Table Of ContentELECTRICA L INTERACTION S
IN MOLECULA R BIOPHYSIC S
An Introductio n
RAYMOND GABLE R
Millipore Corporation
Bedford, Massachusetts
1978
ACADEMIC PRES S New York San Francisc o Londo n
A Subsidiary of Harcourt Brace Jovanovich, Publishers
COPYRIGHT © 1978, BY ACADEMI C PRESS, INC.
ALL RIGHTS RESERVED.
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ACADEMI C PRESS, INC.
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United Kingdom Edition published by
ACADEMI C PRESS, INC. (LONDON ) LTD.
24/28 Oval Road, London NW1 7DX
Library of Congress Cataloging in Publication Data
Gabler, Raymond.
Electrical interactions in molecular biophysics.
Includes bibliographies.
1. Biological chemistry. 2. Biological physics.
3. Molecular biology. 4. Electrostatics.
I. Title. [DNLM: 1. Molecular biology. 2. Bio-
physics. QT34 Gll5m]
QP514.2.G3 574.1*92 77-6595
ISBN 0-12-271350-8
PRINTED IN THE UNITED STATES OF AMERICA
PREFAC E
This book is designed to act as an introductio n to the subject of electrical
interaction s between biomolecules. As such, it encompasses two subjects,
molecular biology and physics. The text is written to give the reader an
appreciatio n of the fact that at the fundamenta l level, all biochemical or
molecular biological reactions are electrical in nature , and as such must
conform with the laws of physics and electrostatics . In this same respect, no
new laws of physics have to be postulated to explain the interaction s of bio-
molecules. The text also illustrate s the wide variety of ways in which bio-
molecules can interact with one another , and how use is made of the electrical
propertie s of biomolecules in order to obtain informatio n about their
structur e and function. The principles of electrostatic s are used to explain
some of the basic units of structur e on a molecular level.
Although both biochemistr y and physics are covered, the reader need not
have prior experience in either. The introductor y chapter s on both subjects
are written to give a simple and basic understandin g of the backgroun d
materia l necessary to understan d the rest of the text. Familiarit y with the
elements of differentia l and integral calculus is, however, assumed. It is
possible to take a concurren t course in calculus, but this is thought to be
more difficult. In this respect then, this text could be used by any person in
the fields of biochemistry , molecular biology, physics, or microbiology.
The book is organized in such a fashion that the chapter s generally build
on the material s in the previous chapters . Mathematic s is kept to a minimum ,
although Chapter s 4 and 7 will requir e special attention . Chapter 1 presents
the concepts and structure s of biochemistr y needed to understan d the rest
of the text. Emphasis is placed on knowing how to recognize different
structura l biochemical groups and how they are used as building blocks in
forming molecules. Very few chemical reactions are considered, so the reader
need not memorize a great deal. Chapter 2 covers the basics of elementar y
electrostatics . This chapter lays the foundation for the formulatio n of the
physical theories presented in the other chapters . Both Chapter s 1 and 2
assume no prior experience in these subjects. Chapter 3 builds on Chapter s
ix
÷ PREFACE
1 and 2 and considers dielectrics and dipole moments. This is the first chapter
in which biochemistr y and physics are mixed. Chapter 4 deals with the dipole
moments of biomolecules, how the dipole moment is measured , and what
it means in terms of the physical structur e of a molecule. In Chapter 5,
equations are derived to describe strength and other propertie s for a wide
variety of different electrical interactions . Biological examples are given to
illustrat e the importanc e and occurrenc e of each. Chapter 6 is devoted
entirely to van der Waals forces, which are another form of electrical
interaction . These unique and generally unappreciate d forces are described
in general terms and several examples are given to illustrat e their importance .
Chapter 7 is also devoted to one subject, Debye-Huckel theory. This theory
seeks to describe the much more realistic situation of electrical interaction s
in a solution containing counterions . Chapter 8 treats water and water
structur e from a physical standpoin t and indicates water' s role in the overall
scheme of molecular biology. Finally, Chapter 9 describes experimenta l
techniques that rely upon the electrical propertie s of biomolecules. This
chapter also discusses what types of informatio n can be obtained from each
experimenta l form and emphasizes that the electrical natur e of biomolecules
can be used to yield informatio n of importanc e to research workers and also
to physicians in dealing with disease.
