Water: The Medium of Life
Where there’s water, there’s life.
If there is magic on this planet, it is contained in water.
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• Life on Earth is totally dependent on water, and water may be essential to the
existence of life anywhere in the universe.
• Life originated, evolved, and thrives in the seas.
• Organisms invaded and occupied terrestrial and aerial niches, but none gained
true independence from water.
• Typically, organisms are 70% to 90% water.
• Normal metabolic activity can occur only when cells are at least 65% H2O.
• This dependency of life on water is not a simple matter, but it is because of some
unusual chemical and physical properties of H2O.
• Water and its ionization products, hydrogen ions and hydroxide ions, are critical
determinants of the structure and function of many biomolecules, including
amino acids and proteins, nucleotides and nucleic acids, and even phospholipids
and membranes.
• In another essential role, water is an indirect participant—a difference in the
concentration of hydrogen ions on opposite sides of a membrane represents an
energized condition essential to biological mechanisms of energy transformation.
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Water Has Unusual Properties
• Water has a substantially higher boiling point, melting point, heat of vaporization, and
surface tension.
• All of these physical properties are anomalously high for a substance of this molecular
weight that is neither metallic nor ionic.
• These properties suggest that intermolecular forces of attraction between H2O
molecules are high.
• Thus, the internal cohesion of this substance is high.
• Water also has an unusually high dielectric constant, its maximum density is found in
the liquid (not the solid) state, and it has a negative volume of melting (that is, the
solid form, ice, occupies more space than does the liquid form, water).
• It is truly remarkable that so many eccentric properties occur together in this single
substance.
• Explanation for these apparent eccentricities is in the structure of water.
• The key to its intermolecular attractions lies in its atomic constitution.
• The extraordinary ability to form hydrogen bonds is the crucial fact to understanding
its properties.
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Hydrogen Bonding in Water Is Key to Its Properties
• The two hydrogen atoms of water are linked covalently
to oxygen, each sharing an electron pair, to give a
nonlinear arrangement.
• This “bent” structure of the H2O molecule has
enormous influence on its properties.
• If H2O were linear, it would be a nonpolar substance.
• In the bent configuration, however, the electronegative
O atom and the two H atoms form a dipole that renders
the molecule distinctly polar.
• Furthermore, this structure is ideally suited to H-bond
formation.
• Water can serve as both an H donor and an H acceptor in H-bond formation.
• The potential to form four H bonds per water molecule is the source of the
strong intermolecular attractions that endow this substance with its
anomalously high boiling point, melting point, heat of vaporization, and
surface tension.
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• Each molecule of water can form hydrogen bonds with as many as
four other water molecules, producing a highly interconnected
network of molecules.
• Each hydrogen bond is formed when the partially positive-charged
hydrogen of one water molecule becomes aligned next to a partially
negative-charged oxygen atom of another water molecule.
• Because of their extensive hydrogen bonding, water molecules have
an unusually strong tendency to adhere to one another.
• This feature is most evident in the thermal properties of water.
• For example, when water is heated, most of the thermal energy is
consumed in disrupting hydrogen bonds rather than contributing to
molecular motion (which is measured as an increased temperature).
• Similarly, evaporation from the liquid to the gaseous state requires
that water molecules break the hydrogen bonds holding them to their
neighbors, which is why it takes so much energy to convert water to
steam.
• Mammals take advantage of this property when they sweat because
the heat required to evaporate the water is absorbed from the body,
which thus becomes cooler.
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• Hydrogen bonding in water is cooperative.
• That is, an H-bonded water molecule serving as an acceptor is a better H-bond
donor than an unbonded molecule (and an H2O molecule serving as an H-bond
donor becomes a better H-bond acceptor).
• Thus, participation in H bonding by H2O molecules is a phenomenon of mutual
reinforcement.
• The H bonds between neighboring molecules are weak (23 kJ/mol
each) relative to the H—O covalent bonds (420 kJ/mol).
• As a consequence, the hydrogen atoms are situated asymmetrically
between the two oxygen atoms along the O-O axis.
• There is never any ambiguity about which O atom the H atom is
covalently bound to, nor to which O it is H bonded.
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Molecular Interactions in Liquid Water Are Based on H Bonds
• Water molecules are connected by uninterrupted H-bond paths running in every
direction, spanning the whole sample.
• The participation of each water molecule in an average state of H bonding to its neighbors
means that each molecule is connected to every other in a fluid network of H bonds.
• The average lifetime of an H-bonded connection between two H2O molecules in water is
9.5 psec (picoseconds, where 1 psec = 10-12 sec).
