AST 376/381 Scalo: Planets and Life
Readings on Biochemical Bonding
This document contains three exerpts having to do with chemical bonding,
with an emphasis on simple biomolecules, taken from cell biology or
biochemistry texts. The three selections overlap, but with different-enough
presentations that you should read all three, at least up to the “optional” stop
points that are marked. The fonts are different among the three selections, to
help you tell them apart if you are reading on a computer monitor (that is
why the font size is large), or if your hard copy blows away and randomizes.
If you decide to print them, you should probably reduce the 16 or 14 pt fonts
to 12 pt.
Besides acquainting you with the most important chemical bond
mechanisms, the reading should also be used to gain a familiarity with
biology background that you are probably lacking (I’m assuming the
average student has had no more biology than a distant memory in high
school, although I realize there are one or two exceptions; these readings
should still be accessible to upper division students--later in the semester I
will ask you to look at (but not read!) some more “serious” papers and parts
of books. Remember--most of any problem you have with this reading is
only a terminology problem: There is almost no theoretical underpinning to
anything that is described; biology is primarily a phenomenological (nice
way to say “not much quantitative theory”) science, although that is rapidly
changing as more and more physicists (and even a few researchers in
planetary sciences and astrobiology) try to take over biophysics and make it
boring and so technical that no one can understand it except themselves.
Just (~ halfway) kidding.
Start with the general survey of bonding, including water, from the greatest
of cell biology books, Alberts, Molecular Biology of the Cell. This
presentation also will give you a bit of an introduction to some of the
biomolecules we will be discussing. This exerpt is followed by a little more
advanced reading, sec. 1.3 from Berg et al., Biochemistry. Notice the
emphasis put on the weak noncovalent bonds--keep a list going. Finally, a
section covering mostly the same material, from a book by Lodish et al.,
Biochemistry.
Depending on your background, you will probably encounter unfamiliar
terms. When you do, take advantage of Wikipedia or some other internet
resource to gradually close up terminological holes in your background. Be
careful!
Note: If you want to link to the full text of this book,
go to the free NCBI/NIH bookshelf at:
http://www.ncbi.nlm.nih.gov/entrez/query.fcgi?db=Books
The problem is that you cannot view the figures at the same time as the
text, or easily navigate the text. Here is the best procedure I have found.
Click on the book of your choice (see below). Then click on the table of
contents to see the full contents. Type the name of the chapter or subsection
whose text or images (will be listed to the left) you want in the box to the
upper right corner. You can also search by key word--try typing DNA and stand
back! These are not the most recent editions, but trust me, I have compared
with many other sources in these fields, and these are all excellent and do not
suffer from their publication date except in a few places if you were a
specialist. These books are meant for beginning biology majors (for example,
they are a step less detailed than biochemistry textbooks), so later chapters
are almost certainly too detailed to be of interest to you. I have already made
Word versions of the first few chapters of most of them (and many more),
besides stealing some of their great illustrations for lecture use, so my advice is
to just read what I assign, beginning with this chapter.
Other books you can access here, some of which I have already copied
into Word documents like this one (and will put at the course web site or print
and pass out in class when the time arrives), are:
Berg, J. et al. Biochemistry, 2002, 5e (Sec. 1.3 is the third reading here)
Cooper, G. M. The Cell: A molecular approach 2000, 2e. Great book--I
will have you read a section from this book on biomacromolecules after you
have had a chance to digest the current reading).
Alberts, B. et al. Molecular Biology of the Cell, 4e (First reading here is
from ch.2)
Lodish, H. et al., Molecular Cell Biology--1999 (Part of chapter 2 copied
in here)
Brown, T. A. Genomes, 2e
Griffiths, A. et al. Modern Genetic Analysis 1999.
An important way in which biology textbooks are, out of necessity, superior to
physics and astronomy books, is to have great graphics and refined organization
and writing style, in order to compensate for having no explanations for the
subject matter. You will find, as the deceased baseball player Yogi Berra said
“You can learn a lot by just looking,” so take advantage of the pictures--I have
copied all of the ones that go with the present reading into this document, and
there are links to a few tables of “panels” that I couldn’t include easily.
You may also find that a trip to half-price books will score you a $10-$15
copy of one of these $100-$200 books in pretty good condition. That is how I
started.
Besides the books below, I will be providing links to free access
textbooks and the best tutorials in all the areas we discuss. One way you can
become more comfortable with a new field of knowledge is to get a pile of
books on the subject and just look through it to see how it is organized. Today
we can do that on the internet, except that there is a huge danger of wasting
much time with sites that are low on content, something like the material you
see at most NASA or university or even course websites, meant for the public
(except more press releases at NASA), or high-school teaching aide, not as an
upper-level teaching tool. I have already weeded through 100s of these, but
feel free to try if you want. I concluded that I would know much more had I
settled on a high-level sight and then actually concentrated on the material
instead of spending so much time collecting materials fo this class, especially
the 1000s of images I accumulated. So try to leave the searching to me unless
there is some special topic that you are interested in that we are not covering.
Now on to biochemical bonding...
From
B. Alberts et al., Molecular Biology of the Cell
(online at NIH bookshelf)
The Outermost Electrons Determine How Atoms Interact
To understand how atoms bond together to form the molecules that
make up living organisms, we have to pay special attention to their
electrons. Protons and neutrons are welded tightly to one another
in the nucleus and change partners only under extreme
conditions—during radioactive decay, for example, or in the
interior of the sun or of a nuclear reactor. In living tissues, it is
only the electrons of an atom that undergo rearrangements. They
form the exterior of an atom and specify the rules of chemistry by
which atoms combine to form molecules.
Electrons are in continuous motion around the nucleus (really?
[JMS] , but motions on this submicroscopic scale obey different
laws from those we are familiar with in everyday life. These laws
dictate that electrons in an atom can exist only in certain discrete
states, called orbitals, and that there is a strict limit to the number
of electrons that can be accommodated in an orbital of a given
type—a so-called electron shell. The electrons closest on average
to the positive nucleus are attracted most strongly to it and occupy
the innermost, most tightly bound shell. This shell can hold a
maximum of two electrons. The second shell is farther away from
the nucleus, and its electrons are less tightly bound. This second
shell can hold up to eight electrons. The third shell contains
electrons that are even less tightly bound; it can also hold up to
eight electrons. The fourth and fifth shells can hold 18 electrons
each. Atoms with more than four shells are very rare in biological
molecules.
The electron arrangement of an atom is most stable when all the
electrons are in the most tightly bound states that are possible for
them—that is, when they occupy the innermost shells. Therefore,
with certain exceptions in the larger atoms, the electrons of an
atom fill the orbitals in order—the first shell before the second, the
second before the third, and so on. An atom whose outermost shell
is entirely filled with electrons is especially stable and therefore
chemically unreactive. Examples are helium with 2 electrons, neon
with 2 + 8, and argon with 2 + 8 + 8; these are all inert gases.
Hydrogen, by contrast, with only one electron and therefore only a
half-filled shell, is highly reactive. Likewise, the other atoms found
in living tissues all have incomplete outer electron shells and are
therefore able to donate, accept, or share electrons with each other
to form both molecules and ions (Figure 2-4).
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Figure 2-4. Filled and unfilled electron shells in some common
elements. All the elements commonly found in living organisms
have unfilled outermost shells (red) and can thus participate in
chemical reactions with other atoms. For comparison, some
elements that have only filled shells (yellow) are shown; these are
chemically unreactive.
Because an unfilled electron shell is less stable than a filled one,
atoms with incomplete outer shells have a strong tendency to
interact with other atoms in a way that causes them to either gain
or lose enough electrons to achieve a completed outermost shell.
This electron exchange can be achieved either by transferring
electrons from one atom to another or by sharing electrons between
two atoms. These two strategies generate two types of chemical
bonds between atoms: an ionic bond is formed when electrons are
donated by one atom to another, whereas a covalent bond is
formed when two atoms share a pair of electrons (Figure 2-5).
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Figure 2-5. Comparison of covalent and ionic bonds. Atoms can
attain a more stable arrangement of electrons in their outermost
shell by interacting with one another. An ionic bond is formed
when electrons are transferred from one atom to the other. A
covalent bond is formed when electrons are shared between atoms.
The two cases shown represent extremes; often, covalent bonds
form with a partial transfer (unequal sharing of electrons), resulting
in a polar covalent bond (see Figure 2-43).
Often, the pair of electrons is shared unequally, with a partial
transfer between the atoms; this intermediate strategy results in a
polar covalent bond, as we shall discuss later.
An H atom, which needs only one more electron to fill its shell,
generally acquires it by electron sharing, forming one covalent
bond with another atom; in many cases this bond is polar. The
other most common elements in living cells—C, N, and O, with an
incomplete second shell, and P and S, with an incomplete third
shell (see Figure 2-4)—generally share electrons and achieve a
filled outer shell of eight electrons by forming several covalent
bonds. The number of electrons that an atom must acquire or lose
(either by sharing or by transfer) to attain a filled outer shell is
known as its valence.
The crucial role of the outer electron shell in determining the
chemical properties of an element means that, when the elements
are listed in order of their atomic number, there is a periodic
recurrence of elements with similar properties: an element with,
say, an incomplete second shell containing one electron will
behave in much the same way as an element that has filled its
second shell and has an incomplete third shell containing one
electron. The metals, for example, have incomplete outer shells
with just one or a few electrons, whereas, as we have just seen, the
inert gases have full outer shells.
Ionic Bonds Form by the Gain and Loss of
Electrons
Ionic bonds are most likely to be formed by atoms that have just
one or two electrons in addition to a filled outer shell or are just
one or two electrons short of acquiring a filled outer shell. They
can often attain a completely filled outer electron shell more easily
by transferring electrons to or from another atom than by sharing
electrons. For example, from Figure 2-4 we see that a sodium (Na)
atom, with atomic number 11, can strip itself down to a filled shell
by giving up the single electron external to its second shell. By
contrast, a chlorine (Cl) atom, with atomic number 17, can
complete its outer shell by gaining just one electron. Consequently,
if a Na atom encounters a Cl atom, an electron can jump from the
Na to the Cl, leaving both atoms with filled outer shells. The
offspring of this marriage between sodium, a soft and intensely
reactive metal, and chlorine, a toxic green gas, is table salt (NaCl).
When an electron jumps from Na to Cl, both atoms become
electrically charged ions. The Na atom that lost an electron now
has one less electron than it has protons in its nucleus; it therefore
has a single positive charge (Na+). The Cl atom that gained an
electron now has one more electron than it has protons and has a
single negative charge (Cl-). Positive ions are called cations, and
negative ions, anions. Ions can be further classified according to
how many electrons are lost or gained. Thus sodium and potassium
(K) have one electron to lose and form cations with a single
positive charge (Na+ and K+), whereas magnesium and calcium
have two electrons to lose and form cations with two positive
charges (Mg2+ and Ca2+).
Because of their opposite charges, Na+ and Cl- are attracted to each
other and are thereby held together in an ionic bond. A salt crystal
contains astronomical numbers of Na+ and Cl- (about 2 × 1019 ions
of each type in a crystal 1 mm across) packed together in a precise
three-dimensional array with their opposite charges exactly
balanced (Figure 2-6).
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Figure 2-6. Sodium chloride: an example of ionic bond
formation. (A) An atom of sodium (Na) reacts with an atom of
chlorine (Cl). Electrons of each atom are shown schematically in
their different energy levels; electrons in the chemically reactive
(incompletely filled) shells are red. The reaction takes place with
transfer of a single electron from sodium to chlorine, forming two
electrically charged atoms, or ions, each with complete sets of
electrons in their outermost levels. The two ions with opposite
charge are held together by electrostatic attraction. (B) The product
of the reaction between sodium and chlorine, crystalline sodium
chloride, consists of sodium and chloride ions packed closely
together in a regular array in which the charges are exactly
balanced. (C) Color photograph of crystals of sodium chloride.
Substances such as NaCl, which are held together solely by ionic
bonds, are generally called salts rather than molecules. Ionic bonds
are just one of several types of noncovalent bonds that can exist
between atoms, and we shall meet other examples.
Because of a favorable interaction between water molecules and
ions, ionic bonds are greatly weakened by water; thus many salts
(including NaCl) are highly soluble in water—dissociating into
individual ions (such as Na+ and Cl-), each surrounded by a group
of water molecules. In contrast, covalent bond strengths are not
affected in this way.
Covalent Bonds Form by the Sharing of Electrons
All the characteristics of a cell depend on the molecules it contains.
A molecule is defined as a cluster of atoms held together by
covalent bonds; here electrons are shared between atoms to
complete the outer shells, rather than being transferred between
them. In the simplest possible molecule—a molecule of hydrogen
(H2)—two H atoms, each with a single electron, share two
electrons, which is the number required to fill the first shell. These
shared electrons form a cloud of negative charge that is densest
between the two positively charged nuclei and helps to hold them
together, in opposition to the mutual repulsion between like
charges that would otherwise force them apart. The attractive and
repulsive forces are in balance when the nuclei are separated by a
characteristic distance, called the bond length.
A further crucial property of any bond—covalent or noncovalent—
is its strength. Bond strength is measured by the amount of energy
that must be supplied to break that bond. This is often expressed in
units of kilocalories per mole (kcal/mole), where a kilocalorie is
the amount of energy needed to raise the temperature of one liter of
water by one degree centigrade. Thus if 1 kilocalorie must be
supplied to break 6 × 1023 bonds of a specific type (that is, 1 mole
of these bonds), then the strength of that bond is 1 kcal/mole. An
equivalent, widely used measure of energy is the kilojoule, which
is equal to 0.239 kilocalories.
To get an idea of what bond strengths mean, it is helpful to
compare them with the average energies of the impacts that
molecules are constantly undergoing from collisions with other
molecules in their environment (their thermal, or heat, energy), as
well as with other sources of biological energy such as light and
glucose oxidation (Figure 2-7).
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Figure 2-7. Some energies important for cells. Note that these
energies are compared on a logarithmic scale.
Typical covalent bonds are stronger than the thermal energies by a
factor of 100, so they are resistant to being pulled apart by thermal
motions and are normally broken only during specific chemical
reactions with other atoms and molecules. The making and
breaking of covalent bonds are violent events, and in living cells
they are carefully controlled by highly specific catalysts, called
enzymes. Noncovalent bonds as a rule are much weaker; we shall
see later that they are important in the cell in the many situations
where molecules have to associate and dissociate readily to carry
out their functions.
Whereas an H atom can form only a single covalent bond, the
other common atoms that form covalent bonds in cells—O, N, S,
and P, as well as the all-important C atom—can form more than
one. The outermost shell of these atoms, as we have seen, can
accommodate up to eight electrons, and they form covalent bonds
with as many other atoms as necessary to reach this number.
Oxygen, with six electrons in its outer shell, is most stable when it
acquires an extra two electrons by sharing with other atoms and
therefore forms up to two covalent bonds. Nitrogen, with five outer
electrons, forms a maximum of three covalent bonds, while carbon,
with four outer electrons, forms up to four covalent bonds—thus
sharing four pairs of electrons (see Figure 2-4).
When one atom forms covalent bonds with several others, these
multiple bonds have definite orientations in space relative to one
another, reflecting the orientations of the orbits of the shared
electrons. Covalent bonds between multiple atoms are therefore
characterized by specific bond angles as well as bond lengths and
bond energies (Figure 2-8).
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Figure 2-8. The geometry of covalent bonds. (A) The spatial
arrangement of the covalent bonds that can be formed by oxygen,
nitrogen, and carbon. (B) Molecules formed from these atoms have
a precise three-dimensional structure, as shown here by ball and
stick models for water and propane. A structure can be specified by
the bond angles and bond lengths for each covalent linkage.
The four covalent bonds that can form around a carbon atom, for
example, are arranged as if pointing to the four corners of a regular
tetrahedron. The precise orientation of covalent bonds forms the
basis for the three-dimensional geometry of organic molecules.
There Are Different Types of Covalent Bonds
Most covalent bonds involve the sharing of two electrons, one
donated by each participating atom; these are called single bonds.
Some covalent bonds, however, involve the sharing of more than
one pair of electrons. Four electrons can be shared, for example,
two coming from each participating atom; such a bond is called a
double bond. Double bonds are shorter and stronger than single
bonds and have a characteristic effect on the three-dimensional
geometry of molecules containing them. A single covalent bond
between two atoms generally allows the rotation of one part of a
molecule relative to the other around the bond axis. A double bond
prevents such rotation, producing a more rigid and less flexible
arrangement of atoms (Figure 2-9 and Panel 2-1, pp. 111–112).
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Figure 2-9. Carbon-carbon double bonds and single bonds
compared. (A) The ethane molecule, with a single covalent bond
between the two carbon atoms, illustrates the tetrahedral
arrangement of single covalent bonds formed by carbon. One of
the CH3 groups joined by the covalent bond can rotate relative to
the other around the bond axis. (B) The double bond between the
two carbon atoms in a molecule of ethene (ethylene) alters the
bond geometry of the carbon atoms and brings all the atoms into
the same plane (blue); the double bond prevents the rotation of one
CH2 group relative to the other.
Some molecules share electrons between three or more atoms,
producing bonds that have a hybrid character intermediate between
single and double bonds. The highly stable benzene molecule, for
example, comprises a ring of six carbon atoms in which the
bonding electrons are evenly distributed (although usually depicted
as an alternating sequence of single and double bonds, as shown in
Panel 2-1).
When the atoms joined by a single covalent bond belong to
different elements, the two atoms usually attract the shared
electrons to different degrees. Compared with a C atom, for
example, O and N atoms attract electrons relatively strongly,
whereas an H atom attracts electrons more weakly. By definition, a
polar structure (in the electrical sense) is one with positive charge
concentrated toward one end (the positive pole) and negative
charge concentrated toward the other (the negative pole). Covalent
bonds in which the electrons are shared unequally in this way are
therefore known as polar covalent bonds (Figure 2-10).
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Figure 2-10. Polar and nonpolar covalent bonds. The electron
distributions in the polar water molecule (H2O) and the nonpolar
oxygen molecule (O2) are compared (δ+, partial positive charge; δ-,
partial negative charge).
For example, the covalent bond between oxygen and hydrogen, O-H, or between nitrogen and hydrogen, -N-H, is polar, whereas
that between carbon and hydrogen, -C-H, has the electrons
attracted much more equally by both atoms and is relatively
nonpolar.
Polar covalent bonds are extremely important in biology because
they create permanent dipoles that allow molecules to interact
through electrical forces. Any large molecule with many polar
groups will have a pattern of partial positive and negative charges
on its surface. When such a molecule encounters a second
molecule with a complementary set of charges, the two molecules
will be attracted to each other by permanent dipole interactions that
resemble (but are weaker than) the ionic bonds discussed
previously for NaCl.
An Atom Often Behaves as if It Has a Fixed Radius
When a covalent bond forms between two atoms, the sharing of
electrons brings the nuclei of these atoms unusually close together.
But most of the atoms that are rapidly jostling each other in cells
are located in separate molecules. What happens when two such
atoms touch?
For simplicity and clarity, atoms and molecules are usually
represented in a highly schematic way—either as a line drawing of
the structural formula or as a ball and stick model. However, a
more accurate representation can be obtained through the use of
so-called space-filling models. Here a solid envelope is used to
represent the radius of the electron cloud at which strong repulsive
forces prevent a closer approach of any second, non-bonded
atom—the so-called van der Waals radius for an atom. This is
possible because the amount of repulsion increases very steeply as
two such atoms approach each other closely. At slightly greater
distances, any two atoms will experience a weak attractive force,
known as a van der Waals attraction. As a result, there is a
distance at which repulsive and attractive forces precisely balance
to produce an energy minimum in each atom's interaction with an
atom of a second, non-bonded element (Figure 2-11).
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Figure 2-11. The balance of van der Waals forces between two
atoms. As the nuclei of two atoms approach each other, they
initially show a weak bonding interaction due to their fluctuating
electric charges. However, the same atoms will strongly repel each
other if they are brought too close together. The balance of these
van der Waals attractive and repulsive forces occurs at the
indicated energy minimum.
Depending on the intended purpose, we shall represent small
molecules either as line drawings, ball and stick models, or space
filling models throughout this book. For comparison, the water
molecule is represented in all three ways in Figure 2-12.
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Figure 2-12. Three representations of a water molecule. (A)
The usual line drawing of the structural formula, in which each
atom is indicated by its standard symbol, and each line represents a
covalent bond joining two atoms. (B) A ball and stick model, in
which atoms are represented by spheres of arbitrary diameter,
connected by sticks representing covalent bonds. Unlike (A), bond
angles are accurately represented in this type of model (see also
Figure 2-8). (C) A space-filling model, in which both bond
geometry and van der Waals radii are accurately represented.
When dealing with very large molecules, such as proteins, we shall
often need to further simplify the representation used (see, for
example, Panel 3-2, pp. 138–139).
Water Is the Most Abundant Substance in Cells
Water accounts for about 70% of a cell's weight, and most
intracellular reactions occur in an aqueous environment. Life on
Earth began in the ocean, and the conditions in that primeval
environment put a permanent stamp on the chemistry of living
things. Life therefore hinges on the properties of water.
In each water molecule (H2O) the two H atoms are linked to the O
atom by covalent bonds (see Figure 2-12). The two bonds are
highly polar because the O is strongly attractive for electrons,
whereas the H is only weakly attractive. Consequently, there is an
unequal distribution of electrons in a water molecule, with a
preponderance of positive charge on the two H atoms and of
negative charge on the O (see Figure 2-10). When a positively
charged region of one water molecule (that is, one of its H atoms)
comes close to a negatively charged region (that is, the O) of a
second water molecule, the electrical attraction between them can
result in a weak bond called a hydrogen bond. These bonds are
much weaker than covalent bonds and are easily broken by the
random thermal motions due to the heat energy of the molecules,
so each bond lasts only an exceedingly short time. But the
combined effect of many weak bonds is far from trivial. Each
water molecule can form hydrogen bonds through its two H atoms
to two other water molecules, producing a network in which
hydrogen bonds are being continually broken and formed (Panel 22, pp. 112–113). It is only because of the hydrogen bonds that link
water molecules together that water is a liquid at room
temperature, with a high boiling point and high surface tension—
rather than a gas.
Molecules, such as alcohols, that contain polar bonds and that can
form hydrogen bonds with water dissolve readily in water. As
mentioned previously, molecules carrying plus or minus charges
(ions) likewise interact favorably with water. Such molecules are
termed hydrophilic, meaning that they are water-loving. A large
proportion of the molecules in the aqueous environment of a cell
necessarily fall into this category, including sugars, DNA, RNA,
and a majority of proteins. Hydrophobic (water-hating) molecules,
by contrast, are uncharged and form few or no hydrogen bonds,
and so do not dissolve in water. Hydrocarbons are an important
example (see Panel 2-1, pp. 110–111). In these molecules the H
atoms are covalently linked to C atoms by a largely nonpolar bond.
Because the H atoms have almost no net positive charge, they
cannot form effective hydrogen bonds to other molecules. This
makes the hydrocarbon as a whole hydrophobic—a property that is
exploited in cells, whose membranes are constructed from
molecules that have long hydrocarbon tails, as we shall see in
Chapter 10.
Some Polar Molecules Form Acids and Bases in
Water
One of the simplest kinds of chemical reaction, and one that has
profound significance in cells, takes place when a molecule
possessing a highly polar covalent bond between a hydrogen and a
second atom dissolves in water. The hydrogen atom in such a
molecule has largely given up its electron to the companion atom
and so exists as an almost naked positively charged hydrogen
nucleus—in other words, a proton (H + ). When the polar molecule
becomes surrounded by water molecules, the proton is attracted to
the partial negative charge on the O atom of an adjacent water
molecule and can dissociate from its original partner to associate
instead with the oxygen atoms of the water molecule to generate a
hydronium ion (H 3
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O
(Figure 2-13A).
+)
Figure 2-13. Acids in water. (A) The reaction that takes place
when a molecule of acetic acid dissolves in water. (B) Water
molecules are continuously exchanging protons with each other to
form hydronium and hydroxyl ions. These ions in turn rapidly
recombine to form water molecules.
The reverse reaction also takes place very readily, so one has to
imagine an equilibrium state in which billions of protons are
constantly flitting to and fro from one molecule in the solution to
another.
Substances that release protons to form H3O+ when they dissolve in
water are termed acids. The higher the concentration of H3O+, the
more acidic the solution. H3O+ is present even in pure water, at a
concentration of 10-7 M, as a result of the movement of protons
from one water molecule to another (Figure 2-13B). By tradition,
the H3O+ concentration is usually referred to as the H+
concentration, even though most H+ in an aqueous solution is
present as H3O+. To avoid the use of unwieldy numbers, the
concentration of H+ is expressed using a logarithmic scale called
the pH scale, as illustrated in Panel 2-2 (pp. 112–113). Pure water
has a pH of 7.0.
Because the proton of a hydronium ion can be passed readily to
many types of molecules in cells, altering their character, the
concentration of H3O+ inside a cell (the acidity) must be closely
regulated. Molecules that can give up protons will do so more
readily if the concentration of H3O+ in solution is low and will tend
to receive them back if the concentration in solution is high.
The opposite of an acid is a base. Just as the defining property of
an acid is that it donates protons to a water molecule so as to raise
the concentration of H3O+ ions, the defining property of a base is
that it raises the concentration of hydroxyl (OH-) ions—which are
formed by removal of a proton from a water molecule. Thus
sodium hydroxide (NaOH) is basic (the term alkaline is also used)
because it dissociates in aqueous solution to form Na+ ions and OHions. Another class of bases, especially important in living cells,
are those that contain NH2 groups. These groups can generate OHby taking a proton from water: -NH2 + H2O → -NH3+ + OH-.
Because an OH- ion combines with a H3O+ ion to form two water
molecules, an increase in the OH- concentration forces a decrease
in the concentration of H3O+, and vice versa. A pure solution of
water contains an equally low concentration (10-7 M) of both ions;
it is neither acidic nor basic and is therefore said to be neutral with
a pH of 7.0. The inside of cells is kept close to neutrality.
[Stop here if you do not have time to finish--I would prefer that you read all
three selections up to the places marked like this, and leave the somewhat
more difficult or detailed material that follows. In that case skip to the
selection by Berg et al.]
[I included this material on noncovalent interactions although we
are concentrating on covalent bonding before we turn to what will
be our main target, the noncovalent bonds. So read it if you are
eager.]
Four Types of Noncovalent Interactions Help Bring
Molecules Together in Cells
In aqueous solutions, covalent bonds are 10 to 100 times stronger
than the other attractive forces between atoms, allowing their
connections to define the boundaries of one molecule from
another. But much of biology depends on the specific binding of
different molecules to each other. This binding is mediated by a
group of noncovalent attractions that are individually quite weak,
but whose bond energies can sum to create an effective force
between two separate molecules. We have already introduced three
of these noncovalent forces: ionic bonds, hydrogen bonds and van
der Waals attractions. In Table 2-2, the strengths of these three
types of bonds are compared to that of a typical covalent bond,
both in the presence and the absence of water. Because of their
fundamental importance in all biological systems, we shall
summarize their properties here.
Ionic bonds. These are purely electrostatic attractions between
oppositely charged atoms. As we saw for NaCl, these forces
are quite strong in the absence of water. However, the polar
water molecules cluster around both fully charged ions and
polar molecules that contain permanent dipoles (Figure 214).
on water molecules orient to
reduce the affinity of
oppositely charged ions or
polar groups for each other.
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Figure 2-14. How the dipoles
This greatly reduces the potential attractiveness of these charged
species for each other (see Table 2-2).
Hydrogen bonds. The structure of a typical hydrogen bond is
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illustrated in Figure 2-15.
Figure 2-15. Hydrogen bonds. (A) Ball- and-stick model of a
typical hydrogen bond. The distance between the hydrogen
and the oxygen atom here is less than the sum of their van der
Waals radii, indicating a partial sharing of electrons. (B) The
most common hydrogen bonds in cells.
This bond represents a special form of polar interaction in which
an electropositive hydrogen atom is partially shared by two
electronegative atoms. Its hydrogen can be viewed as a
proton that has partially dissociated from a donor atom,
allowing it to be shared by a second acceptor atom. Unlike a
typical electrostatic interaction, this bond is highly
directional—being strongest when a straight line can be
drawn between all three of the involved atoms. As already
discussed, water weakens these bonds by forming competing
hydrogen-bond interactions with the involved molecules (see
Table 2-2).
van der Waals attractions. The electron cloud around any
nonpolar atom will fluctuate, producing a flickering dipole.
Such dipoles will transiently induce an oppositely polarized
flickering dipole in a nearby atom. This interaction generates
an attraction between atoms that is very weak. But since
many atoms can be simultaneously in contact when two
surfaces fit closely, the net result is often significant. These
so-called van der Waals attractions are not weakened by
water (see Table 2-2).
The fourth effect that can play an important part in bringing
molecules together in water is a hydrophobic force. This force is
caused by a pushing of nonpolar surfaces out of the hydrogenbonded water network, where they would physically interfere with
the highly favorable interactions between water molecules.
Because bringing two nonpolar surfaces together reduces their
contact with water, the force is a rather nonspecific one.
Nevertheless, we shall see in Chapter 3 that hydrophobic forces are
central to the proper folding of protein molecules.
Panel 2-3 provides an overview of the four types of interactions
just described. (You should be able to link to it from this
document.)
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Figure 2-16. How two macro-molecules with complementary
surfaces can bind tightly to one another through noncovalent
interactions. In this schematic illustration, plus and minus are
used to mark chemical groups that can form attractive interactions
when paired.
And Figure 2-16 above illustrates, in a schematic way, how many
such interactions can sum to hold together the matching surfaces of
two macromolecules, even though each interaction by itself would
be much too weak to be effective.
[Chap. 2 of Alberts continues with “A cell is formed from carbon
compounds, which introduces a little organic chemistry before going into the
major classes of macromolecules: Save for another reading collection, and
keep this close to bonding.]
_____________________________________________
NEXT section of this document: A parallel but slightly more extended
discussion of “Chemical bonds in biochemistry,” Sec. 1.3 in Berg et al.,
Biochemistry. Note the emphasis on the various weak noncovalent forces
and the roles they play. Keep thinking--what sort of structures or
phenomena could they give rise to in nonbiological mesoscale
macromolecular assemblies? (Note: This is NOT a detailed treatment of
each kind of molecule--nearly each topic, like protein folding, is treated in
separate chapters later in the book.)
Berg et al. Biochemistry
1.3. Chemical Bonds in Biochemistry
[I have included the several pages discussing the noncovalent
interactions which are of primary interest to us because it is a
short and painless introduction. I have omitted the (longer)
sections on this topic from the other two reading selections. By
the time you read this we should have covered the interactions
from a physics point of view, so this will be complementary.]
The essence of biological processes—the basis of
the uniformity of living systems—is in its most
fundamental sense molecular interactions; in other
words, the chemistry that takes place between
molecules. Biochemistry is the chemistry that takes
place within living systems. To truly understand
biochemistry, we need to understand chemical
bonding. We review here the types of chemical bonds
that are important for biochemicals and their
transformations.
The strongest bonds that are present in
biochemicals are covalent bonds, such as the bonds
that hold the atoms together within the individual
bases shown in Figure 1.3.
Figure 1.3. Watson-Crick Base Pairs. Adenine pairs
with thymine (A-T), and guanine with cytosine (G-
C). The dashed lines represent hydrogen bonds.
A covalent bond is formed by the sharing of a pair
of electrons between adjacent atoms. A typical
carbon-carbon (C-C) covalent bond has a bond length
of 1.54 Å and bond energy of 85 kcal mol-1 (356 kJ
mol-1). Because this energy is relatively high,
considerable energy must be expended to break
covalent bonds. More than one electron pair can be
shared between two atoms to form a multiple
covalent bond. For example, three of the bases in
Figure 1.4
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Figure 1.4. Base-Pairing in DNA. The base-pairs A-T
(blue) and C-G (red) are shown overlaid. The
Watson-Crick base-pairs have the same overall size
and shape, allowing them to fit neatly within the
double helix.
include carbon-oxygen (C=O) double bonds. These
bonds are even stronger than C-C single bonds, with
energies near 175 kcal mol-1 (732 kJ mol-1).
For some molecules, more than one pattern of
covalent bonding can be written. For example,
benzene can be written in two equivalent ways
called resonance structures. Benzene's true
structure is a composite of its two resonance
structures. A molecule that can be written as
several resonance structures of approximately equal
energies has greater stability than does a molecule
without multiple resonance structures. Thus,
because of its resonance structures, benzene is
unusually stable.
Chemical reactions entail the
breaking and forming of covalent
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the course of a reaction can be
depicted by curved arrows, a
method of representation called
“arrow pushing.” Each arrow
represents an electron pair.
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1.3.1. Reversible Interactions of Biomolecules Are
Mediated by Three Kinds of Noncovalent Bonds
Readily reversible, noncovalent molecular
interactions are key steps in the dance of life.
Such weak, noncovalent forces play essential roles
in the faithful replication of DNA, the folding of
proteins into intricate three-dimensional forms,
the specific recognition of substrates by enzymes,
and the detection of molecular signals. Indeed, all
biological structures and processes depend on the
interplay of noncovalent interactions as well as
covalent ones. The three fundamental noncovalent
bonds are electrostatic interactions, hydrogen
bonds, and van der Waals interactions. They differ
in geometry, strength, and specificity.
Furthermore, these bonds are greatly affected in
different ways by the presence of water. Let us
consider the characteristics of each:
1. Electrostatic interactions. An electrostatic
interaction depends on the electric charges on
atoms. The energy of an electrostatic
interaction is given by Coulomb's law:
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where E is the energy, q1 and q2 are the charges on
the two atoms (in units of the electronic charge),
r is the distance between the two atoms (in
angstroms), D is the dielectric constant (which
accounts for the effects of the intervening
medium), and k is a proportionality constant (k =
332, to give energies in units of kilocalories per
mole, or 1389, for energies in kilojoules per
mole). Thus, the electrostatic interaction between
two atoms bearing single opposite charges separated
by 3 Å in water (which has a dielectric constant of
80) has an energy of 1.4 kcal mol-1 (5.9 kJ mol-1).
2. Hydrogen bonds. Hydrogen bonds are relatively
weak interactions, which nonetheless are
crucial for biological macromolecules such as
DNA and proteins. These interactions are also
responsible for many of the properties of water
that make it such a special solvent. The
hydrogen atom in a hydrogen bond is partly
shared between two relatively electronegative
atoms such as nitrogen or oxygen. The hydrogenbond donor is the group that includes both the
atom to which the hydrogen is more tightly
linked and the hydrogen atom itself, whereas
the hydrogen-bond acceptor is the atom less
tightly linked to the hydrogen atom (Figure
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1.9).
Figure 1.9. Hydrogen Bonds that Include Nitrogen
and Oxygen Atoms. The positions of the partial
charges (δ+ and δ-) are shown.