The electrical natur e of biomolecules and their interaction s can be an
exceedingly complex subject, and this book is written to give the reader a
fundamenta l overview of the basics, and perhap s to spark an interest in
further study of the physical reasons for biomolecules behaving as they do.
ACKNOWLEDGMENTS
I want to thank Dr. Norman Ford for reading the manuscrip t and supply-
ing me with many useful comments and criticisms. I would also like to thank
my wife Rose Mary for proofreadin g the manuscript , suggesting changes,
and for endurin g my many nights of writing and rewriting .
Thank s are also extended to Drs. Edward Westhead and Kenneth Langley
for numerou s helpful conversations , and to Dr. Irwin Bendet for initially
introducin g me to biophysics.
xi
1
BIOCHEMISTRY
INTRODUCTION
The science of biophysics is concerned with the use of physical principles
and techniques to solve biological problems. Biophysics is an interdisci-
plinary science that require s a knowledge of physics, biology, and chemistry.
It is the type of discipline where various fields of knowledge overlap, and
only by considering them all can a complete and organized overview be
obtained. Molecular biophysics is that branch dealing with the structur e
and function of the basic biomolecules that are responsible for the different
forms of life systems. Even though biology is usually thought of as a disci-
pline completely independen t of physics, when one examines biological
reactions, on either a macroscopic or submicroscopi c level, it is known that
biomolecules, cells, etc. must obey the fundamenta l physical laws. Just
because an entity is involved with a life system does not exempt it from the
rules of physical behavior exhibited by the rest of the universe. In consider-
ing the specific case of molecular behavior it should be emphasized that at
the most basic and fundamenta l level all interaction s are electrical in nature ,
and that structura l and functional aspects of biological macromolecule s
are determine d by electrical laws. The fact that the basic electrical laws
cannot be applied with certaint y to all cases of interest is an indication of
the complexity of biomolecular systems or of the inadequac y of our under-
standing of the physical universe.
This textbook is an attempt to make the reader more aware of the im-
portanc e of electrical interactions , and in particula r to emphasize their
application to molecular biology. It is an effort to illustrat e how basic
biological processes can be governed, influenced, and controlled by physical
forces.
This chapter will cover the basic biochemical informatio n the reader
should be familiar with before proceeding to other chapters . To those who
are familiar with biochemistry , this will be a review. To those who have
never had exposure to biochemistry , this chapter represent s the minimum
knowledge that should be mastered . The emphasis here will be on the ability
é
2 1 BIOCHEMISTRY
to recognize different types of chemical compound s that are importan t in
biophysics and biochemistry . Very few actual chemical reactions will be
discussed in detail since this knowledge is not going to be needed to any
large extent. The types of electrical interaction s a biomolecule can enter
into and the strength of these interaction s are ultimately determine d by the
chemical makeup of the molecule itself. It is therefor e importan t that the
reader be familiar with the various parts of a biomolecule from a chemical
standpoint . The materia l in this chapter is of a general nature , and any
amplification s of details that are necessary will be made throughou t the
text as they are needed. For those who are interested in learning more or
who would like to see another presentation , several organic chemistry and
biochemistr y textbooks are listed at the end of the chapter .
BIOELEMENTS
Before getting into the biochemical aspects directly it is instructiv e to
take a look at the basic building blocks of all biological molecules, namely,
the bioelements. Any molecule, whether large or small, organic or inorganic,
etc. will have physical and chemical propertie s that are determine d by the
individual atoms composing it. Biological molecules tend to be thought of
as being rather special because they are involved in a process called life,
but they too are constructe d from a collection of atoms. In examining a
wide variety of these molecules it has been noticed that they are built from
a preponderenc e of a certain few atoms, rather than from a wide distributio n
of all the elements seen in the periodic chart. That this is true indicates that
this small group of elements apparentl y possesses certain propertie s that
make them especially useful to the constructio n of biological molecules.
These bioelements, then, deserve a special look to see just what propertie s
they have that makes them so special.