• Thus, about every 10 psec, the average H2O molecule moves, reorients, and interacts with
new neighbors.
• In summary, pure liquid water consists of H2O molecules held in a disordered, three-
dimensional network that has a local preference for tetrahedral geometry, yet contains a
large number of strained or broken hydrogen bonds.
• The presence of strain creates a kinetic situation in which H2O molecules can switch H-
bond allegiances; fluidity ensues.
• The sum of all the hydrogen bonds between H2O molecules confers great internal
cohesion on liquid water.
• Extended networks of hydrogen-bonded water molecules also form bridges between
solutes (proteins and nucleic acids, for example) that allow the larger molecules to
interact with each other over distances of several nanometers without physically
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The Structure of Ice
• As water cools, its molecular motion slows and
the molecules move gradually closer to one
another.
• The density of any liquid increases as its
temperature decreases. For most liquids, this
continues as the liquid freezes; the solid state is
denser than the liquid state.
• However, water behaves differently.
• It actually reaches its highest density at
about 4°C.
Between 4°C and 0°C , the density gradually decreases as the hydrogen
bonds begin to form a network characterized by a generally hexagonal
structure with open spaces in the middle of the hexagons
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The Structure of Ice Is Based On H-Bond Formation
• In ice, the hydrogen bonds form a space-filling, three-dimensional
network.
• These bonds are directional and straight; that is, the H atom lies on
a direct line between the two O atoms.
• This linearity and directionality mean that the H bonds in ice are
strong.
• In addition, the directional preference of the H bonds leads to an
open lattice structure.
• For example, if the water molecules are thought to be rigid spheres
centered at the positions of the O atoms in the lattice, then the
observed density of ice is actually only 57% of that expected for a
tightly packed arrangement of such spheres.
• The H bonds in ice hold the water molecules apart.
• Melting involves breaking some of the H bonds that maintain the crystal structure of ice so
that the molecules of water (now liquid) can actually pack closer together.
• Thus, the density of ice is slightly less than that of water. Ice floats, a property of great
importance to aquatic organisms in cold climates.
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• In liquid water, the rigidity of ice is replaced by fluidity and the crystalline
periodicity of ice gives way to spatial homogeneity.
• The H2O molecules in liquid water form a disordered H-bonded network.
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Water Forms Hydrogen Bonds with Polar Solutes
• Hydrogen bonds are not unique to water.
• They readily form between an electronegative atom (the hydrogen acceptor, usually
oxygen or nitrogen) and a hydrogen atom covalently bonded to another
electronegative atom (the hydrogen donor) in the same or another molecule.
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• Hydrogen atoms covalently bonded to carbon atoms do not participate in
hydrogen bonding, because carbon is only slightly more electronegative than
hydrogen and thus the C—H bond is only very weakly polar.
• The distinction explains why butanol (CH3(CH2)2CH2OH) has a relatively high
boiling point of 117 °C, whereas butane (CH3(CH2)2CH3) has a boiling point of only
-0.5 °C.
• Butanol has a polar hydroxyl group and thus can form intermolecular hydrogen
bonds.
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• Uncharged but polar biomolecules such as sugars dissolve readily in water
because of the stabilizing effect of hydrogen bonds between the hydroxyl groups
or carbonyl oxygen of the sugar and the polar water molecules.
• Alcohols, aldehydes, ketones, and compounds containing N—H bonds all form
hydrogen bonds with water molecules and tend to be soluble in water.
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• Hydrogen bonds are strongest when the bonded molecules are oriented to
maximize electrostatic interaction, which occurs when the hydrogen atom and
the two atoms that share it are in a straight line—that is, when the acceptor
atom is in line with the covalent bond between the donor atom and H, putting
the positive charge of the hydrogen ion directly between the two partial negative
charges.
• Hydrogen bonds are thus highly directional and capable of holding two hydrogen-
bonded molecules or groups in a specific geometric arrangement.
• This property of hydrogen bonds confers very precise three-dimensional
structures on protein and nucleic acid molecules, which have many
intramolecular hydrogen bonds.
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Water Interacts Electrostatically with Charged Solutes
• Water is a polar solvent also known as the Universal Solvent.
• It readily dissolves most biomolecules, which are generally charged or polar compounds;
compounds that dissolve easily in water are hydrophilic (Greek, “water-loving”).
• In contrast, nonpolar solvents such as chloroform and benzene are poor solvents for polar
biomolecules but easily dissolve those that are hydrophobic—nonpolar molecules such as
lipids and waxes.
