Transcript Water, acids, bases and buffers

Water, acids, bases and buffers
Water him no get enemy. Fela kuti
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THE BIOLOGICAL IMPORTANCE OF WATER
Water is an ideal biological solvent: it dissolves and
transports a wide variety of organic and inorganic molecules
Water influences the conformations of many biomolecules
Water is a reactant or a product in many reactions
Water removes excess heat from the body
Total body water is roughly 50 to 60% of body weight in
adults and 75% of body weight in children
Because fat has relatively little water associated with it,
obese people tend to have a lower percentage of body water
than thin people, women tend to have a lower percentage
than men, and older people have a lower percentage than
younger people
Approximately 40% of the total body water is intracellular
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and 60% extracellular
• The extracellular water includes the fluid in plasma (blood after
the cells have been removed) and interstitial water (the fluid in
the tissue spaces, lying between cells)
• Transcellular water is a small, specialized portion of
extracellular water that includes saliva, gastrointestinal
secretions, ,urine, sweat, cerebrospinal fluid,….
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Fluid compartments in the body based on an average 70kg man
• The unique properties of water are due to its structure
Hydrogen bonding
• A water molecule is an irregular, slightly skewed
tetrahedron with oxygen at its center
• The 1050 angle between the hydrogens differs slightly from
the ideal tetrahedral angle, 109.50
• Water is a dipole, a molecule with electrical charge
distributed asymmetrically about its structure
• The strongly electronegative oxygen atom pulls electrons
away from the hydrogen nuclei, leaving them with a partial
positive charge (δ+), while its two unshared electron pairs
constitute a region of local negative charge (δ-)
• The hydrogen nuclei on one molecule of water interacts
with the lone pair on an oxygen atom on another water 4
The tetrahedral structure of the
water molecule
Hydrogen bonding between
water molecules
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• Hydrogen bonding favors the self-association of water
molecules into ordered arrays
• