies

of products and reactants

AH = ZiHiX + kHkZk (116)

products reactants

where each sum is taken over each of the different components

for products and reactants and X. and Zk are the respective

mole fractions. In view of equation (1.15) the enthalpy

change for the adsorption reaction is

AH = H X + EkHk Zk (1.17)

AH = AH (1.18)

where AHo is the standard enthalpy change for the complete

reaction. Thus, under the rather ideal conditions in which

there is no interaction between solute molecules in solution or

on the surface at 1 atm and 2980K AH may substitute for AH.

The value of AH can be used with AGo to find at least an approximate

value for AS0.

The experimental conditions in this work meet the contraints

of pressure and temperature. The constraint of non-interaction

of solute molecules holds only for dilute solutions. In as

much as enthalpy values tended towards.constant rather than

steadily decreasing values,lateral interaction on the surface

between adsorbed molecules does not appear to have occurred.

Interaction of solute molecules in solution can only be assumed

to be small in the concentration ranges used, typically 10- to
10 M.

Heterogeneous Adsorption

As mentioned earlier, from the standpoint of adsorption

studies, the surface of a substrate is often not uniform.

Heats of adsorption may vary at different positions on the

surface. If the condition is maintained that the different

sites are non-interacting, the adsorption onto a heterogenous

surface may be regarded as simultaneous independent reactions

of the type expressed by equation (1.3)

Sl + R Sol
S + R + So1

S2 + R So2

S + + R Soj

where the subscript j enumerates the different types of sites.

The concentration of a single type of solute molecules, R,

is common to all surface sites. Each adsorption reaction,

according to equation (1.19), would evolve a reaction heat

Qj. The total reaction heat, Qt, would be the sum of Qj for

each reaction on the different types of sites.

Qt Ql + Q2 + .. .Qj (1.20)
If AHi is the enthalpy change per mole of adsorbed molecules

on sites of type.j and Nj is the number of moles adsorbed then

equation (1.20) can be expressed according to equation (1.6) as

-Qt = AH1 N1 + AH2N2 + ....AHjNj (1.21)
In a calorimetric measurement it is the value of Qt

which is measured. If the solution is analyzed to determine

the total number of moles adsorbed, Nt a total apparent

enthalpy is found

AH = -Qt/Nt (1.22)

From equation (1.20) then

AHt = -Qt/Nt = 1/Nt (AH1 N1 + AH2 N2 +...AHj N )


AHt = AH1 Y + AH2 Y2 +....AH. yj (1.23)

where Yj is the fraction of the total number of moles of

molecules bound to sites of type j. Applying equation (1.18)

to each reaction

AHt = EjAH9 yj

where AHO is the weighted sum of the standard enthalpy changes

for all adsorption processes. The standard free energy

change for each reaction expressed by equation (1.19) is given


AG = AH? TAS? (1.25)
3 1 1
Rewriting equation (1.25) for A and substituting into (1.24)
one obtains

AH = E.(AG9 + TAS9) y (1.26)
t i I I
= AC yj + TASjyj


AH = AGo + TASo
t t t

AGo = G. AG y.
L I 3 I



AS = ES? yj (1.28)

express the weighted sums of the standard free energy and

standard entropy changes for the complete adsorption reactions.

For each reaction expressed in equation (1.19) there

is an equilibrium constant Kj which can be written according to

equation (1.4) as

K. = 6. (1.29)

(1-06) [R]

or related to AGQ according to equation (1.9) as

AG? = -RT In K. (1.30)
J 1
Equations (1.29) and (1.30) express the fact that each site

carries on an equilibrium reaction independently of all others.

The fraction 6. = N./N? is the ratio of occupied sires of type

j to the total number of sites of type j.

For any equilibrium concentration [R] of solute molecules

the total fraction of occupied sites of all types, 8t, can be

written as

6 = N /No
t t s

S = X1 61 + X2 0 + ...X. (1.31)

where X. = N9/No is the fraction of sites of type j and is a
] 3 s

If equation (1.30) is substituted into equation (1.27)

one obtains


AGo = -E. (RT In Kj) Yi
t. J


= -RT'ln K' K^ ..... K (1.33)
12 2
Thus, even though each AG? for the individual.reactions is

constant, the overall free energy change will vary as each

fraction y. varies.

The net effect is that if any single type site adsorbs

a large percentage of all molecules adsorbed then

AGt approaches AG? as v 1

Generally, however, there is no simple number AGo which can

be expressed in terms of a single equilibrium constant Kt,

having the form

K = KJ K2 ........KY (1.34)
t 1 2 3
The value of Kt would vary as the fraction of occupied sites


The physical interpretation is that each site contributes

a specific amount of energy to the total energy change. At

very low concentrations only a very few sites react, presumably

those of higher energy. At higher concentrations a greater

number of sites react, but of overall lower energy. The result

is a lowering in the average energy change as the concentration

is increased.

Under the conditions of independently acting sites

evaluation of K would give AGo exactly. However, to do this,
t t
knowledge of each Ki and yi has to be available. In the absence

of such knowledge, the evaluation of K can be approximated by

evaluation of equation (1.29) by replacing 6j by Ot. The number

found from this method, K', could be used for determination of

AGO under suitable conditions.

Those conditions may be determined by estimated

values of Kj for calculation of K' and Kt and using each for

evaluation of AG The value of K' is found from

K' = Ot

(1-et) [R]

where 0t is defined by equation (1.31). The percent errors in

AGt is found from

% error AGt = n K In K' x 100

In K'

The percent error is calculated by setting values for K. and

Xj and directly calculating K' from equation (1.34) and Kt

from equation (1.40). If the percent error is acceptable, AGt

may be calculated.

An example for the case of two different types of

adsorption sites is given. The equations necessary to carry

out this calculation are given below for convenience

61 = K1 [R]/(l + K1 ER])

2 = K2 [R]/(l + K2 [R])

yl = 01No /(O1NO + 0 NO)
11 IN 1 2 2
y = 2N /(62No + 2NO)
Kt = K1 K2

X NO / (NO + NO)

X= N / (No + N
2 2 1 2
ot = X161 + X202
K' = t/(1 + t) [R]

In the example the values of [R] were carried over several

orders of magnitude while the total number of sites NO =
Nl + N2 was kept constant at 10-6 moles. The results are

shown in Figure 1

Two cases are presented. In (a) the fraction Ni/No

0.1 and K = 5000 are held constant, while K1 is given

values of 1000 and 5000. In (b) K1 and K2 are held constant

at 1000 and 2000 respectively while X varies over three

values; 0.002, 0.02, and 0.2.

While K2 is less than five times the value of K1 and

X1 is a good approximation of Kt with the error remaining

within 1%. As X1 approaches lthe error goes to zero. Also,

at very low concentrations, it is assumed that most molecules

would adsorb only onto the highest energy sites so all values

of y (equation 1.34) go to zero except Y and the error again

goes to zero. For oxides the overall fraction of highly

reactive sites is small [85,71]. Moreover, the hydrated

oxide surface will be of lower energy than a perfectly dry

surface, aiding in meeting the condition that K1 not be too

much larger than K2 [60,28]. Under these constraints and using

equations (1.8) and (1.7) with Q = Qt, AHt can be calculated.

For the calculations in the later chapters, the

difference between the values of NO for use in equation (1.8)
is kept small so that variation of AH in that concentration

interval is small. The values of AGo calculated tend to remain

within 20% of the highest to lowest values. This corresponds

S K2 = 5000

Sx = 0.1
10 K',Kt ( K1=50000) ,\ 4

9 (K = 50000)

( ) 8 -

(K1 = 10000)
6 0

5 K', Kt ( K1 = 10000) o

SK2 =2000 4
0 -x --"--" l =0 2 K1 =10000 0
.- ."' 2


(h) 6 0
Kt (.2)

4 Kt (.02)


7 6 5 4 3 2 1

-LOG (R)

Figure 1. Examples for the calculation of K' and Kt
for constant K2 and xl (a), and for constant K2 and K1 (b).

to a range of K. varying by about a factor of seven. Under

these conditions the maximum error in AGt is less than 2%.