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Water Has a High Dielectric Constant
• Water dissolves salts such as NaCl by hydrating and stabilizing the Na+ and Cl- ions,
weakening the electrostatic interactions between them and thus counteracting their
tendency to associate in a crystalline lattice.
• Similarly water can dissolve other
charged biomolecules, compounds
with functional groups such as
ionized carboxylic acids (—COO-),
protonated amines (—NH3 + ), and
phosphate esters or anhydrides.
• Water readily dissolves such
compounds by replacing solute-
solute hydrogen bonds with
solute-water hydrogen bonds, thus
screening the electrostatic
interactions between solute
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(I)
• Water is effective in screening/shielding the electrostatic interactions between
dissolved ions because it has a high dielectric constant, a physical property that
reflects the number of dipoles in a solvent.
• The strength, or force (F), of ionic interactions in a solution depends on the
magnitude of the charges (Q), the distance between the charged groups (r), and the
dielectric constant, (D), which is dimensionless of the solvent in which the
interactions occur:
• For water at 25 °C, is 78.5, and for the very nonpolar solvent benzene, is 2.3.
• Thus, ionic interactions between dissolved ions are much stronger in less polar
environments.
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(II)
• The attractions between the water molecules interacting with, or hydrating, ions
are much greater than the tendency of oppositely charged ions to attract one
another.
• Water’s ability to surround ions in dipole interactions and diminish their attraction
for each other is a measure of its dielectric constant, D ().
• Thus, ionization in solution depends on the dielectric constant of the solvent;
otherwise, the strongly attracted positive and negative ions would unite to form
neutral molecules.
• The strength of the dielectric constant is related to the force,
F, experienced between two ions of opposite charge
separated by a distance, r, as given in the relationship
where e1 and e2 are the charges on the two ions.
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Dielectric Constants of Some Common Solvents at 25°C
Note that the dielectric constant for water is more than twice that of methanol and
more than 40 times that of hexane.
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• The dielectric constant of water is among the highest of any pure liquid, whereas
those of nonpolar substances, such as hydrocarbons, are relatively small.
• The force between two ions separated by a given distance in nonpolar liquids
such as hexane or benzene is therefore 30 to 40 times greater than that in water.
• Consequently, in nonpolar solvents (low D), ions of opposite charge attract each
other so strongly that they coalesce to form a salt, whereas the much weaker
forces between ions in water solution (high D) permit significant quantities of the
ions to remain separated.
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Water based - Hydrophobic Interactions
• Nonpolar solutes (or nonpolar functional groups on biological macromolecules) do not
readily H bond to H2O, and as a result, such compounds are only sparingly soluble in water.
• The process of dissolving such substances is accompanied by significant reorganization of
the water surrounding the solute like “making structures with particular shapes”.
• Because nonpolar solutes must occupy space,
the random H-bonded network of water must
reorganize to accommodate them.
• At the same time, the water molecules
participate in as many H-bonded interactions
with one another as the temperature permits.
• Consequently, the H-bonded water network
rearranges toward formation of a local cage-like
(clathrate) structure surrounding each solute
molecule, as shown for a long-chain fatty acid in
Figure.
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• This fixed orientation of water molecules around a hydrophobic “solute”
molecule results in a hydration shell.
• As a result of this rearrangement, the molecules of H2O participating in the cage
layer have markedly reduced options for orientation in three-dimensional space.
• Water molecules tend to straddle the nonpolar solute such that two or three
tetrahedral directions (H-bonding vectors) are tangential to the space occupied
by the inert solute.
• “Straddling” allows the water molecules to retain their H-bonding possibilities
because no H-bond donor or acceptor of the H2O is directed toward the caged
solute.
• The water molecules forming these clathrates are involved in highly ordered
structures.
• That is, clathrate formation is accompanied by significant ordering of structure
or negative entropy.
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• Multiple nonpolar molecules
tend to cluster together, because
their joint solvation cage
involves less total surface area
and thus fewer ordered water
molecules than in their separate
cages.
• It is as if the nonpolar molecules
had some net attraction for one
another.
• This apparent affinity of nonpolar structures for one another is called hydrophobic
interactions.
• In actuality, the “attraction” between nonpolar solutes is an entropy-driven process
due to a net decrease in order among the H2O molecules.
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Interaction of water with Amphiphilic Molecules
• Amphipathic compounds contain regions that are polar (or charged) and regions
that are nonpolar.
• When an amphipathic compound is mixed with water, the polar, hydrophilic
region interacts favorably with the solvent and tends to dissolve, but the nonpolar,
hydrophobic region tends to avoid contact with the water.
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• The nonpolar regions of the molecules cluster together to present the smallest hydrophobic
area to the aqueous solvent, and the polar regions are arranged to maximize their interaction
with the solvent (Figure b).
