CHEMISTRY OF WATER
Water makes up 70 to 90 percent of the weight of most life forms, and is often regarded as an inert
liquid. Actually, it is a highly reactive substance with unusual properties that make it essential for
life. The strong intermolecular forces in liquid water are caused by the specific distribution of
electrons within the molecule. Each of the hydrogen atoms shares a pair of electrons with the
oxygen atom, and the resulting bond angles gives the resulting compound an electrical asymmetry.
The high electronegativity of oxygen tends to withdraw the single electrons from the hydrogen
atom, leaving the hydrogen nuclei comparatively bare. As a result, the hydrogen atoms have a
partial positive charge, and the oxygen atom has a partial negative charge. Thus, although there is
no net charge to the molecule, it is an electric dipole.
In solution, electrostatic attraction occurs between the partial negative charge on the oxygen atom
of one water molecule and the partial positive charges on adjacent hydrogen molecules. This
results in a charge redistribution that greatly enhances their interaction (hydrogen bonding).
Because of the nearly tetrahedral arrangement of the electrons around the oxygen atom, each water
molecule is potentially capable of bonding with four neighboring water molecules, conferring great
internal cohesion to liquid water.
Hydrogen bonds are extremely weak as compared with covalent bonds; the bond energy of these
bonds in liquid water is about 4.5 kcal mol-1 as compared with 110 kcal mol-1 for covalent H-O
bonds. In the most common form of ice (ice I) each water molecule is hydrogen bonded with
exactly four neighboring water molecules in a regular lattice arrangement. In liquid water, each
water molecule is hydrogen bonded at any given time to about 3.6 water molecules. Liquid water
therefore has considerable short-range order but no long-range structure. The small difference in
hydrogen bonding between ice and water is due to the high rate at which hydrogen bonds are made
and broken in liquid water. The half-life of each hydrogen bond is only about 10-11 s.
Water is also a much better solvent than most common liquids. Many crystalline salts dissolve in
water readily. For example, the water easily dissolves crystalline sodium chloride because the
strong electrostatic attraction between water dipoles and the Na+ and Cl- ions greatly exceeds the
tendency of Na+ and Cl- to attract each other, leading to very stable hydrated ions.
Water also tends to oppose the electrostatic attraction between positive and negative ions. The
strength of this tendency is given by the dielectric constant D, and defined by Coulomb's Law:
F = kq1q2/Dr2
where:
F = the force between two electrical charges (q1 and q2)
r = distance separating the two charges
k = proportionality constant
Water has a dielectric constant of 80, about 40 times higher than benzene and hexane. Thus, the
attractive force between Na+ and Cl- in benzene would be 40x stronger than in water. The second
class of substances readily dissolved by water includes the nonionic but polar compounds,
including sugars, simple alcohols, aldehydes, and ketones. Their solubility is due to the propensity
of the hydroxyl groups of sugars and the carbonyl oxygen of aldehydes and ketones to form
hydrogen bonds with water molecules.
The presence of dissolved solutes (as in biological fluids) causes the structure and properties of
liquid water to change. For example, hydrated Na+ and Cl- ions create regions that are more highly
ordered and regular in structure. In .15 M salt solutions (isotonic with body fluids), dissolved ions
are only separated by about 1.9 nm. Since these hydrated ions are 0.5 to 0.7 nm in diameter when
hydrated and a tetrahedral cluster of five water molecules is about 0.5 nm in diameter, these
dissolved salts significantly alter the normal structure of liquid water.
The effect of solutes on the solvent are manifest in the colligative properties of that solution, which
depend on the number of solute particles per unit volume of solvent. Typically, solutes depress the
freezing point, elevate the boiling point and depress the vapor pressure of water. They also create
the property of osmotic pressure.
Because the mass of the hydrogen atom is so small and its electron is drawn so tightly to the oxygen
atom, there is a tendency for a hydrogen ion to dissociate from the oxygen atom to which it is
covalently bonded in one water molecule and reassociate with an oxygen atom in an adjacent water
molecule (proton jumping). This reaction yields a hydronium ion (H3O+) and a hydroxide ion (OH). Hydronium ions are hydrated to form H9O4+ ions, as well as more hydrated, but less stable,
forms. In a liter of pure water at 25 C at any given time there is only 1.0 x 10-7 mol of each ion, as
shown by electrical conductivity measurements. Thus, water can be classified as an extremely
weak electrolyte. The ionic mobilities of both hydronium and hydroxide ions are extremely high, an
important property for acid-base reactions as well as for maintaining electrical neutrality in
solutions.
The concentration of water in pure water is very high (55.5 M) and affected only slightly by the
dissociation into hydronium and hydroxide ions. Since ion concentrations are about 1.0 x 10-7 each
at 25 C, the ion product for water, or:
Kw = [H3O+][OH-]
-14
equals 1.0 x 10 . In acidic solutions, therefore, hydronium ions will exceed hydroxide ions, and
the reverse will be true for basic solutions. Sorenson devised the pH system to simplify the
determination of concentrations, and succeeded in confusing the entire issue for nearly 100 years.
pH = log10(1/[H+]) = -log10[H+]
in a precisely neutral solution at 25 C, pH = 7.0. Since the pH scale is logarithmic, a change of 1
pH unit is equal to a 10-fold change in hydrogen-ion concentration.