The value of AS calculated from

ASo = AGo AH
t t


can be given only simple interpretations. It is known that

values of AS will be between -20 and +20 cal/mole-deg [49,51].

A decrease in entropy is typically explained as a loss of

freedom of solute molecules as they adsorb. Increases in

entropy, generally found in experiments using macromolecules,

are explained as solvent molecules gaining additional freedom

as the largestructuring molecules are removed from solution.

Exceptions to this general rule are present.

The various plots of AG, AH, and AS presented in the

following chapters, in accordance with the previous discussion,

are to be understood as the composite values AGt, AHt, and

ASt. The values for these parameters are closer to single

values of AH?, AG?, and AS? in those regions of the curves

where they tend towards constant values. In these regions the

percent error is generally less than .5%.

The determination of the thermodynamic properties for

the adsorption of collagen on hydroxyapatite is presented as

an example of the calculations made in the following chapters.

The first step in the analytical procedure is to plot Q vs. C

(Figure 2). From this graph two values of Q and Co are chosen

for the sample calculation. In this case values of




c1 c22

I I I i

0,5 1,0 1,5 2.0


Figure 2. The first step in the calculation of the thermo-
dynamic functions is plotting the reaction heat, Q, vs the
original concentration Co. Here, two points are taken from
the measurements of the adsorption of collagen on alumina.

Q1 = 5.0 x 10"3 cal
Q2 = 5.9 x 103 cal

C = 1 x 10-6 M

C = 1.33 x 10-6 M

are arbitrarily chosen. The values of CI and C2 correspond


N = 2 x 10-9 moles

N = 2.66 x 10-9 moles
The number of moles of surface sites, No, in this example
is taken as 2.5 x 10-9. This number was determined from

consideration of the surface area occupied by a collagen mole-

cule, the number of sites as calculated by the program, and

study of the reaction heat curve. Substitution of these

values into equation 1.8 yields -3.3 x 106 cal/mole for AH.

Substitution of AH, Q1, No and V (4 x 10 -3 l)into equation 1.7

gives 4.3 x 105 for K or -7.6 kcal/mole for AGo. This in turn

yields -1.1 x 104 e.u. for AS0.


Calorimetric measurements were made using an LKB model

10700 batch microcalorimeter [86]. The basic calorimetric unit

consists of two identical gold cells situated in an aluminum

heat sink (see Figure 3). Each cell has two compartments

capable of holding 2 and 4 ml of fluid. Mixing of the fluids

in each compartment is accomplished by rotation of the entire

calorimeter. There is no stirring, and after rotation, the full






Figure-3. Schematic of the operation of an LKB model 10700
microcalorimiter. A-calorimeter, B-reaction cell, C-thermo-
pyle, D-heat sink, E-amplifier, F-chart recorder, G- reaction
heat curve. The reaction heat, Q, is proportional to the
area under the curve.

4 ml of fluid are contained in the forward compartment. In

all experiments, 2 ml of fluid were used in each compartment.

Measurement of the heat loss or gain incurred by the

mixing procedure is made through multiple thermopiles located

between the heat sink and cells on two sides of each cell.

The thermopiles are connected in opposition so that the signals

from reactions producing equal amounts of heat are cancelled

electronically. One cell is then arbitrarily chosen as a

reaction cell and the other as a reference cell in which

unwanted heats can be cancelled.

Determination of the reaction heat is made by manual

integration of the voltage vs. time curve produced during the

course of a reaction. The energy is calibrated against a

known heat produced in the reaction cell using a precision

resistor and a known current generated for a specific time


Each reaction in this work is of the type


A BI A + B

where A is a solution of organic molecules and B is a slurry

of powdered oxide used as a substrate. In the mixing operation

several reactions are possible, each contributing to the over-

all heat produced. They are due to 1) wetting of the cell

wall; 2) dilution of the organic molecule; 3) friction of

mixing; 4) chemical reaction. Only the last heat is desired.

The others must be eliminated.

The first three heats are accounted for in various

ways. The cells are first wet with the solvent being used

and emptied prior to filling with reacting components. This

eliminates the heat of wetting of the cell wall. The heat of

dilution of the organic molecules is accounted for in the

reference cell, while the heat of dilution of the oxide powders

and frictional heat are measured separately and subtracted

from the reaction heat.

The heat measured in this way gives a measure of the

reaction A + B C where C is a complex of solid particulates

and adsorbed organic molecules.

Calibration of the calorimeter was carried out as suggested

by the manufacturer and consisted of two procedures. The first

procedure determined the sensitivity of the thermopiles. The

manufacturer listed the sensitivity of the thermopiles as 28.0

and 30.0 microvolts for a constant current of 30 miliamps

through the calibration heaters. The measured values were

consistently within 28.0 L .05 microvolts and 30.0 1 .05 micro-

volts. The second procedure judged the accuracy of the unit's

calibration mechanism. The heat of dilution of a six percent

sucrose solution was measured periodically to be 6.36 .05

kcal/mole. The literature value is 6.36 t .03 kcal/mole.

No literature data could be found for the heats of

adsorption of surfactants from aqueous solution on prewetted

surfaces. However the heat of adsorption of sodium dodecyl

sulfonate (SDS) measured from the heat of immersion of dry

alumina (Linde A) was estimated to be 12 kcal/mole 1 kcal/

mole [23]. Using the method described in this work the

heat of adsorption of SDS on Linde A alumina was found to

be 10.2 kcal/mole of SDS. The discrepancy, -1 kcal/mole,

is thought to be due to the surface being wet prior to contact

with SDS. This would prevent any possibility of SDS coming

into contact with a drier, and presumably, higher energy


A Perkin-Elmer-Hitachi model 139 U.V. Vis spectro-

photometer was used for concentration determinations. Separa-

tion of particulates from supernatent solutions was carried

out by centrifuging or by filtering through micropore glass

filters. In those cases where filtering was used, the filter

was first saturated with a solution of the organic molecule to

be analyzed, then rinsed thoroughly. In all cases a standard

was used which had been put through identical procedures as the


Protein and polypeptide determination was made through

the use of Biuret reagent [87 ] and measuring at 550 millimicrons.

Carbohydrates were determined using a phenol solution and

measuring at 490 millimicrons [88]. Amino acid concentrations

were determined by use of ninhydrin [87]. Carboxylic acids

were titrated with phenothalein as the indicator.

Distilled water, having an initial pH near 7, was used

for preparing solutions. The pH was varied by adding HC1 or

NaOH. Three solutions were used as solvents: a low ionic strength

solution (LISS) in which the only ions present were those added

by pH adjustment, a 0.165 M salt solution, and a buffered solution.

The concentration of NaCI in the low ionic strength solution

did not exceed 0.001 M., The 0.165 M salt (Ringers) solution

consisted of 9 gm/l of NaC1, 0.25 gm/1 of CaC12, and 0.42 gm/1

of KC1 and is known. The buffered solution was a 0.2 M solution

of mono- and di-basic phosphate. The phosphate buffer was

used only in Ringers solution bringing the total molarity to


Except in those instances where solids were not used

at all, 0.1 gm solid particulates were added to the solution.

The solids were placed into suspension by mixing 1.0 gm of solid

powder with about 18 ml of distilled water, adjusting the pH

and then bringing final volume to 20 ml. One hour was allowed

for the pH to equilibrate. With stirring, 2 ml aloquots were

distributed into 10 tubes by pipetting. Using this method 0.1

gm 0.01 gm of solid was delivered to each tube.

The gold calorimeter reaction cells were washed daily

with detergent. The washing procedure included injecting a

solution of the surfactant into each compartment, rotating the

calorimeter and withdrawing the solution by aspiration. The cells

were then continually rinsed with distilled water while being

evacuated. By moving the tip of the aspirator tube from the

top of the cell to the bottom in one compartment, while filling

the other compartment, a good turbulent rinsing reaction

developed. Experience showed that five minutes of such procedure

cleaned the cell. Approximately once a week, the detergent

was left in the cells overnight to permit it to react more


Between daily experiments, the same procedure was used

to clean the cells. The oxide powder slurry was removed for

analysis after mixing, however, and the detergent was not used.

Cleaning with 1 M HC1 or NaOH was found to be necessary only

Organic constituents were weighed out to 1 -'.01 mg

and mixed volumetrically. Adjustment of pH was the same as

with solids. Incremental concentrations were made from standard

batches and measured volumetrically by pipet. Distribution of

organic, and solid, solutions into the reaction cell was made

by syringe. The lowest concentration was always used first.

The materials used as substrates include silica, alumina,

and tricalcium phosphate. The silica [89] was described by the

manufacturer as amorphous. It has. a specific surface area of

0.7 m2/gm and a pzc of 3. The alumina used [90] consisted of

two types: Linde A, a a-alumina of specific surface area 15 m2/gm

and a pzc of near 9 and Linde B, a mixture of y and a alumina

of specific surface area 82 m2/gm and a pzc near 9. The tri-

calcium phosphate, referred to as hydroxyapatite in the text,

consisted of 85 volume per cent hydroxyapatite and had a specific

surface area of 57 m2/gm. Surface measurements were made in

this laboratory using a multi-point B.E.T. nitrogen adsorption

isotherm. Measurements of pzc were also conducted in this

laboratory using a Zeta meterR [91].

All experiments were run at 250 C.

All solids were used without modification. Prior to

weighing, large (approximately 10 gm) batches of solids were

rinsed in distilled water, decanted and evacuated for 24

hours at 10'3 atm. Prepared powders were stored under vacuum

at room temperature.

All organic substances used were stored under refrigeration

prior to use. None was repurified or modified in any way.

Batch solutions were used within one week and were stored at

4C. Particular information on each substance used is given

in the appropriate chapter.

Reaction heat data were plotted against Co and fed into

a statistical analysis program available through the Northeast

Florida Regional Computing Center. The reaction heat, Q, was

held as the independent variable. The program generated an

approximating function which was used to calculate the thermo-

dynamic data. In all chapters, referral to "calculated thermo-

dynamic data" refers to this procedure. The graphs of Q versus

Co presented are original data.





Several of the amino acids which make up collagen

possess charged side groups, principally amines and carboxyl.

The polysaccharide, chondroitin sulfate, often associated with

collagen, also possesses the carboxyl groups as well as the

sulfate group. Because each of these charged species has the

potential to interact with an aqueous oxide surface, knowledge

of the type and strength of the surface reaction is sought.

The thermodynamic and adsorption behavior of molecules

containing the sulfate group have been previously described.

The adsorption of sodium dodecyl sulfate on alumina has been

studied by calorimetric techniques [31] and by solution depletion

techniques [91]. Heats of adsorption for this molecule have

values near -6 kcal/mole at low concentrations when only ionic

forces drive the adsorption reaction. The free energy change

lies near -11 kcal/mole of ions. The adsorption of sulfate

ions onto solid barium sulfate from aqueous solution has been

measured by the thermometric titration technique. The heat of

precipitation was found to be -4.5 kcal/mole .[84].

The adsorption of alkyl ammonium acetate on quartz

has been followed as a-function of temperature and concen-

tration [92-93].It was found that at 25C at neutral pH the

isoteric heat of adsorption was between zero and 2 kcal/mole

of ions. The positive enthalpy was expected because of charge

repulsion of the quartz surface and the negatively charged

ions in solution. The free energy change at 250 remained near

-3.5 kcal/mole. In most cases the hydrocarbon chain of the

molecules caused abrupt changes in the thermodynamic properties

due to lateral interaction of the adsorbing molecules as the

equilibrium concentration increased.

Binding studies of sulfates, citrates, and amino acids

on calcium oxalate and calcium phosphate have been carried

out in relation to bone formation and kidney stone growth

[94,95]. Thermodynamic data are not readily available for these

reactions. Equilibrium constants of magnesium oxalate [94 ],

however, have been shown to be near 4000 which would correspond

to a free energy change near -5 kcal/mole for this precipitation


The purpose of this chapter is to discuss the variation

of the heat of adsorption of simple and polymeric carboxlic

acids on alumina and hydroxyapatite in relation to surface charge,

molecular conformation and ionization and solution pH. Qualita-

tive results indicate that while electrostatic interactions

are required to initiate adsorption, other interactions such

as hydrogen bonding, take place.

The sodium salts of acetic acid, oxalic acid, citric

acid, and polyacrylic acid (PAA), containing one, two, three

and multiple carboxyl groups were selected for study. The

different number of charged groups, different degrees of

ionization and molecular structure of each of these molecules

should provide a sufficient variety of detectable changes in

the calorimetric measurements. Analysis of the various changes

should furnish a clear understanding of the adsorption


The structural: formulas of each of the molecules used

in the work described in this chapter are given below:


Acetic Acid Oxalic Acid

I \ I I I


Citric Acid Repeat Unit of
Polyacrylic Acid


Sodium salts of the carboxlic acids were purchased

from Sigma Scientific, Inc. [96]. Poly(acrylic acid) in a

65% aqueous solution was purchased from Aldrich Chemical

Company [97] and was reported by them to have a molecular weight

of 2000. Linde A alumina, and hydroxyapatite were used as


Titrations of the carboxlic acids were carried out

with HC1 or NaOH. Determination of the amount of acid

adsorbed was based on calibration against known concentrations.

Other methods and procedures were described earlier.


Reaction heats for the adsorption of the three simple

carboxylic acids on alumina are presented in Figure 4. For

each acid, the maximum reaction heat was found to occur when

the solution was near pH 5. The reaction heats at pH 3 were

second and the lowest curves were recorded for pH 7. The

highest heat was recorded for sodium oxalate which also has

the lowest dissociation constant (see Table 2).

The enthalpy change upon adsorption was first found

by determining the amount of each acid adsorbed by titration.

The method was suitable only for pH 5. At pH 3 and 7 poor

precision resulted because of the small amount of each acid

adsorbed. Results are given below:

acetic oxalic citric

-AH (kcal/mole molec.) 4.2 4.8 4.6

(Co = .001M)

Thermodynamic data were calculated for pH 5 using a

smaller concentration range (C = 10'4 to 10-3 M) (see Figure

5). These results show that there are differences in the

enthalpy change for the three acids used. Each curve displays

a tendency towards more negative exothermicc) values at the

lower end of the concentration range. The heat of adsorption

S/ /' 3

20- /5 / 5
20 -
o/ ./ 3

10/ /


4 3 2 1 0


Figure 4. Reaction heats for the adsorption of three
carboxylid acids on alumina.


Table 2

Dissociation Constants
of Carboxylic Acids

Acetic Oxalic Citric

4.75 1.23 3.14
4.19 4.17





25 5

S- 50

o 20 c C 0


15 30
S- 30

o10 20


-4- -. .- .

0.4 0.6


0. 2

lies between -6 and -24 kcal/mole. The free energy changes

are nearly equal at about -5.5 kcal/mole, while the entropy

changes are each negative and lie between -10 and -20 cal/mole-

deg (1 cal/mole'deg = 1 entropy unit or e.u.).

Calculated values for the thermodynamic functions for

the adsorption of the three carboxlyic acids on hydroxyapatite

are presented in Table 3. The values show less variation for

hydroxyapatite than they did for alumina. The enthalpy

change tends to be less negative than for alumina, while the

free energy and entropy changes are about the same.

Polyacrylic acid (PAA) had such a high affinity for

alumina and hydroxyapatite that a cotton-like gel formed at

pH 5 and 7, causing great difficulty in cleaning the gold

reaction cells. Three concentrations were run with alumina,

however, and were repeated several times to improve precision.

Titration determinations of the amount of PAA adsorbed leads

to an enthalpy of 82 cal/gm of PAA for the adsorption reaction.

Taking 72 gm/mole as the molecular weight of the monomer, the

determined enthalpy change is -5.7 kcal/mole of acrylic acid

monomer assuming all acid groups participate in the adsorption


The calculated thermodynamic data for the adsorption

of PAA onto aluminaare presented in Figure 6. The enthalpy

change per residue lies between -7 and -12 kcal/mole, while

the .free energy change is concentration invarient at -5 kcal/

mole. The entropy change is negative as it is with the other

carboxylic acids used in this section.