• These stable structures of amphipathic compounds in water, called micelles, may contain
hundreds or thousands of molecules.
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• The forces that hold the nonpolar regions of the molecules together are
called hydrophobic interactions.
• The strength of hydrophobic interactions is not due to any intrinsic
attraction between nonpolar moieties.
• Rather, it results from the system’s achieving greatest thermodynamic
stability by minimizing the number of ordered water molecules required to
surround hydrophobic portions of the solute molecules.
Micelle formation by amphiphilic molecules in
aqueous solution. Because of their negatively
charged surfaces, neighboring micelles repel
one another and thereby maintain a relative
stability in solution
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• Many biomolecules are amphipathic; proteins, pigments, certain vitamins, and the
sterols and phospholipids of membranes all have both polar and nonpolar surface
regions.
• Structures composed of these molecules are stabilized by hydrophobic
interactions among the nonpolar regions.
• Hydrophobic interactions among lipids, and between lipids and proteins, are the
most important determinants of structure in biological membranes.
• Hydrophobic interactions between nonpolar amino acids also stabilize the three-
dimensional structures of proteins.
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Release of ordered water favors formation of
an enzyme-substrate complex.
• Hydrogen bonding between water and polar
solutes also causes an ordering of water
molecules, but the energetic effect is less
significant than with nonpolar solutes.
• Part of the driving force for binding of a polar
substrate (reactant) to the complementary polar
surface of an enzyme is the entropy increase as
the enzyme displaces ordered water from the
substrate, and as the substrate displaces ordered
water from the enzyme surface
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Water Can Ionize to Form H+ and OH—
• Water shows a small but finite tendency to form ions.
• This tendency is demonstrated by the electrical conductivity of pure water, a
property that clearly establishes the presence of charged species (ions).
• Water ionizes because the larger, strongly electronegative oxygen atom strips the
electron from one of its hydrogen atoms, leaving the proton to dissociate.
• Two ions are thus formed: (1) protons or hydrogen ions, H+, and (2) hydroxyl ions,
OH—. Free protons are immediately hydrated to form hydronium ions, H3O+:
H+ + H2O → H3O+
• Because most hydrogen atoms in liquid water are hydrogen bonded to a neighboring
water molecule, this protonic hydration is an instantaneous process and the ion
products of water are H3O+ and OH—:
• This property of water can be demonstrated by a Proton Mobility.
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Proton Mobility
• When an electrical current is passed through an ionic solution, the ions migrate
toward the electrode of opposite polarity at a rate proportional to the electrical
field and inversely proportional to the frictional drag experienced by the ion as it
moves through the solution.
• This latter quantity, varies with the size of the ion.
• However, the ionic mobilities of both H3O+ and OH— are anomalously large
compared to those of other ions.
Ionic Mobilitiesa in H2O at 25°C
aIonic mobility is the distance an ion
moves in 1 s under the influence of an
electric field of 1 V ∙ cm — 1.
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• For H3O+ (the hydronium ion, which is abbreviated
H+; a bare proton has no stable existence in aqueous
solution), this high migration rate results from the
ability of protons to jump rapidly from one water
molecule to another.
• Although a given hydronium ion can physically
migrate through solution in the manner of, say, an
Na+ ion, the rapidity of the proton-jump mechanism
makes the H3O+ ion’s effective ionic mobility much
greater than it otherwise would be (the mean
lifetime of a given H3O+ ion is 10-12 s at 25°C).
• The anomalously high ionic mobility of the OH— ion is likewise
accounted for by the proton-jump mechanism but, in this case, the
apparent direction of ionic migration is opposite to the direction of
proton jumping.
• Proton jumping is also responsible for the observation that acid– base
reactions are among the fastest reactions that take place in aqueous
solutions.
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• The ionization of water can be measured by its electrical conductivity; pure water
carries electrical current as H3O+ migrates toward the cathode and OH— toward
the anode.
• The movement of hydronium and hydroxide ions in the electric field is extremely
fast compared with that of other ions such as Na⁺, K⁺, and Cl⁻.
• This high ionic mobility results from the kind of “proton hopping”.
• No individual proton moves very far through the bulk solution, but a series of
proton hops between hydrogen-bonded water molecules causes the net
movement of a proton over a long distance in a remarkably short time.
• As a result of the high ionic mobility of H (and of OH, which also moves rapidly by
proton hopping, but in the opposite direction), acid-base reactions in aqueous
solutions are exceptionally fast.
• As noted above, proton hopping very likely also plays a role in biological proton-
transfer reactions
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