Hydrogen bonds are fundamentally electrostatic
interactions. The relatively electronegative atom
to which the hydrogen atom is covalently bonded
pulls electron density away from the hydrogen atom
so that it develops a partial positive charge (δ+).
Thus, it can interact with an atom having a partial
negative charge (δ-) through an electrostatic
interaction.
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Hydrogen bonds are much weaker than covalent bonds.
They have energies of 1–3 kcal mol-1 (4–13 kJ mol-1)
compared with approximately 100 kcal mol-1 (418 kJ
mol-1) for a carbon-hydrogen covalent bond. Hydrogen
bonds are also somewhat longer than are covalent
bonds; their bond distances (measured from the
hydrogen atom) range from 1.5 to 2.6 Å; hence,
distances ranging from 2.4 to 3.5 Å separate the
two nonhydrogen atoms in a hydrogen bond. The
strongest hydrogen bonds have a tendency to be
approximately straight, such that the hydrogen-bond
donor, the hydrogen atom, and the hydrogen-bond
acceptor lie along a straight line.
3. van der Waals interactions. The basis of a van
der Waals interaction is that the distribution of
electronic charge around an atom changes with time.
At any instant, the charge distribution is not
perfectly symmetric. This transient asymmetry in
the electronic charge around an atom acts through
electrostatic interactions to induce a
complementary asymmetry in the electron
distribution around its neighboring atoms. The
resulting attraction between two atoms increases as
they come closer to each other, until they are
separated by the van der Waals contact distance
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(Figure 1.10).
Figure 1.10. Energy of a van der Waals Interaction
as Two Atoms Approach One Another. The energy is
most favorable at the van der Waals contact
distance. The energy rises rapidly owing to
electron- electron repulsion as the atoms move
closer together than this distance.
At a shorter distance, very strong repulsive forces
become dominant because the outer electron clouds
overlap.
Energies associated with van der Waals interactions
are quite small; typical interactions contribute
from 0.5 to 1.0 kcal mol-1 (from 2 to 4 kJ mol-1)
per atom pair. When the surfaces of two large
molecules come together, however, a large number of
atoms are in van der Waals contact, and the net
effect, summed over many atom pairs, can be
substantial.
1.3.2. The Properties of Water Affect the Bonding
Abilities of Biomolecules
Weak interactions are the key means by which
molecules interact with one another—enzymes with
their substrates, hormones with their receptors,
antibodies with their antigens. The strength and
specificity of weak interactions are highly
dependent on the medium in which they take place,
and the majority of biological interactions take
place in water. Two properties of water are
especially important biologically:
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1. Water is a polar molecule. The water molecule is
bent, not linear, and so the distribution of charge
is asymmetric. The oxygen nucleus draws electrons
away from the hydrogen nuclei, which leaves the
region around the hydrogen nuclei with a net
positive charge. The water molecule is thus an
electrically polar structure.
2. Water is highly cohesive. Water molecules
interact strongly with one another through hydrogen
bonds. These interactions are apparent in the
structure of ice (Figure 1.11).
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Figure 1.11. Structure of Ice. Hydrogen bonds
(shown as dashed lines) are formed between water
molecules.
Networks of hydrogen bonds hold the structure
together; simi-lar interactions link molecules in
liquid water and account for the cohesion of liquid
water, although, in the liquid state, some of the
hydrogen bonds are broken. The highly cohesive
nature of water dramatically affects the
interactions between molecules in aqueous solution.
What is the effect of the properties of water on
the weak interactions discussed in Section 1.3.1?
The polarity and hydrogen-bonding capability of
water make it a highly interacting molecule. Water
is an excellent solvent for polar molecules. The
reason is that water greatly weakens electrostatic
forces and hydrogen bonding between polar molecules
by competing for their attractions. For example,
consider the effect of water on hydrogen bonding
between a carbonyl group and the NH group of an
amide.
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A hydrogen atom of water can replace the amide
hydrogen atom as a hydrogen-bond donor, whereas the
oxygen atom of water can replace the carbonyl
oxygen atom as a hydrogen-bond acceptor. Hence, a
strong hydrogen bond between a CO group and an NH
group forms only if water is excluded.
The dielectric constant of water is 80, so water
diminishes the strength of electrostatic
attractions by a factor of 80 compared with the
strength of those same interactions in a vacuum.
The dielectric constant of water is unusually high
because of its polarity and capacity to form
oriented solvent shells around ions. These oriented
solvent shells produce electric fields of their
own, which oppose the fields produced by the ions.
Consequently, the presence of water markedly
weakens electrostatic interactions between ions.
The existence of life on Earth depends critically
on the capacity of water to dissolve a remarkable
array of polar molecules that serve as fuels,
building blocks, catalysts, and information
carriers. High concentrations of these polar
molecules can coexist in water, where they are free
to diffuse and interact with one another. However,
the excellence of water as a solvent poses a
problem, because it also weakens interactions
between polar molecules. The presence of water-free
microenvironments within biological systems largely
circumvents this problem. We will see many examples
of these specially constructed niches in protein
molecules. Moreover, the presence of water with its
polar nature permits another kind of weak
interaction to take place, one that drives the
folding of proteins (Section 1.3.4) and the
formation of cell boundaries (Section 12.4).
The essence of these interactions, like that of all
interactions in biochemistry, is energy. To
understand much of biochemistry—bond formation,
molecular structure, enzyme catalysis—we need to
understand energy. Thermodynamics provides a
valuable tool for approaching this topic. We will
revisit this topic in more detail when we consider
enzymes (Chapter 8) and the basic concepts of
metabolism (Chapter 14).
[Stop here if you do not have time to finish--I
would prefer that you read all three selections up
to the places marked like this, and leave the
somewhat more difficult or detailed material that
follows. In that case skip to the selection by
Lodish et al.
________________________________________
[I included this material on thermodynamical view
of bond formation in case you are interested or if
you have never had thermodynamics, as an
opportunity to see this point of view. However
sec. 1.3.3 and 1.3.4 below (on protein folding) are
optional reading, and you won’t be tested on it and
no assigned homework will depend on it.]
1.3.3. Entropy and the Laws of Thermodynamics
The highly structured, organized nature of living
organisms is apparent and astonishing. This
organization extends from the organismal through
the cellular to the molecular level. Indeed,
biological processes can seem magical in that the
well-ordered structures and patterns emerge from
the chaotic and disordered world of inanimate
objects. However, the organization visible in a
cell or a molecule arises from biological events
that are subject to the same physical laws that
govern all processes—in particular, the laws of
thermodynamics.
How can we understand the creation of order out of
chaos? We begin by noting that the laws of
thermodynamics make a distinction between a system
and its surroundings. A system is defined as the
matter within a defined region of space. The matter
in the rest of the universe is called the
surroundings. The First Law of Thermodynamics
states that the total energy of a system and its
surroundings is constant. In other words, the
energy content of the universe is constant; energy
can be neither created nor destroyed. Energy can
take different forms, however. Heat, for example,
is one form of energy. Heat is a manifestation of
the kinetic energy associated with the random
motion of molecules. Alternatively, energy can be
present as potential energy, referring to the
ability of energy to be released on the occurrence
of some process. Consider, for example, a ball held
at the top of a tower. The ball has considerable
potential energy because, when it is released, the
ball will develop kinetic energy associated with
its motion as it falls. Within chemical systems,
potential energy is related to the likelihood that
atoms can react with one another. For instance, a
mixture of gasoline and oxygen has much potential
energy because these molecules may react to form
carbon dioxide and release energy as heat. The
First Law requires that any energy released in the
formation of chemical bonds be used to break other
bonds, be released as heat, or be stored in some
other form.
Another important thermodynamic concept is that of
entropy. Entropy is a measure of the level of
randomness or disorder in a system. The Second Law
of Thermodynamics states that the total entropy of
a system and its surroundings always increases for
a spontaneous process. At first glance, this law
appears to contradict much common experience,
particularly about biological systems. Many
biological processes, such as the generation of a
well-defined structure such as a leaf from carbon
dioxide gas and other nutrients, clearly increase
the level of order and hence decrease entropy.
Entropy may be decreased locally in the formation
of such ordered structures only if the entropy of
other parts of the universe is increased by an
equal or greater amount.
An example may help clarify the application of the
laws of thermodynamics to a chemical system.
Consider a container with 2 moles of hydrogen gas
on one side of a divider and 1 mole of oxygen gas
on the other (Figure 1.12).
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Figure 1.12. From Order to Disorder. The
spontaneous mixing of gases is driven by an
increase in entropy.
If the divider is removed, the gases will
intermingle spontaneously to form a uniform
mixture. The process of mixing increases entropy as
an ordered arrangement is replaced by a randomly
distributed mixture.