Table 3

Themodynamic Variables for the Adsorption of
Carboxylic Acids on Hydroxyapatite at pH 5 and 7




(kcal/mole) (kcal/mole) (cal/mole-deg)







15 -30



1 0

5 -- --"- .. ----- --.- i0

0.02 0.04 0.06 0.08

Co, mM

Figure 6. Thermodynamic data for the adsorption of poly-
acrylic acid on alumina at pH 7 in low ionic strength solu-
tion; AG----, AH-- AS.-.-


As molecules are adsorbed from solution, other

molecules will ionize to try to maintain the original concen-

tration. At the same time the alumina or the hydroxyapatite

will act as a buffer to maintain the pH. So long as the pH

is constant, the degree ionization of the solute will remain

the same. This leads to a qualitative explanation for the

reaction heat curves of Figure 4.

The total ionic charge in solution is directly propor-

tional to the degree ionization, a, which is related to the

pH by [98]:

pH = pKo log [ (1 a)/a ] (2.1)

where pK is defined as
pK = -In [H+] [A-] (2.2)
0 [HAJ

This relation holds for single as well as polyelectrolytes


In the absence of any specific interaction between a

solid and electrolytes in solution, the surface potential is

given by [5b]:

S= RT In a/a (2.3)

where z is the valence (including sign) of the potential

ion, F is the Faraday constant, a is the activity of the

potential determining ion in solution and a is the activity

of the potential determining ion at the pzc. As mentioned

in the introduction, the potential determining ion for the

systems under consideration is H+. The surface potential

may then be written from equation (2.3) as

= 2.3 RT (log [H ] log [Ho]) (2.4)
where concentration are substituted for activities and

where [Ho] is the hydrogen ion concentration at the pzc.

The interaction due to the surface potential and the

charge, q, in solution gives rise to an interaction energy,

E [64]

E = o dq (2.5)
q 0
where integration is necessary since the charge concentration

is a differential process. The surface potential affecting

ions in solution will decrease as saturation of the surface by

charged ions is approached. The total interaction energy is

nearly equal to Poq in this model if the concentration of charged

ions in solution is low so that there is little interaction of

the adsorbed ions on the surface. In this case the surface

potential which each ion encounters will be the same. The

energy lost by the ions is transferred to other ions and solvent

in the form of kinetic energy and flows as heat out of the


A plot of o from equation (2.4) and a from equation (2.1)
versus pH for a hypothetical monovalent acid with a pKo near 5

and an oxide with a pcz near 9 is given in Figure 7. At low

pH, a, and thus the charge in solution decreases. At high pH

the surface potential decreases. The highest interaction should

be expected to occur where o and a are not near zero.

2 4 6 8

Figure 7.




Plot of 4o and a for a hypothetical acid.




Figure 8. Values of the reaction heat calculated by use
of equation 2.6 and found by calorimetric experiments.







For example, at pH 5 equation (2.4) gives for a

monovalent electrolyte '(H+)

to = .059 (9 5) = .24 volts
The total charge in solution, using acetic acid as the mono-

valent acid at .001 M is

q = a Ane (2.6)

= 1.78x10-5(6xl023)(2xl0-6)(1.6x10-19)


= .12 coul.

where A is Avogadro's number, n is the total number of mole-

cules in solution, and e is the electric charge. The total

electronic interaction energy is

E = qo q

= (,24) (.12) = .028 j = 6.7 meal

The value of Q found from microcalorimetry is 4.7 meal. The

agreement is fair. Plots for other values of Q and E for other

acids at different pH values are shown in Figure 8 for C =

.001 M. Although the experimental and theoretical results do

not fit well for all cases, the simple model qualitatively

explains the experimental findings, which is what was sought.

Such factors, as treating the ions in solution as other than

point charges, and the potential at the Stern layer are not

taken into account. The most important result is that the

reaction heat and electronic interaction energy decrease as

either o or a decrease. The tentative conclusion develops that

electronic interaction will probably be the essential factor

in enthalpy changes measured for charged adsorbing molecules.

The enthalpy and free energy change values are within the

range expected from the investigations of other workers.

Some idea of the number of active groups of each

molecule actually on the surface can be obtained from the

following analysis. Using the data from Table 4and Table 5

the average number of ionized groups on acetic acid, oxalic

acid, or citric acid can be determined from their dissociation

constants. The numbers are given in Table 4. If we divide

the enthalpy change per molecule by the average number of

ionized groups. AH" is obtained. This assumes that the enthalpy

change for the adsorption of carboxyl groups is about the

same regardless of the molecule being considered. Results are

given in Table 5 for pH 5 and 7 at .001 M.

The values for AH* are fairly consistent in each case,

Considering the enthalpy change of the adsorption for acetate

as an arbitrary baseline, we can argue that if AH* for oxalate

or citrate had been much larger than that for acetate it

would have implied that more groups per molecule were participat-

ing in the surface reaction than expected from the degree of

ionization. Asis,AH* for oxalate is slightly lower and citrate is

slightly higher than the AH* for acetate at both pH values.

Oxalate is the most strongly dissociated of the three acids

implying, perhaps, a slightly stronger reaction with the surface.

It should be realized that a particular group cannot be partially

ionized at a particular instant in time. The invarience in


Table 4

Ionized Groups per Molecule
for Three Carboxylic Acids

pH Acetic Oxalic Citric

3 0.02 1.0 0.42
5 0.63 1.85 1.9
7 1.0 2.0 2.8

Table 5

Enthalpy Change per Mole, AH, and Enthalpy
Change per Mole of Ionized Groups, AH"

pH Substance












AH* is evidence that only ionized groups participate in

adsorption. This supports the supposition that electro-

static attraction is primarily responsible for adsorption.

The same procedure applied to the data for the adsorp-

tion of the three acids on alumina does not give such consis-

tent results. At pH 5 AH* lies between -6 to -11 kcal/mole

for acetate, -2 to -7 kcal/mole for oxalate and -7 to -3 kcal/

mole for citrate.

Near Co = .4mM (see Figure 5), a medium concentration,

we find that AH* lies near -9.8 kcal/mole for acetate, -4.1

kcal/mole for oxalate and -8.4 kcal/mole for citrate. This

indicates that while acetate has one, and citrate two ionized

surface groups, oxalate has on the average only one of its

two ionized groups on the surface.

It was argued earlier that only ionized molecules are

attracted to the surface and that the number of such molecules

drops as the surface charge drops. It should be recalled

that the surface charge on alumina and hydroxyapatite is

produced by the adsorbed H+ ions. This suggests the possibility

that an approaching ionized carboxyl group and surface hydrogen

ion participate in hydrogen bonding.

Zeta potential measurements [100] show that citrate

changes the surface charge of hydroxyapatite at low concentra-

tions. This indicates that electrostatic attraction by itself

is not the only factor in adsorption. If it were, when the

surface charge had been neutralized by adsorption of a sufficient

amount of citrate, adsorption would have ceased and the surface

charge would not reverse.

The single charged acetate ion does not reverse the

surface charge of hydroxyapatite or alumina, this would

indicate the absence of other than electrostatic interaction.

It will be recalled from the introduction that .close approach

of hydrogen to an anion is required for hydrogen bonding to

occur. Since, at pH 7, acetate has one and citrate has three

ionized groups, it is thought that multiple groups are required

to pull the molecule close enough to the surface for hydrogen

bonding to occur. The interplay between surface charge,

molecular size and charge density, ionic and hydrogen bonding

becomes apparent in these situations.

Polyacrylic acid has such multiply-charged groups. It

is known to be a linear molecule which is fully ionized at pH 7

[ ]; therefore, there are no hydrogen bonds to be broken due

to an unfolding of the molecule upon adsorption. The AH between

-5.7 and -6.9 kcal/mole of residues of PAA is close to that

found for the adsorption of the carboxyl group of the other

molecules on alumina and hydroxyapatite. The similarity in all

these instances implies that the same type interaction occurs,

and is not a strict function of surface composition or of

molecular structure.