One might naively think that the percentage s of elements found in bio-
molecules would parallel the percentage s of elements found in the earth' s
crust. This seems natura l since the more abundan t an element, the more
likely it will be incorporate d into a growing system of molecules. However,
this would be true only if all the elements possessed the same properties ,
and none had electrical configuration s that would favor one element over
another . The percentage s for some of the more common elements found
in the earth' s crust are shown in Table 1-1. In terms of biological molecules
the four most commonly found elements are hydrogen (H), carbon (C),
oxygen (O), and nitrogen (N). In all of natur e these four elements make up
well over 90% of all living matter . Two other elements that are extremely
importan t are phosphoru s (P) and sulfur (S). Ions that are importan t in
biological processes are sodium (Na+), potassiu m (K+), magnesiu m (Mg+2),
BIOELEMENTS 3
TABLE 1-1
Common Elements in the Earth' s Crust and Their
Relative Abundance"
Element Percent Element Percent
Oxygen 47 Carbon 0.2
Silicon 28 Potassium 2.5
Aluminum 8 Titanium 0.4
Sodium 2.5 Hydrogen 0.2
Calcium 3.5 Nickel 0.2
Iron 4.5 Copper 0.002
Magnesium 2.5 Zinc 0.001
Phosphorus 0.1
a Data from "Van Nostrand's Scientific Encyclopedia,"
fifth ed. Van Nostrand, Princeton, New Jersey, 1976.
calcium (Ca+2), and chloride (Cl~). Also encountered , but usually in very
small or trace amounts, are manganese (Mn), aluminum (Al), vanadiu m (V),
iron (Fe), cobalt (Co), copper (Cu), molybdenu m (Mo), iodine (I), silicon (Si),
zinc (Zn), and boron (B). So, these 22 elements just listed are the ones that
are most commonly found in biological systems. Considerin g that there are
well over one hundre d elements in the periodic table, it is easily seen that
the incorporatio n of atoms into biomolecules has been restricte d to a fairly
small group.
If we now take a look at the four most common elements and their place-
ment in the periodic table (Fig. 1-1), several things are noticed. Being the
simplest of all elements, hydrogen is in the first period, while C, N, and Ï
are in the second period. This means that the most common bioelements
are among the lightest. The electronic structure s of these atoms are shown
schematically in Fig. 1-2 where each dot represent s an electron in the outer
shell. More will be said in Chapter 8 concerning the geometric arrangemen t
of the orbiting electrons. For our purposes, it is enough to note that the
number of electrons in the outer orbits has a large role in determinin g an
element's ability to form chemical bonds with other elements. The inner
electrons are more protected and hence cannot react as well with other
atoms. The number of electrons an atom has surroundin g it is also importan t
from the standpoin t of stability. It is well known that atoms with certain
number s of electrons tend to possess greater stability than those that do not
have these special configurations . For the lighter elements two or eight
electrons surroundin g a nucleus in the outer shell is an advantag e in terms
of stability. In this sense we see that Ç require s one additiona l electron to.
attain a stable configuration , while Ï , N, and C requir e two, three, and
four electrons, respectively. So, these elements represen t the four lightest
4 1 BIOCHEMISTRY
vra 2 He 4.0026 10 Ne 20.183 18 Ar 39.948 36 Kr 83.80 54 Xe 131.30 86 Rn (222) 71 Lu 174.97 103 Lw (257)
vn 9 F 18.9984 17 CI 35.453 35 Br 79.909 53 I 126.9044 85 At (210) 70 Yb 173.04 102 No (254)
vi 8 Ï 15.9994 16 S 32.064 34 Se 78.96 52 Te 127.60 84 Po (210) 69 Tm 168.934 101 Md (256)
v 7 Í 14.0067 15 Ñ 30.9738 33 As 74.9216 51 Sb 121.75 83 Bi 208.980 68 Er 167.26 100 Fm (252)
iv 6 C 12.01115 14 Si 28.086 32 Ge 72.59 50 Sn 118.69 82 Pb 207.19 67 Ho 164.930 99 Es (254) nts
m 5 Â 10.811 13 Al Transition elements Transition elements 26.9815 24 27 25 26 28 29 30 31 Cr Mn Fe Co Ni Zn Cu Ga 54.9380 55.847 51.996 65.37 58.71 58.9332 63.54 69.72 42 43 44 47 49 45 46 48 Mo Tc Ru Pd Ag Cd In Rh 95.94 (99) 101.07 106.4 112.40 107.870 114.82 102.905 77 74 76 78 81 75 79 80 Tl W Os Ir Pt Au Re Hg 204.37 196.967 183.85 190.2 195.09 186.2 192.2 200.59 59 60 61 62 64 65 66 63 Pr Nd Pm Tb Dy Sm Eu Gd 140.907 144.24 (145) 158.924 162.50 151.96 157.25 150.35 91 94 97 92 93 98 95 96 Pa U Np Pu Bk Cf Am Cm (231) 238.03 (237) (249) (242) (243) (245) (249) Atomic number (Z) Element symbol Atomic mass of the naturally occurring isotopic mixture; for the elemethat are naturally radioactive, the numbers in parentheses are mass numbers of the most stable isotopes of these elements. Fig. 1-1 Periodic table of the elements.