Other processes within this system can decrease the
entropy locally while increasing the entropy of the
universe. A spark applied to the mixture initiates
a chemical reaction in which hydrogen and oxygen
combine to form water:
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If the temperature of the system is held constant,
the entropy of the system decreases because 3 moles
of two differing reactants have been combined to
form 2 moles of a single product. The gas now
consists of a uniform set of indistinguishable
molecules. However, the reaction releases a
significant amount of heat into the surroundings,
and this heat will increase the entropy of the
surrounding molecules by increasing their random
movement. The entropy increase in the surroundings
is enough to allow water to form spontaneously from
hydrogen and oxygen (Figure 1.13).
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Figure 1.13. Entropy Changes. When hydrogen and
oxygen combine to form water, the entropy of the
system is reduced, but the entropy of the universe
is increased owing to the release of heat to the
surroundings.
The change in the entropy of the surroundings will
be proportional to the amount of heat transferred
from the system and inversely proportional to the
temperature of the surroundings, because an input
of heat leads to a greater increase in entropy at
lower temperatures than at higher temperatures. In
biological systems, T [in kelvin (K), absolute
temperature] is assumed to be constant. If we
define the heat content of a system as enthalpy
(H), then we can express the relation linking the
entropy (S) of the surroundings to the transferred
heat and temperature as a simple equation:
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The total entropy change is given by the expression
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Substituting equation 1 into equation 2 yields
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Multiplying by -T gives
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The function -TΔS has units of energy and is
referred to as free energy or Gibbs free energy,
after Josiah Willard Gibbs, who developed this
function in 1878:
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The free-energy change, ΔG, will be used throughout
this book to describe the energetics of biochemical
reactions.
Recall that the Second Law of Thermodynamics states
that, for a reaction to be spontaneous, the entropy
of the universe must increase. Examination of
equation 3 shows that the total entropy will
increase if and only if
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Rearranging gives TΔSsystem > ΔH, or entropy will
increase if and only if
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w
In other words, the free-energy change must be
negative for a reaction to be spontaneous. A
negative free-energy change occurs with an increase
in the overall entropy of the universe. Thus, we
need to consider only one term, the free energy of
the system, to decide whether a reaction can occur
spontaneously; any effects of the changes within
the system on the rest of the universe are
automatically taken into account.
1.3.4. Protein Folding Can Be Understood in Terms
of Free-Energy Changes
The problem of protein folding illustrates the
utility of the concept of free energy. ConsIder a
system consisting of a solution of unfolded protein
molecules in aqueous solution (Figure 1.14).
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Figure 1.14. Protein Folding. Protein folding
entails the transition from a disordered mixture of
unfolded molecules to a relatively uniform solution
of folded protein molecules.
Each unfolded protein molecule can adopt a unique
conformation, so the system is quite disordered and
the entropy of the collection of molecules is
relatively high. Yet, protein folding proceeds
spontaneously under appropriate conditions. Thus,
entropy must be increasing elsewhere in the system
or in the surroundings. How can we reconcile the
apparent contradiction that proteins spontaneously
assume an ordered structure, and yet entropy
increases?
The entropy decrease in the system on folding
is not as large as it appears to be, because of the
properties of water. Molecules in aqueous solution
interact with water molecules through the formation
of hydrogen and ionic interactions. However, some
molecules (termed nonpolar molecules) cannot
participate in hydrogen or ionic interactions. The
interactions of nonpolar molecules with water are
not as favorable as are interactions between the
water molecules themselves. The water molecules in
contact with these nonpolar surfaces form “cages”
around the nonpolar molecule, becoming more well
ordered (and, hence, lower in entropy) than water
molecules free in solution. As two such nonpolar
molecules come together, some of the water
molecules are released, and so they can interact
freely with bulk water (Figure 1.15).
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Figure 1.15. The Hydrophobic Effect. The
aggregation of nonpolar groups in water leads to an
increase in entropy owing to the release of water
molecules into bulk water.
Hence, nonpolar molecules have a tendency to
aggregate in water because the entropy of the water
is increased through the release of water
molecules. This phenomenon, termed the hydrophobic
effect, helps promote many biochemical processes.
How does the hydrophobic effect favor protein
folding? Some of the amino acids that make up
proteins have nonpolar groups. These nonpolar amino
acids have a strong tendency to associate with one
another inside the interior of the folded protein.
The increased entropy of water resulting from the
interaction of these hydrophobic amino acids helps
to compensate for the entropy losses inherent in
the folding process.
Hydrophobic interactions are not the only means of
stabilizing protein structure. Many weak bonds,
including hydrogen bonds and van der Waals
interactions, are formed in the protein-folding
process, and heat is released into the surroundings
as a consequence. Although these interactionws
replace interactions with water that take place in
the unfolded protein, the net result is the release
of heat to the surroundings and thus a negative
(favorable) change in enthalpy for the system.
The folding process can occur when the combination
of the entropy associated with the hydrophobic
effect and the enthalpy change associated with
hydrogen bonds and van der Waals interactions makes
the overall free energy negative.
Lodish, Molecular Cell Biology
Ch. 2 Chemical Foundations
2.1. Covalent Bonds [I have left out sec. 2.2, a great discussion
of noncovalent bonds; we will read that separately.]
Covalent bonds, which hold the atoms within an individual molecule
together, are formed by the sharing of electrons in the outer atomic
orbitals. The distribution of shared as well as unshared electrons in
outer orbitals is a major determinant of the three-dimensional
shape and chemical reactivity of molecules. For instance, as we
learn in Chapter 3, the shape of proteins is crucial to their function
and their interactions with small molecules. In this section, we
discuss important properties of covalent bonds and describe the
structure of carbohydrates to illustrate how the geometry of bonds
determines the shape of small biological molecules.
Each Atom Can Make a Defined Number of Covalent Bonds
Electrons move around the nucleus of an atom in clouds called
orbitals, which lie in a series of concentric shells, or energy levels;
electrons in outer shells have more energy than those in inner
shells. Each shell has a maximum number of electrons that it can
hold. Electrons fill the innermost shells of an atom first; then the
outer shells. The energy level of an atom is lowest when all of its
orbitals are filled, and an atom’s reactivity depends on how many
electrons it needs to complete its outermost orbital. In most cases,
in order to fill the outermost orbital, the electrons within it form
covalent bonds with other atoms. A covalent bond thus holds two
atoms close together because electrons in their outermost orbitals
are shared by both atoms.
Most of the molecules in living systems contain only six different
atoms: hydrogen, carbon, nitrogen, phosphorus, oxygen, and sulfur.
The outermost orbital of each atom has a characteristic number of
electrons:
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These atoms readily form covalent bonds with other atoms and
rarely exist as isolated entities. As a rule, each type of atom forms
a characteristic number of covalent bonds with other atoms.
For example, a hydrogen atom, with one electron in its outer shell,
forms only one bond, such that its outermost orbital becomes filled
with two electrons. A carbon atom has four electrons in its
outermost orbitals; it usually forms four bonds, as in methane
(CH4), in order to fill its outermost orbital with eight electrons. The
single bonds in methane that connect the carbon atom with each
hydrogen atom contain two shared electrons, one donated from the
C and the other from the H, and the outer (s) orbital of each H
atom is filled by the two shared electrons:
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Nitrogen and phosphorus each have five electrons in their outer
shells, which can hold up to eight electrons. Nitrogen atoms can
form up to four covalent bonds. In ammonia (NH3), the nitrogen
atom forms three covalent bonds; one pair of electrons around the
atom (the two dots on the right) are in an orbital not involved in a
covalent bond:
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In the ammonium ion (NH4+), the nitrogen atom forms four covalent
bonds, again filling the outermost orbital with eight electrons:
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Phosphorus can form up to five covalent bonds, as in phosphoric
acid (H3PO4). The H3PO4 molecule is actually a “resonance hybrid,”
a structure between the two forms shown below in which
nonbonding electrons are shown as pairs of dots:
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In the resonance hybrid on the right, one of the electrons from the
P=O double bond has accumulated around the O atom, giving it a
net negative charge and leaving the P atom with a net positive
charge. The resonance hybrid on the left, in which the P atom
forms the maximum five covalent bonds, has no charged atoms.
Esters of phosphoric acid form the backbone of nucleic acids, as
discussed in Chapter 4; phosphates also play key roles in cellular
energetics (Chapter 16) and in the regulation of cell function
(Chapters 13 and 20).
The difference between the bonding patterns of nitrogen and
phosphorus is primarily due to the relative sizes of the two atoms:
the smaller nitrogen atom has only enough space to accommodate
four bonding pairs of electrons around it without creating
destructive repulsions between them, whereas the larger sphere of
the phosphorus atom allows more electron pairs to be arranged
around it without the pairs being too close together.
Both oxygen and sulfur contain six electrons in their outermost
orbitals. However, an atom of oxygen usually forms only two
covalent bonds, as in molecular oxygen, O2:
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Primarily because its outermost orbital is larger than that of
oxygen, sulfur can form as few as two covalent bonds, as in
hydrogen sulfide (H2S), or as many as six, as in sulfur trioxide (SO3)
or sulfuric acid (H2SO4):
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Esters of sulfuric acid are important constituents of the
proteoglycans that compose part of the extracellular matrix
surrounding most animal cells (Chapter 22).