There appears to be no particular differences in the

enthalpy of adsorption of a carboxyl group onto alumina or

hydroxyapatite due to the number of groups on a molecule. In

each case the enthalpy change is near -6 kcal/mole and an

attracting force is required to initiate adsorption. Other-

wise a plot of E or Q versus pH would be similar in shape to

the a versus pH and not bell shaped.

Once adsorption occurs, the influence by multiply-

charged groups on the molecule was evidenced by a change in

surface charge. For this condition to arise, specific adsorp-

tion has to occur which requires forces other than electro-

static. The proximity of oxygen in the carboxyl groups and

H on the surface suggest hydrogen bonding.





Polysaccharides, notably chondroitin sulfate (CS),

which contain carboxyl and sulfate groups, are present in

dentin and enamel [101] and in bone [102,103]. Under proper

conditions of pH and ionic strength, these polysaccharides

will complex with collagen [104]. In the presence of a

foreign surface, these polysaccharides, like other charged

molecules, will adsorb and react not only with collagen but

also with the surface. It is the purpose of the studies of

this chapter to explain the interaction of aqueous solutions

of chondroitin sulfate with alumina, hydroxyapatite, and

silica. A later chapter will discuss the interaction of CS

with collagen.

Chondroitin sulfate is a polysaccharide made up of

basic dimer units of glucoronic acid and galactosamine. Several

simpler carbohydrates were chosen to model CS: (a) galactose,

glucose, D-acetyl galactosamine; (b) glucose-6-sulfate (G6S),

D-galacturonic acid; (c) dextran, and (d) polygalacturonic

acid (PGA). Each carbohydrate was chosen because it possessed

a single feature of the CS molecule: (a) the carbohydrate

residue; (b) a charged carbohydrate; (c) the polymer back-

bone structure; and (d) the charged polymer. The structural

formulaefor these molecules are shown in Figure 9a and 9b.


All carbohydrates were purchased from Sigma Chemical

Company [96] and used without further purification. The

chondroitin sulfate (sodium), dextran, and polygalacturonic

acid were reported to have molecular weights of 45,000, 60-

90,000, and 25,000 respectively. The chondroitin sulfate was

determined to be chondroitin-6-sulfate by infrared spectroscopy

using the KBr pellet technique, and had a molecular weight of

45-60,000. The materials used as substrates, Linde B alumina,

hydroxyapatite, and silica, and the low ionic strength solution

were described previously.


The first set of experiments was conducted with

galactose, glucose, and dextran (see Figure 9b). None of these

molecules possesses charged groups. By measuring the heat of

adsorption of these molecules on the oxides, the contribution

to the total enthalpy change on adsorption of uncharged

carbohydrate monomer and polymer could be estimated. The initial

concentration, C was varied from .01 to .1 moles/liter. By

comparison with the reaction heats produced by the carboxylic

acids in this concentration range, it was estimated that 20 to

40.meal would be considered a significant reaction. The results


OH / 3




.------ 28.7 A .

So S3


Figure 9. Molecular structure of chondroitin sulfate (a)
( D-glucoronic acid N-acety galactosamine-6-sulfate)
(b) the three-fold helix of chondroitin sulfate.












Galacturonic acid

Figure 9b. Molecular structures of the carbohydrates
used in the experiments of this section.

are presented in Figure 10. The maximum reaction heat, Q,

produced was about .3 mcal at a concentration of .1 M on

alumina. The enthalpy change, as determined by solution

depletion, was approximately -100 cal/mole adsorbed for galactose

and glucose on alumina. For dextran, AH was about -20 cal/

mole of residues (.11 cal/gm) and -15 cal/mole of residues

(.8 cal/gm) on hydroxyapatite.

The enthalpy change for the adsorption of D-acetyl

galactosamine on alumina was found to be -430 cal/mole of

molecules. Using the thermometric titration method described

earlier, the free energy change was found to be -5.2 kcal/mole.

The enthalpy change for the adsorption of this molecule on

hydroxyapatite was found to be -400 cal/mole of molecules.

Calorimetric measurements for the adsorption of D-

galacturonic acid on alumina indicated a stronger reaction than

with the uncharged carbohydrates. The reaction heat, Q, reached

a maximum of 8.9 meal. The calculated enthalpy change and

that determined by solution analysis are presented in Figure 11

and lie near -10.5 kcal/mole. There is good agreement for the

enthalpy change using both methods except in the lower concentra-

tion range where the calculated values are more negative. The

free energy change varies between -4.6 and -5.5 kcal/mole

adsorbed. The entropy change is negative and lies between -8

and -30 cal/mole*deg per molecule adsorbed.

Measurements of the adsorption of D-galacturonic acid on

hydroxyapatite showed a similar enthalpy change to that on

r. 0.3


o I I I

14 n0

.02 .04 .06 .08
Co, M

Figure 10. Reaction heats for the adsorption of dextran o ,
galactose ,, and glucose o, on alumina (closed symbols) and
hydroxyapatite ( open symbols) in low ionic strength solution
at pH 7.

- -

- :- -- --Z --

- -




C, M

Figure 11. Thermodynamic data for the adsorption of D-galac-
turonic acid on alumina at pH7 in low ionic strength solution;
AG--- AH- AS---. Enthalpy change determined by solution
analysis -*-




alumina, as determined by solution depletion (see Figure 12).

Values for AH lie between -6 and -8 kcal/mole.

The adsorption of D-galacturonic acid on silica is

endothermic. At low concentration, the enthalpy is +700

cal/mole, becoming more positive at higher concentrations.

The amount adsorbed was determined by solution depletion. The

free energy change and entropy change were not calculated.

The change in enthalpy for the adsorption of glucose-6-

sulfate on alumina, hydroxyapatite, and silica was determined

to be -7.6, -5.4, and 0.3 kcal/mole of adsorbed molecules,



Calorimetric measurements for the adsorption of poly-D-

galacturonic acid on alumina and hydroxyapatite were hampered

by agglomeration of the particles by the polymer. The thermo-

dynamic functions for this reaction were calculated and are

shown below.. In the concentration range used, .001 M to .01 M

of residues, these values were constant (1 0.1 kcal/mole).

(kcal/mole) (kcal/mole) (kcal/mole/
alumina -3.4 -2.54 3.5

hydroxyapatite -4.2 -.31 15.0

In contrast to the decrease in entropy found for the

adsorption of PAA on alumina or hydroxyapatite, the entropy

change is positive for adsorption of poly-D-galacturonic acid

on both alumina and hydroxyapatite indicating an over-all

decrease of ordering.


2 -


o -2 -



C0 M

Figure 12. Heat of adsorption for D-galacturonic acid on
silica-e-- and hydroxyapatite -- in low ionic strength
solution at pH 7.


The results of the calorimetric measurements of the

adsorption of chondroitin sulfate (CS) on alumina are

presented in Figure 13. Comparison of the enthalpy of

adsorption for CS on silica, alumina, and hydroxyapatite are

given below:


Silica +2.46

Alumina -1.85

Hydroxyapatite -2.47

The enthalpy changes for the adsorption of CS on the

uncharged hydroxyapatite is found to be 15-30% more negative

than that for the positively charged alumina. The enthalpy

change for adsorption of CS on silica is found to be positive

endothermicc) and was determined by solution depletion, the

free energy change AG was not calculated.


Glucose, galactose, and dextran are uncharged molecules.

Because of the availability of hydroxyl groups it is possible

that these molecules can undergo hydrogen bonding with an

oxide surface. The low value of the enthalpy change for the

adsorption of these molecules, however, does not indicate very

strong reaction with the oxide surface in comparison with the

charged molecules.


6 30

S 4 .. 20

S 0


H 3
o U


0 I I
0.2 0.4 0.6 0.8

CO, mM

Figure 13. Thermodynamic data for the adsorption of chon-
droitin sulfate on alumina in low ionic strength solution
at pH 7; AG----, AH- AS--.