ACID-BASE DEFINITIONS
Bronsted-Lowry (1923) - An acid is a hydrogen ion donor, containing one or more bound
hydrogen ions. A base is a hydrogen ion acceptor, containing one or more hydrogen ion
binding sites. An aprote is a type of molecule that will neither donate nor accept hydrogen
ions. Proton is used synonymously with hydrogen ion. In aqueous solutions, the hydrogen
ions are hydrated and the "free" hydrogen ions exist mainly as H3O+.
Arrhenius (1887) - applies only to aqueous solutions, while Bronsted also includes non-aqueous
media. Within aqueous media the concepts are similar, with an Arrhenius base (hydroxyl
ion donor) being a Bronsted base because the hydroxyl ion is also a hydrogen ion acceptor.
Lewis (1938) - expands on the Bronsted definition. While "Lewis bases" differ only formally from
"Bronsted bases", the definition of an acid according to Lewis is much more
comprehensive.
Van Slyke (1921) - was the first to employ the anion-cation approach to acid-base balance. This is
in complete disagreement with the Bronsted definition: Anions are bases according to
Bronsted, but acids according to Van Slyke; and most cations (especially cations of the
alkali metals) are under all conditions aprotes according to Bronsted, but bases according to
Van Slyke. Despite this, the concepts acidic and basic as well as the concepts of acidosis
and alkalosis have been consistent in the medical literature, despite indiscriminate use of
both definitions. This paradox is easily explained if the assumption is made that an anion
must always be accompanied by a cation (the law of electroneutrality). Therefore it has
been assumed that an anion acts as an acid only if it is accompanied by a hydrogen ion
(therefore fulfilling the assumptions of a Bronsted acid). If the anion is accompanied by a
cation other than H+, a salt is formed. Similarly, a cation acts as a base only when
accompanied by a hydroxyl ion and thus essentially is a Bronsted base. Thus, the apparent
disagreements between Van Slyke and Bronsted have been minimized.
Using Van Slyke's definitions, a description of the acid-base relationships requires a full
consideration of all the electrolytes, while the Bronsted terminology only requires
consideration of the hydrogen ion. Using the Bronsted terminology an addition of a strong
acid is completely equivalent to the removal of a strong base with respect to acid-base
conditions. Van Slykes definitions were considered to be "relics of the days before the
dissociation theory".
Originally, the term acid was an attribute of taste, associated with aqueous solutions. With
the advent of Arrhenius definitions of the dissociation theory, it became clear that the
degree of acidity was related to the hydrogen ions. The acidity of the solution was usually
determined by titration, until Sorenson (1909) pointed out that two kinds of acidity had to
be distinguished; actual acidity and titratable acidity. Thus, an aqueous solution of 0.1 mole
HCl per liter and an aqueous solution of 0.1 mole of lactic acid per mole show the same
acidity when titrated with a strong base, yet the concentration of free hydrogen ions is
widely different (100 mmol/L and 3.7 mmol/L, respectively). The terms acidity and
titratable acidity thus came to define different values.
Stewart (1981) brought the first new consideration of the basis of acid-base balance in several
decades, and approached acid-base behavior using a systems analysis approach. This
quantitative approach first recognizes [H+] as a dependent variable. This distinction is
crucial to the understanding of this approach. Dependent variables are internal to the
system, and are dependent for change on the activity of the independent variables.
Dependent variables are necessarily correlated (since they are all dependent on the changes
in the independent variables), but are NOT causally related to each other. H+ movements in
or out of a solution do not explain changes in [H+].
Stewart brought back the consideration of electrolytes as critical factors in determining acidbase behavior. In extracellular fluid, the most concentrated strong ions are Na+ and Cl-.
Although K+, Ca2+, Mg2+ and SO42- are strong ions, they are present at much lower
concentrations than Na+ and Cl-. Lactic acid is highly dissociative and behaves as an
organic strong ion in plasma and extracellular fluid. It may be quantitatively important
under anoxic or hypoxic conditions. Intracellularly, the most important strong ion is K+.
Thus, the concept of electrolytes as critical to acid-base behavior was updated by Stewart to
the consideration of strong ion differences [SID]. This is an expansion of the old Van Slyke
definition, and incorporates aspects of Bronsted as well in the inclusion of dissociation
constants. However, the approach by which the Stewart concept is based is radically
different than any of its predecessors. [SID] can be defined as the sum of all strong cation
concentrations in solution minus the sum of all strong anion concentrations in the solution.