The Making or Breaking of Covalent Bonds Involves Large Energy
Changes
Covalent bonds tend to be very stable because the energies
required to break or rearrange them are much greater than the
thermal energy available at room temperature (25 °C) or body
temperature (37 °C). For example, the thermal energy at 25 °C is
less than 1 kilocalorie per mole (kcal/mol), whereas the energy
required to breakQuickTime™
a C—C bond
in ethane is about 83 kcal/mol:
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where ΔH represents the difference in the total energy of all of the
bonds (the enthalpy) in the reactants and in the products.*The
positive value indicates that an input of energy is needed to cause
the reaction, and that the products contain more energy than the
reactants. The high energy needed for breakage of the ethane bond
means that at room temperature (25 °C) well under 1 in 1012 ethane
molecules exists as a pair of ·CH3 radicals. The covalent bonds in
biological molecules have ΔH values similar to that of the C—C bond
in ethane (Table 2-1).
Table 2-1. The Energy Required to Break Some Important
Covalent Bonds Found in Biological Molecules*
Type of Bond
SINGLE BOND
O—H
H—H
P—O
C—H
C—O
C—C
S—H
C—N
C—S
N—O
S—S
Energy
(kcal/mol)
110
104
100
99
84
83
81
70
62
53
51
Type of Bond
Energy
(kcal/mol)
DOUBLE BOND
C=O
C=N
C=C
P=O
170
147
146
120
TRIPLE BOND
C≡O
195
* Note that double and triple bonds are stronger than single
bonds.
* A calorie is defined as the amount of thermal energy required to
heat 1 cm3 of water by 1 °C from 14 °C to 15 °C. Many biochemistry
textbooks use the joule (J), but the two units can be interconverted
quite readily (1 cal = 4.184 J). The energy changes in chemical
reactions, such as the making or breaking of chemical bonds, are
measured in kilocalories per mole in this book (1 kcal = 1000 cal).
One mole of any substance is the amount that contains 6.02 × 1023
items of that substance, which is known as Avogadro’s number.
Thus, one can speak of a mole of photons, or 6.02 × 1023 photons.
The weight of a mole of a substance in grams (g) is the same as its
molecular weight. For example, the molecular weight of water is
18, so the weight of 1 mole of water, or 6.02 × 1023 water
molecules, is 18 g.
Covalent Bonds Have Characteristic Geometries
When two or more atoms form covalent bonds with another central
atom, these bonds are oriented at precise angles to one another.
The angles are determined by the mutual repulsion of the outer
electron orbitals of the central atom. These bond angles give each
molecule its characteristic shape (Figure 2-2).
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In methane, for example, the central carbon atom is bonded to
four hydrogen atoms, whose positions define the four points of a
tetrahedron, so that the angle between any two bonds is 109.5°.
Like methane, the ammonium ion also has a tetrahedral shape. In
these molecules, each bond is a single bond, a single pair of
electrons shared between two atoms. When two atoms share two
pairs of electrons — for example, when a carbon atom is linked to
only three other atoms — the bond is a double bond:
Qu ic kTi me™ a nd a
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are n ee de d to s ee th is pi ctu re .
In this case, the carbon atom and all three atoms linked to it lie in
the same plane (Figure 2-3).
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Atoms connected by a double bond cannot rotate freely about
the bond axis, while those in a single bond generally can. The
rigid planarity imposed by double bonds has enormous significance
for the shape of large biological molecules such as proteins and
nucleic acids. (In triple bonds, two atoms share six electrons. These
are rare in biological molecules.) [Conversely, atoms connected
by a single bond can rotate freely and allow the molecules to
adopt a variety of shapes which are called conformational
isomers. This will become an important theme as we approach
more complex molecules.]
All outer electron orbitals, whether or not they are involved in
covalent bond formation, contribute to the properties of a
molecule, in particular to its shape. For example, the outer shell of
the oxygen atom in a water molecule has two pairs of nonbonding
electrons; the two pairs of electrons in the H—O bonds and the two
pairs of nonbonding electrons form an almost perfect tetrahedron.
However, the orbitals of the nonbonding electrons have a high
electron density and thus tend to repel each other, compressing the
angle between the covalent H—O—H bonds to 104.5° rather than
the 109.5° in a tetrahedron (see Figure 2-2). [More on molecular
shape provided separately.]
Electrons Are Shared Unequally in Polar Covalent Bonds
In a covalent bond, one or more pairs of electrons are shared
between two atoms. In certain cases, the bonded atoms exert
different attractions for the electrons of the bond, resulting in
unequal sharing of the electrons. The power of an atom in a
molecule to attract electrons to itself, called electronegativity, is
measured on a scale from 4.0 (for fluorine, the most
electronegative atom) to a hypothetical zero (Figure 2-4).
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Figure 2-4. Electronegativity values of main-group elements in
the periodic table.
Atoms located to the upper right tend to have high
electronegativity, fluorine being the most electronegative.
Elements with low electronegativity values, such as the metals
lithium, sodium, and potassium, are often called electropositive.
The electronegativities of several atoms abundant in biological
molecules differ enough that they form polar covalent bonds (e.g.,
O—H, N—H) or ionic bonds (e.g., Na+Cl−). Because the inert gases
(He, Ne, etc.) have complete outer shells of electrons, they neither
attract nor donate electrons, rarely form covalent bonds, and have
no electronegativity values.
Knowing the electronegativity of two atoms allows us to predict
whether a covalent bond can form between them; if the differences
in electronegativity are considerable — as in sodium and chloride —
an ionic bond, rather than a covalent bond, will form. This type of
interaction is discussed in a later section.
In a covalent bond in which the atoms either are identical or have
the same electronegativity, the bonding electrons are shared
equally. Such a bond is said to be nonpolar. This is the case for C—C
and C—H bonds. However, if two atoms differ in electronegativity,
the bond is said to be polar. One end of a polar bond has a partial
negative charge (δ−), and the other end has a partial positive charge
(δ+). In an O—H bond, for example, the oxygen atom, with an
electronegativity of 3.4, attracts the bonded electrons more than
does the hydrogen atom, which has an electronegativity of 2.2. As a
result, the bonding electrons spend more time around the oxygen
atom than around the hydrogen. Thus the O—H bond possesses an
electric dipole, a positive charge separated from an equal but
opposite negative charge. We can think of the oxygen atom of the
O—H bond as having, on average, a charge of 25 percent of an
electron, with the H atom having an equivalent positive charge. The
dipole moment of the O—H bond is a function of the size of the
positive or negative charge and the distance separating the charges.
[Later we will see that dipole moments, whether permanent or
induced or fluctuating around zero, play the major role in
controlling the formation of the macromolecules and
supramolecular assemblies or macromolecules that we label as
grains, PAH assemblies, aerosols, nanoparticles, and biomolecules.]
In a water molecule both hydrogen atoms are on the same side of
the oxygen atom. As a result, the side of the molecule with the two
H atoms has a slight net positive charge, whereas the other side has
a slight net negative charge. Because of this separation of positive
and negative charges, the entire molecule has a net dipole moment
(Figure 2-5). Some molecules, such as the linear molecule CO2, have
two polar bonds:
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Figure 2-5. The water molecule has two polar O—H bonds and a
net dipole moment. The symbol δ represents a partial charge (a
weaker charge than the one on an electron or a proton), and each
of the polar H—O bonds has a dipole moment. The net dipole
moment of the molecule is determined by the sizes and directions
of the dipole moments of each of the bonds.
Because the dipole moments of the two C=O bonds point in opposite
directions, they cancel each other out, resulting in a molecule
without a net dipole moment.
NOTE: Water is so essential, along with carbon, that we will devote
many pages to understanding its amazing behavior and interactions
with other molecules. I thought this was completely overblown a
year ago, but no longer.
Asymmetric Carbon Atoms Are Present in Most Biological
Molecules
A carbon (or any other) atom bonded to four dissimilar atoms or
groups is said to be asymmetric. The bonds formed by an
asymmetric carbon atom can be arranged in three dimensional
space in two different ways, producing molecules that are mirror
images of each other. Such molecules are called optical isomers, or
stereoisomers. One isomer is said to be right-handed and the other
left-handed, a property called chirality. Most molecules in cells
contain at least one asymmetric carbon atom, often called a chiral
carbon atom. The different stereoisomers of a molecule usually
have completely different biological activities.
Amino Acids
Except for glycine, all amino acids, the building blocks of the
proteins, have one chiral carbon atom, called the α carbon, or Cα,
which is bonded to four different atoms or groups of atoms. In the
amino acid alanine, for instance, this carbon atom is bonded to —
NH2, —COOH, —H, and —CH3 (Figure 2-6).