The only difference in the experimental conditions

in using the uncharged versus charged molecules is

attributable to the functional groups. Therefore the large

differences in AH observed are seen as due to the presence of

the charged groups.

If it is assumed that, in the case of dextran, only

two or'three points of contact are made per molecule, then

the enthalpy change could be on the order of 3-4 kcal/mole

molecules. Comparison of dextran with poly-d-galacturonic

acid (PGA), however, still shows that dextran is much less

strongly bound than PGA.

In the concentration ranges used, 10-6 M dextran

molecules (not residues), there is little hydrogen bonding

of dextran chains to one another [76,98]. Therefore, the

breaking of interchain hydrogen bonds should not contribute

substantially to the low enthalpy change actually measured.

If the hydrogen bonding does take place between the oxide sur-

face and many dextran residues, it is not manifested in the

measured enthalpy change. In the absence of other interactions,

it is concluded that the charged groups of the carbohydrate

molecules are necessary to provide sufficient attraction of

the entire molecule to the surface.

The increase in the heat of adsorption of D-acetyl

galactosamine over that of galactose is attributed to the

presence of the NH2COCH3 side chain. The exact cause can only

be speculated. Perhaps the nonpolar methyl group is forced

from solution by more polar solvent ions, drawing the molecule

to the surface. Whatever the mechanism, if these uncharged

molecules are strongly bound to the surface, it is not

reflected in the enthalpy determination. The type of bond

which would occur would almost certainly be hydrogen.

In any event, the adsorption of D-galacturonic acid

on alumina and hydroxyapatite is much more energetic than that

of galactose or galactosamine. The similarity in the molecules

and in the adsorption experiment strongly suggests that the

charged carboxyl group is responsible for the higher enthalpy

change and that binding between solute molecules or desolvation

effects do not account for the noted change. The negative

entropy change indicates an overall increase in ordering. This

increase in ordering may be due to confinement of the carbo-

hydrate molecules to the surface and subsequent loss of freedom


Adsorption of D-galacturonic acid on silica produces a

positive enthalpy change. This can be accounted for by the

charge repulsion which exists between the surface and the molecule.

There was a finite amount of acid adsorbed, however. This

would indicate perhaps a second stronger force necessary to

overcome the charge repulsion or that the negatively charged

molecules are occupying the fewer positive siteS on the silica

surface, or reaction with high energy sites. Since the reaction

heat and adsorption measurements leveled off quickly, the last

two possibilities appear more likely; especially in view of the

finding that charge attraction appears necessary for strong


Likewise the adsorption and calorimetric measurements

of glucose-6-sulfate suggest that the presence of a charged

group on the carbohydrate, opposite to that of the surface,

is required for stronger (more exothermic) reactions. The

negative sulfate group is attracted to the positive alumina

surface. As with'D-acetyl galacturonic acid on silica, the

adsorption of glucose-6-sulfate on silica is endothermic.

Poly-D galacturonic acid is obviously strongly attached

to the alumina and hydroxyapatite surfaces. The enthalpy

change is more negative for alumina, than for hydroxyapatite,

demonstrating the greater attraction for this surface. If

we consider that the enthalpy change is produced by the

charged groups bonding to the surface, then, using a figure

of -9kcal/mole as the enthalpy change found for the adsorption

of D-galacturonic acid on alumina, we can estimate that one

in three residues bonds to the surface. For hydroxyapatite,

this figure is perhaps one in twenty or thirty.

Chondroitin Sulfate

The chondroitin sulfate molecule is known to exist in

a threefold or eightfold helix [105,106] which is rigid in

solution [107] (see Figure 9). It possesses the ability to

change the conformation of positively charged polypeptides

from an extended coil to a helical structure (see Chapter 5).

The sulfate group extends further away from the carbohydrate

backbone than does the carboxyl group which is located on the

other side of the same dimer. The acetyl amine group extends

slightly further away from the backbone than does the

carboxyl and is located on the same side of the backbone.

From these considerations--helical structure, position of

the charged groups, and possible steric hinderance--it is

reasonable to assume that not all the charged groups participate

in bonding with the surface at the same time,

To help analyze the binding of CS to a surface, consider

that the interaction of a single dimer with the surface

permits interaction of both the carboxyl group and sulfate

group with the surface. Both groups would then contribute to

the enthalpy of the reaction. From the data in this chapter

and the previous one, it is seen that the change in enthalpy

is fairly constant for each type molecule, as it is between

carboxyl groups and that only charged groups contribute

significantly to the reaction heat, Q. The reaction heat due

to the adsorption of a carboxyl and sulfate group would be

between -12 and -16 kcal/mole. The measured enthalpy value is

about -2 kcal/mole of dimers for alumina and between -2.2 and

-2.8 kcal/mole for hydroxyapatite. Dividing the total enthalpy

possible by the measured value would indicate that between one

in three to one in seven dimers interact with the surface.

We may assume that only one of the charged groups inter-

acts per dimer. Using an enthalpy change between -6 and -12

kcal/mole of charged groups then one in three to one in five

groups would be indicated as interacting with the surface.

From the physical picture and the calorimetric data it seems

plausible to conclude that the.chondroitin sulfate molecule is

positioned horizontally on the surface with approximately

one out of every four dimers on the average coming into

contact with the oxide surface. In this instance, the charged

group interacting with the surface would be the sulfate group

since it extends further away from the CS backbone.

The entropy change due to adsorption of CS on alumina

is positive, indicating an increase in entropy or a decrease

in the order of the system. Since the molecule is rigid in

solution and is not likely to greatly change conformation on

the surface, no entropy contribution is attributable to a

change in shape. An increase in entropy can be attributed to

a release of solvent molecules from around the molecule or

from the surface into solution [106,107].

The adsorption of CS on silica can also be explained by

an increase in entropy. The surface charge on the silica is

the same as that on the carbohydrate. Overall, there is an

electrostatic repulsion between the surface and the charged

molecule. There must be some other energy supplied to over-

come this repulsion. Since it is the free energy change which

drives the reaction, and AH is positive, there must be at

least an equivalent positive entropy change so that TAS is

greater than AH. This entropy change, as suggested above,

can be supplied by the solvent ions.

It is difficult to speculate on the conformation of

CS on the silica surface as was done above with alumina. This

is so because the carbohydrate monomers are not attracted to

the surface of the silica as they are to alumina because of

charge repulsion.


In this section we have determined some of the thermo-

dynamic features of the adsorption of chondroitin sulfate on

alumina, hydroxyapatite, and silica by the use of model

carbohydrates. The results show that for positively charged

alumina the enthalpy change for the adsorption of charged

carbohydrates is about the same as that for the carboxylic

acids and lies between -7 and -9 kcal/mole of adsorbed species.

In these experiments the entropy change is negative and the

enthalpy change forms the major portion of the free energy

change. The enthalpy change for the adsorption of the charged

monomeric carboxylated carbohydrates on hydroxyapatite is close

to -7.6 kcal/mole. The adsorption of carbohydrates containing

the sulfate group on alumina or hydroxyapatite is about -5.4

kcal/mole. The similarity in enthalpy for the reaction suggests

a similar type reaction. The adsorption of charged carbohydrate

monomers on silica is weak and endothermic. The uncharged

monomers produce an enthalpy change only a fraction of that of

the charged monomers.

It was found that a model of the uncharged polymer backbone

of CS does not produce a large reaction heat or an enthalpy

change, suggesting that the presence of charged groups is

required to enhance the adsorption reaction. The data on

adsorption of polygalacturonic acid supports this conclusion.

Polygalacturonic acid adsorbed strongly, producing

agglomeration of the solid particles at high concentration (.1 M).

The enthalpy change per residue is lower than that found

for its monomeric counterpart. The conclusion drawn here,

as with polyacrylic acid in the previous chapter, is that

fewer points of contact are made, but that each point of

contact contributes essentially the same heat change as the

monomer. The entropy change was positive further increasing

the driving force.