If strong anions predominate, [SID] will be negative. [H+] in strong ion solutions is
determined by the value of [SID]. So is [OH-], but differently. If [SID] is negative, [H+] is
larger than [OH-], and the solution is acidic. The basic premise of Bronsted-Lowry is that
H+ behaves as a strong ion, when in fact it only behaves this way in acidic solutions. Since
physiological fluids are almost never acidic, hydrogen ions almost never behave as strong
ions. Any changes in [SID] alter both [OH-] and [H+], so that it is not possible to change
only the [H+] (and therefore the pH) of a solution. In most physiological solutions, the
magnitudes of [OH-] and its changes are at least 10 times greater than the corresponding
changes for [H+]. The magnitudes of changes for the other ions in physiological solutions
are generally several orders of magnitude above that. The implications of these facts on pH
dependent phenomena are of tremendous importance. In strong ion solutions, all one can
conclude from pH-dependent data is that the data is dependent on [SID]. Measurements of
pH in alkaline (or physiological) solutions can never be an indication of how much H+ has
moved into or out of a solution.
In physiological solutions, the three independent variables which ultimately determine the
value of [H+] are [SID], Pco2, and the total concentration of weak acids (i.e. proteins in
plasma). It should therefore be obvious that during the immediate postnatal period, with
large fluctuations in all three of the independent variables discussed above, acid-base
balance is much more difficult to interpret than in the adult, when Pco2 and the total
concentration of plasma proteins is relatively stable. Looking at adult values, [H+] increases
as Pco2 increases, so that CO2 is acting as an acid. [H+] also increases as [SID] increases.
By this definition, respiratory and metabolic acidosis and alkalosis become more clearly
defined. Respiratory would be due only to a change in Pco2, but metabolic would be as a
result of [SID] or plasma proteins. If pH and Pco2 values are available for a plasma sample,
and if plasma proteins are near normal, then the effective [SID] can be mathematically
determined by the position of the [H+]-Pco2 point on Figure 7 or the pH-Pco2 point on
Figure 8. If plasma electrolytes are available, so that the inorganic component of [SID]
([Na+] + [K+] - [Cl-]) is known, then the difference between the measured inorganic
components and the effective [SID] value from the graph would determine the value for
unidentified strong ions present in the plasma. During lactic acidosis, for example, this
difference should be equal to the lactate concentration in the plasma, while in diabetic
acidosis, the difference should agree with the total keto acids present. In the newborn, the
changes in plasma proteins would alter the position, shape, and slope of these curves and
therefore affect the interpretation of the acid-base data.
The effects of temperature on dependent ion concentrations is important, as the dissociation
constants for both the dependent and independent variables. Physiologically, interpretation
of temperature effects is made more complicated by changes in metabolic and
cardiovascular function which alter the independent variables, especially Pco2 and [SID] by
large amounts.
The major mechanism by which the kidney can control plasma [H+] is by differential
removal of Na+, K+, and Cl- from plasma to urine. Since plasma and urine are separate
solutions, the [H+] of urine is irrelevant to the [H+] of plasma, and the conventional
treatment of urine as a vehicle for the removal of H+ from the body is incorrect. If the
kidneys are to lower the [H+] of plasma, they must excrete more Cl- than Na+ from it, so as
to raise plasma [SID]. Synthesizing NH3 and excreting it as NH4+ can have no effect on
plasma [H+], although it may be very important in keeping the urine [H+] from rising too
high due to the excessive Cl- (and therefore the decreased [SID] value) in the urine.
Similarly, movement of HCO3- into or out of kidney tubules cannot alter plasm [H+], so
long as carbonic anhydrase is effective. In some cases, as in depressed carbonic anhydrase
activity, equilibration with Pco2 is never achieved, complicating the calculations necessary
even further.
Stomach acid presents a similar situation, in that oxyntic cells can only alter acid content by
adjusting [SID], not by pumping H+. They may excrete Cl-, or absorb Na+ or K+, or both,
but the end result must be a highly negative [SID] (-100 Meq/L).
The regulation of pH can be summarized in the following four statements:
1) The [H+], pH, [HCO3-], etc., of each body fluid is determined by the Pco2, [SID]
and total weak acid content (i.e. plasma proteins) in that fluid.
2) Most membranes separating body fluids are not normally permeable to proteins,
so that protein interactions are not normally significant in adult situations,
therefore limiting independent variables of interest under these conditions to
two.
3) The Pco2 of most body fluids is determined, either directly or indirectly, by the
lungs (or the placenta in fetal conditions), and is the point of the most rapid
regulation of acid-base behavior.
4) Strong ions are moved between body fluids through membranes, and the resulting
changes in [SID] values provide the major mechanism for acid-base
interactions between body fluids. Changes in plasma strong ion activity
produced by kidney and gut activity are communicated to other body fluids,
with appropriate modifications by Donnan effects, since capillary
membranes are generally permeable to strong ions but not proteins. The
kidney and the gut therefore are the primary controllers of body fluid [SID]
values, and are involved in chronic control of acid-base alterations.