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Figure 2-6. Stereoisomers of the amino acid alanine. The
asymmetric α carbon is black. Although the chemical properties of
such optical isomers are identical, their biological activities are
distinct.
By convention, the two mirror-image structures are called the D
(dextro) and the L (levo) isomers of the amino acid. The two
isomers cannot be interconverted without breaking a chemical
bond. With rare exceptions, only the L forms of amino acids are
found in proteins. We discuss the properties of amino acids and the
covalent peptide bond that links them into long chains in Chapter 3.
Stop here if your time or interest is too limited.
_____________________________________________
Carbohydrates
The three-dimensional structures of carbohydrates provide another
excellent example of the structural and biological importance of
chiral carbon atoms, even in simple molecules. A carbohydrate is
constructed of carbon (carbo-) plus hydrogen and oxygen (-hydrate,
or water). The formula for the simplest carbohydrates — the
monosaccharides, or simple sugars — is (CH2O)n, where n equals 3,
4, 5, 6, or 7. All monosaccharides contain hydroxyl (—OH) groups
and either an aldehyde or a keto group:
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In the linear form of D-glucose (C6H12O6), the principal source of
energy for most cells in higher organisms, carbon atoms 2, 3, 4, and
5 are asymmetric (Figure 2-7, top). If the hydrogen atom and the
hydroxyl group attached to carbon atom 2 (C2) were interchanged,
the resulting molecule would be a different sugar, D-mannose, and
could not be converted to glucose without breaking and making
covalent bonds. Enzymes can distinguish between this single point
of difference.
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Figure 2-7. Three alternative configurations of D-glucose. The
ring forms, shown as Haworth projections, are generated from the
linear molecule by reaction of the aldehyde at carbon 1 with the
hydroxyl on carbon 5 or carbon 4.
D-Glucose can exist in three different forms: a linear structure and
two different hemiacetal ring structures (see Figure 2-7). If the
aldehyde group on carbon 1 reacts with the hydroxyl group on
carbon 5, the resulting hemiacetal, D-glucopyranose, contains a sixmember ring. Similarly, condensation of the hydroxyl group on
carbon 4 with the aldehyde group results in the formation of Dglucofuranose, a hemiacetal containing a five-member ring.
Although all three forms of D-glucose exist in biological systems,
the pyranose form is by far the most abundant.
The planar depiction of the pyranose ring shown in Figure 2-7 is
called a Haworth projection. When a linear molecule of D-glucose
forms a pyranose ring, carbon 1 becomes asymmetric, so two
stereoisomers (called anomers) of D-glucopyranose are possible.
The hydroxyl group attached to carbon 1 “points” down (below the
plane of projection) in α-D-glucopyranose, as shown in Figure 2-7,
and points up (above the plane of projection) in the β anomer. In
aqueous solution the α and β anomers readily interconvert
spontaneously; at equilibrium there is about one-third α anomer
and two-thirds β, with very little of the open-chain form. Because
enzymes can distinguish between the α and β anomers of D-glucose,
these forms have specific biological roles.
Most biologically important sugars are six-carbon sugars, or hexoses,
that are structurally related to D-glucose. Mannose, as noted, is
identical with glucose except for the orientation of the substituents
on carbon 2. In Haworth projections of the pyranose forms of
glucose and mannose, the hydroxyl group on carbon 2 of glucose
points downward, whereas that on mannose points upward (Figure
2-8). Similarly, galactose, another hexose, differs from glucose only
in the orientation of the hydroxyl group on carbon 4.
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The Haworth projection is an oversimplification be-cause the actual
pyranose ring is not planar. Rather, sugar molecules adopt a
conformation in which each of the ring carbons is at the center of a
tetrahedron, just like the carbon in methane (see Figure 2-2). The
preferred conformation of pyranose structures is the chair (Figure
2-9). In this conformation, the bonds going from a ring carbon to
nonring atoms may take two directions: axial (perpendicular to the
ring) and equatorial (in the plane of the ring).
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Figure 2-9. Chair conformations of glucose, mannose, and
galactose in their pyranose forms. The chair is the most stable
conformation of a six-membered ring. (In an alternative form,
called the boat, both carbon 1 and carbon 4 lie above the plane of
the ring.) The four bonds at each of the ring carbon atoms are
tetrahedral. As shown in the generalized pyranose ring at the top
left, bonds that extend nearly perpendicular to the plane of the
ring are said to be axial (a); those that extend nearly parallel to the
ring are said to be equatorial (e). In α-D-glucopyranose, all the
hydroxyl groups except the one bonded to carbon 1 are equatorial.
In α-D-mannopyranose, the hydroxyl groups bonded to carbons 1
and 2 are axial. In α-D-galactopyranose, the hydroxyl groups bonded
to carbons 1 and 4 are axial. Note that, as in Figure 2-8, the
hydroxyl groups with orientations different from those in glucose
are highlighted.
The L isomers of sugars are virtually unknown in biological systems
except for L-fucose. One of the unsolved mysteries of molecular
evolution is why only D isomers of sugars and L isomers of amino
acids were utilized, and not the chemically equivalent L sugars and
D amino acids.
α and β Glycosidic Bonds Link Monosaccharides
In addition to the monosaccharides discussed above, two common
disaccharides, lactose and sucrose, occur naturally (Figure 2-10). A
disaccharide consists of two monosaccharides linked together by a
C—O—C bridge called a glycosidic bond. The disaccharide lactose is
the major sugar in milk; sucrose is a principal product of plant
photosynthesis and is refined into common table sugar.
In the formation of any glycosidic bond, the carbon 1 atom of one
sugar molecule reacts with a hydroxyl group of another. As in the
formation of most biopolymers, the linkage is accompanied by the
loss of water. In principle, a large number of different glycosidic
bonds can be formed between two sugar residues. Glucose could be
bonded to fructose, for example, by any of the following linkages:
α(1 → 1), α(1 → 2), α(1 → 3), α(1 → 4), α(1 → 6), β(1 → 1), β(1 →
2), β(1 → 3), β(1 → 4), or β(1 → 6), where α or β specifies the
conformation at carbon 1 in glucose and the number following the
arrow indicates the fructose carbon to which the glucose is bound.
Only the α(1 → 2) linkage occurs in sucrose because of the
specificity of the enzyme (the biological catalyst) for the linking
reaction.
Glycosidic linkages also join chains of monosaccharides into longer
polymers, called polysaccharides, some of which function as
reservoirs for glucose. The most common storage carbohydrate in
animal cells is glycogen, a very long, highly branched polymer of
glucose units linked together mainly by α(1 → 4) glycosidic bonds.
As much as 10 percent by weight of the liver can be glycogen. The
primary storage carbohydrate in plant cells, starch, also is a glucose
polymer with α(1 → 4) linkages. It occurs in two forms, amylose,
which is unbranched, and amylopectin, which has some branches. In
contrast to glycogen and starch, some polysaccharides, such as
cellulose, have structural and other nonstorage functions. An
unbranched polymer of glucose linked together by β(1 → 4)
glycosidic bonds, cellulose is the major constituent of plant cell
walls and is the most abundant organic chemical on earth. Because
of the different linkages between the glucose units, cellulose forms
long rods, whereas glycogen and starch form coiled helices. Human
digestive enzymes can hydrolyze α(1 → 4) glycosidic bonds, but not
β(1 → 4) bonds, between glucose units; for this reason humans can
digest starch but not cellulose. The synthesis and utilization of
these polysaccharides are described in later chapters.
SUMMARY
 Covalent bonds, which bind the atoms composing a molecule in
a fixed orientation, consist of pairs of electrons shared by two
atoms. Relatively high energies are required to break them
(50 – 200 kcal/mol).
 In covalent bonds between unlike atoms that differ in
electronegativity, the bonding electrons are distributed
unequally. In such polar bonds, one end has a partial positive
charge and the other end has a partial negative charge (see
Figure 2-5).
 Most molecules in cells contain at least one chiral (asymmetric)
carbon atom, which is bonded to four dissimilar atoms. Such
molecules can exist as optical isomers, designated D and L,
which have identical chemical properties but completely
different biological activities. In biological systems, nearly all
amino acids are L isomers and nearly all sugars are D isomers.
 Glucose and other hexoses can exist in three forms: an openchain linear structure, a six-member (pyranose) ring, and a
five-member (furanose) ring (see Figure 2-7). In biological
systems, the pyranose form of D- glucose predominates. The
two possible stereoisomers of D-glucopyranose (the α and β
anomers) differ in the orientation of the hydroxyl group
attached to carbon 1.
 Glycosidic bonds link carbon 1 of one monosaccharide to a
hydroxyl group on another sugar, leading to formation of
disaccharides and polysaccharides. Many different glycosidic bonds
are theoretically possible between two sugar residues, but the
enzymes that make and break
these bonds are specific for the α
or β anomer of one sugar and a particular hydroxyl group
on
the other.