Chondroitin sulfate was shown to adsorb to each of the

oxide powder substrates. The negative enthalpy change for

alumina and hydroxyapatite indicates that adsorption is strongly

enhanced by the opposite charges. Since the chondroitin sulfate

molecule is comparatively bulky relative to the models used,

fewer points of contact would be expected. The thermodynamic

calculations show, however, that perhaps as much as one-third

to one-fourth of possible bonding sites touch the surface.

The results for CS on silica are more speculative. The

enthalpy change is positive and it.is assumed that the major

contribution to the free energy change of adsorption is from

a positive entropy change related to the molecule size. Smaller

molecules were not found to produce a positive entropy change.

Relying on the results of the previous chapter, it

is assumed that once the molecules are attached to the surface,

that hydrogen bonding will take place.

In a subsequent chapter, we will investigate the type

interaction which CS and other molecules undergo with collagen

and collagen models. In the next chapter a study of the adsorp-

tion of molecules possessing amine side groups is discussed.




The calorimetric measurements discussed in the previous

chapters were related to the adsorption of molecules which

possessed a carboxyl group. The results showed that the mole-

cules on which the carboxyl group was ionized displayed greater

reaction heat and adsorption density than those molecules

which were not ionized. In this chapter, molecules containing

charged amine side groups are studied for the possible informa-

tion they can give on the adsorption of collagen onto silica,

alumina, and hydroxyapatite.

Several of the primary amino acids and their respective

polymers were investigated as collagen models. The molecules

used were alanine, poly-l-alanine (PA), proline, poly-l-

proline (PLP), poly-1-hydroxyproline (PLHP), lysine, poly-1-

lysine (PLP), and poly-l-arginine (PLA). Lysine and arginine

and their polymers possess basic side chains. The structural

formulaeof the monomers are given below. The polymers are linked

at the carboxyl and amino groups.

+NH3 H H

Alanine Proline

()coO- +H3N(CH2)4CHOO-


Hydroxyproline Lysine

+NH2 NH2


The amino acids are dipolar ions. For the dibasic

amino acids, arginine and lysine, adsorption onto negative

surfaces will be enhanced. In polymeric form only poly-l-

lysine and poly-l-arginine retain any charge in neutral


The solubility in water of the other amino acids will

decrease as a result of their polymerization. Acidic amino

acids, aspartic acid and glutamic acid were not studied because

the carboxyl group has been discussed in the previous chapters.

Furthermore, interpretation of calorimetric data would be

difficult because of the presence of three charged groups.


The amino acids used in this section were purchased

from Sigma Chemical Company. Both amino acid monomers and

polypeptides were chloride salts, except for poly-l-lysine,

which was a bromide salt. The molecular weights of the poly-

peptides were reported to be: 1,000-5,000, poly-l-alanine;

15,000-50,000, poly-1-arginine; 70,000, poly-l-lysine;

10,000-30,000, poly-l-proline; and 10,000-30,000, poly-l-


The substrates and the solvent are the same as used

in the previous chapter. A few experiments were performed in

the .165 M salt solution (and will be indicated as such in

the text).

Mixing and calculation procedures were described




Calorimetric and adsorption results for alanine and

PA on the three oxides are presented in Figure 14. The

amino acid monomer adsorbs strongly on all three surfaces at

this pH. It acts much as the carboxylic acids do. There is

no sharp endpoint in the Q vs. C curve for the monomer, but

the slope of the surve is greater than that for the polymer.

The enthalpy change for each surface tends towards -5 to -8

kcal/mole of monomers.




alanine TCP

3 alanine A1203




PA Al 203

I p I

0.0 0.1 0.2 0.3 0.4 0.5

Co mM

Figure 14. Heats of adsorption for alanine and poly-l-alanine
(PA) onto alumina, silica, and hydroxyapatite (TCP) in low
ionic strength solution at pH 7. run in .165M salt solution.

When the dissolved salt (R-C1) concentration in the

solvent is increased to .16 M, the enthalpy change of the

adsorption of alanine on alumina decreases in the higher

concentrations range of alanine, but tends toward the -6

kcal/mole in the lower end.

The polymer, PA, exhibits much different behavior

showing no specific tendencies to adsorb. Reaction heats are

less than .5 kcal/mole and enthalpy changes are of magnitude

less than -1 kcal/mole.


Measurements made with proline and PLP and hydroxy-

proline and PLHP (see Figure 15) show results similar to

results for alanine and PA. The monomer again produces a

higher enthalpy change than the polymer (-4 to -8 kcal/mole).

However, PLP and PLHP are apparently more strongly attracted

than the poly-l-alanine with heats of adsorption lying between

-0.5 and -2.5 kcal/mole of residues. PLHP is somewhat more

strongly attracted on all three surfaces than is PLP. An

increase in ionic strength is noted by a decrease in the

enthalpy change for PLHP on alumina.

Lysine and Arginine

The addition of charged side groups causes a marked

change in the enthalpy curves, as shown in Figure 16. The

heat of adsorption of lysine on silica is, as expected,


1. P-A
2. HP-A
3. HP-S
1 4. HP-HA
5. PLP-S
2 9. PLHP-A*
10.PLP -S

3 0

2 7


*1 -

.02 .04 .06 .08
Co, mM
Figure 15., Heats of adsorption for proline (P), hydroxyproline
(HP), poly-1-hydroxyproline (PLHP), and poly-l-proline (PLP),
on alumina (A), silica (S), and on hydroxyapatite (HA) in low
ionic strength solution and Ringers solution (*).






~ 6




0.2 00.4 0.6 0.8
Co, mM

Figure 16. Heats of adsorption for lysine (L), poly-1-lysine
(PLL), and poly-1-arginine (PLA), on silica (S), alumina (A),
and hydroxyapatite (HA) in low ionic strength solution at
pH 7. Values determined by solution analysis -e- others
by calculation.

greater in magnitude than on alumina and is more pronounced

in the lower concentration end. The enthalpy change lies

in the range of -2 to -8 kcal/mole of monomers. The greatest

enthalpy change, however, for the polypeptides was recorded

for PLA on alumina. The second largest was on hydroxyapatite

and third on silica. This order also happens to be the order

of decreasing specific surface area. The enthalpy of adsorp-

tion of PLA was greater than that for PLL on hydroxyapatite.

The free energy change for PLA and PLL lie in the

range between -4 and -7 kcal/mole of residues (see Figure 17).

There is a sharp decrease in the free energy change of PLA

on hydroxyapatite at low concentrations. The entropy changes

for these molecules are small because of the similarity of

AH and AG.



The higher enthalpy change for the adsorption of

alanine, compared to poly-l-alanine, is due to the electro-

static attraction of the amine group or the carboxyl group

to the surface. PA, having no charge except for its terminal

groups is not strongly adsorbed despite its greater molecular

weight. PA is in a helical form [20], not coiled.

Since the molecule is uncharged,.the reaction heat of

the entire molecule due to adsorption only would be small in


H 6

S/ PLL-SiO 2

4 PLL-Al /



0.0 0.2 0 .4 0.6 0.8 1.0

Co, mM

Figure 17. Free energy change for the adsorption of poly-1-
lysine (PLL) and poly-l-arginine (PLA) on silica, alumina and
hydroxyapatite at pH 7.

comparison with ionized molecules [50]. A large conformational

change would then cause the overall reaction to be endo-

thermic. The molecular concentrations are low. Therefore,

breaking of intermolecular hydrogen bonds should contribute

little to the enthalpy change.


The charged monomers of proline and hydroxyproline

are also more strongly attracted to the oxide surfaces than

their polymers (Figure 15). This is taken as a result of

electrostatic attraction. The heat of adsorption, measured

by solution depletion and found to be between -4 and -6 kcal/

mole for both monomers, results from reaction of the charged

groups with the oxide surface.

Calorimetric measurements for the adsorption of PLP

and PLHP did not show a specific pattern for any of the

surfaces. Both of these molecules have a helical conformation

in solution [20]. There was no attempt to determine whether

or not this structure was grossly disturbed upon adsorption,

or if it was, what contribution to the enthalpy change such

a disruption would make. Neither was there an attempt to

determine how many points of contact were made. The definite

conclusions which can be drawn from these data are relatively

few. There are some reasonable assumptions, however, that can

be made which, if accepted, will further explain the situation.