23 V 50.942 41 Nb 92.906 73 Ta 180.948 (105) 58 Ce 140.12 90 Th 232.038 — 47-
6 e 8
22 Ti 47.90 40 Zr 91.22 72 Hi 178.49 (104) 57 La 138.91 89 Ac (227) 2 F 55.
ð 4 Be 9.0122 12 Mg 24.312 20 21 Ca Sc 40.08 44.956 39 38 Y Sr 87.62 88.905 57-71 56 Ba 137.34 88 (89-103) Ra t (227) *Lanthanide *Lanthanide *Lanthanide elements elements elements tActinide tActinide tActinide elements elements elements Key: Key: Key:
Group I d 1 Ç 1.00797 3 Li 6.939 11 Na 22.9898 19 Ê 39.102 37 Rb 85.47 55 Cs 132.905 87 Fr (223)
Perio 1 1 1 1 1 1 1 2 2 2 2 2 2 2 3 3 3 3 3 3 3 4 4 4 4 4 4 4 5 5 5 5 5 5 5 6 6 6 6 6 6 6 7 7 7 7 7 7 7
BIOELEMENT S 5
·· ·· ··
Ç· c: · Í : : O :
Fig. 1-2 Schematic electronic arrangement of hydrogen, carbon, nitrogen, and oxygen atom.
Each dot represents an electron in the outer shell.
atoms that requir e the fewest number of additiona l electrons to form stable
electronic configurations .
At this point the reader may well wonder why a stable configuratio n of
electrons is important , how these additiona l electrons are acquired , and
what this all has to do with molecular biology. The stability question is
rather easy to understand , in that biological systems would not be of much
lasting value if they were not stable, but could literally fall apart at any time.
This implies that the atoms composing the systems must be stable. Biological
systems do not last forever; they are constantly undergoin g an evolutionar y
process both on the macroscopic and molecular level. But they do possess
stability of structur e for significant periods of time. In general the additiona l
electrons needed for stability are gained by having two atoms come to-
gether to form a bond in such a manner that both atoms share mutual
electrons so both atoms feel they have the full complement require d to form
a stable configuration . This particula r type of association is called a covalent
bond and is the most frequentl y found bond in biological systems. Naturally ,
the number of other atoms one atom may combine with depends on how
many additiona l electrons are needed to complete a stable set. For instance,
Ç needs one electron to complete its stable configuration ; so two hydrogen
atoms may form a covalent bond with each other where each atom shares
its own single electron plus its neighbor's . Both hydrogens then think they
have a total of two electrons. In this instance the covalent bond is represente d
as H:H or Ç—Ç where the dash represent s two shared electrons. In this
configuratio n both Ç atoms are happy, and we have formed molecular
hydrogen H . Usually, but not always, an atom will be able to donate one
2
electron to be shared in a covalent bond for every atom with which it is
interacting . In this case the covalent bond is known as a single bond.
Another type of bond that is also encountere d is the ionic bond. In this
type of union one atom actually donates an electron to its partner . The first
atom then acquires a positive charge, and the receiving atom becomes
negatively charged. The two atoms are then held together by electrostati c
forces. An example of this is seen in the union between Na and CI, shown
below in equation form:
•Na • Na+ + ·
: ci: + · —•: ci:
·· ··
CP + Na+ • NaCl