Since these molecules are in an extended conformation,

not coiled, there are no intramolecular hydrogen bonds. At

low concentrations (2 x 10-5 M of residues) there should be

little intermolecular hydrogen bonding [50]. The disruptions

which would primarily occur, then, would correspond to

rotational movement of the molecule [19], There is no

reason to suspect that these molecules should undergo grotesque

distortion on the surface since there are no strong attractive

forces. Therefore, the contribution to energetic changes due

to conformation alterations should be small.

Poly-l-proline and poly-l-hydroxyproline possess a

ring structure which is relatively nonpolar compared to the

polar solvent. Because of this, it is plausible to assume

that these less polar structures are in a lower energy state

on the surface, rather than in the solution. This is termed

hydrophobic bonding and could possibly account for the energetic

changes measured if most of the residues were near the surface

and not surrounded by the mobile polar ions in solution.

By comparison with the carbohydrates studied earlier

and in absence of detailed information on the geometric

smoothness of the surface, it would also be possible to suggest

that only a few points of contact are made on the surface [50].

Each of these contacts would assume a higher energy than

indicated by the average of 1-2 kcal/mole of residues measured.

If these contacts are hydrogen bonds made up of hydrogen

atoms on the ring structure and oxygen atoms on the surface,

then each bond would entail an energy change of about

-7kcal/mole. On the average then one out of 7 residues

would be in contact with the surface.

Comparison of the adsorption of these uncharged

molecules with those that possess charged functional groups

indicates that the role of PLP or PLHP would be minor in

comparison. Although the mechanism of adsorption has not

been fully explained in this case, there should be little

doubt that when positioned next to a charged molecule in a

peptide chain, the latter will play the dominant role in

adsorption to an oxide surface.

Lysine and Arginine

The adsorption heats of lysine on silica and alumina

are similar (Figure 16). Since this molecule possesses two

basic and one acidic group at neutral pH, it is reasonable

to assume that the charged carboxyl group is attracted to

the positive alumina surface, and that the positive amine

groups are attracted to the negative silica surface. Because

of the more basic properties of this molecule, it might be

suspected that the reaction with the silica surface would be

somewhat stronger than with the alumina surface. This appears

to be the case.

Poly-l-lysine and poly-l-arginine are known to exist

in an extended charged coil conformation in solution at

neutral pH [8,9]. If the coils, which are stabilized by

hydrogen bonding, were to break down, the enthalpy change

would be due to both the adsorption and the unfolding processes.

Enthalpy changes measured in this and the two previous chap-

ters indicate that enthalpy changes between -6 and -8

kcal/mole of single, charged groups are to be expected. The

magnitude of the unfolding process, the breaking of hydrogen

bonds, would lie between 5 and 7 kcal/mole [19,82], an

endothermic process. The resulting enthalpy changes for both

processes would be comparatively small with a value near 0

kcal/mole. Such a situation is encountered in experiments

presented in the next chapter where the coil-to-helix

transition is known to take place. Instead,the enthalpy

change is much more negative, decreasing in magnitude from

-14 kcal/mole to -6'kcal/mole. The variation is thought to

be due to adsorption on the fewer negatively charged sites

on the alumina and hydroxyapatite surface. These sites possess

a distribution of high to low energy within their own group.

Adsorption of PLL or PLA to neutral high energy sites

can be eliminated since PA, PLHP, and PLP each have heats of

adsorption which are smaller in magnitude. If the adsorption

had not depended on surface charge the neutral molecules

should have been just as strongly attracted to the surface.

The order of decreasing enthalpy change (alumina >

hydroxyapatite > silica) corresponds to the specific surface

area of the solids. The higher specific area of the alumina

(87m2/gm) provides more edges and peaks which are assumed to

form high energy sites. For equal amounts adsorbed, a

greater fraction of adsorbed molecules would be on these

sites for alumina, than for either hydroxyapatite or silica.

Negatively charged polysaccharides also displayed an increase

in enthalpy change at lower concentrations attributable to

high energy sites. More than a single species of these

specially adsorbing areas appears likely [72]. The free

energy changes are of the same order of magnitude as those

found for the carboxylic acids and polysaccharides.


As found in the previous chapters, molecules with

charged ionic side groups react more energetically with oxide

surfaces than those without. The enthalpy change per

charged group lies between -4 and -10 kcal/mole. For those

polyamino acids which possess no charged side groups, the

magnitude of the enthalpy change is found to be less, near

-1 kcal/mole of residues. Although the mechanism for uncharged

molecules is not clearly defined, it is believed that conforma-

tional deformation does not contribute significantly to the

enthalpy changes measured. It is possible than an uncharged

polyamino acid could be bound to an oxide surface by a few

relatively high energy contacts. Polyamino acids with charged

side groups, however, will play the dominant role in adsorption.





In the earlier chapters, an investigation of the

adsorption of molecules onto oxide surfaces has been discussed.

These calorimetric studies of compounds which are models for

collagen provided some information on the state of the adsorbed

molecules, the energy changes on adsorption and the type of

interaction they undergo.

In vivo, it is unlikely that a collagen molecule will

come into contact with a clean surface. In general, there

will be other substances present in the body fluids which

will adsorb first because of factors such as greater concentra-

tion. Also, collagen may not be present at all when the

hydrated oxide surface is first exposed to body fluids [102].

We should have some indication, therefore, of how these adsorbed

molecules will affect the adsorption of collagen. In order

to understand their interaction at a liquid-solid interface,

it would be helpful to first investigate their .interactions

in solution.

It is the purpose of this chapter to provide further

insight to the interaction of collagen, chondroitin sulfate,

and serum albumin in solution.

The same compounds which have been used previously

to model collagen and chondroitin sulfate have been used

here. For collagen they are poly-l-arginine (PLA), poly-l-

lysine (PLL), poly-lalanine (PA), and poly-l-proline (PLP).

For chondroitin sulfate they are dextran, galactose, galacturonic

acid, and polygalacturonic acid.

The reaction of chondroitin sulfate with collagen has

been studied by model systems, as mentioned in earlier

chapters. The reaction of collagen with dyes containing acidic

groups and with CS have been studied in regard to the under-

standing of the role of CS and collagen in connective tissue

[108]. It was found that the cationic groups of collagen bond

with the anions of the dyes and CS in a pH range of 1.5 to

7 with a sharp drop in the number of anions fixed below pH 2

and a more gradual decrease from maximum adsorption at pH 3

to zero at pH 7.

In experiments to determine the role of CS in the

calcifying mechanisms of bone [109], it was found that calcifi-

cation would not occur or would occur more slowly in an aqueous

collagen mixture when CS was not present. These experiments

were performed near pH 7. The authors suggested that binding

of CS to collagen at neutral pH would aid the natural calcifica-

tion process.

In another experiment with chondroitin sulfate [13]

and cationic dyes, it was concluded that aggregation of

dyes on the surface of the CS rather than ionic interaction

was mainly responsible for bonding. The thermodynamic functions

indicated an enthalpy change of -7 to -12 kcal/mole of dye


The binding of cobalt hexammine (Co(NH3)6+3) to connective

tissue, micropolysaccharides, heparine, and sulfated chitosans

has been studied by a spectrophotometric procedure [110]. The

cobalt hexammine was used to represent amino functions of fibrin.

Ion pair formation was found to be the primary binding mechanism,

but was influenced by local binding factors, electrostatic attrac-

tion of neighboring charged groups, and competition with other

cations for binding sites.

Such previous work generally indicates that CS-collagen

or CS-polypeptide binding will be primarily ionic, pH, and

structure dependent. These previous results and the results

discussed in this chapter will serve as an aid in understanding

later calorimetric measurements.


In these calorimetric measurements, a solid substrate

was not used. In the first set of experiments, the concentra-

tion of saccharides was held constant at a concentration Co

of 10-3 M of saccharide monomers or residues. In the case of

CS, this refers to dimer residues. Aquisition of the organic

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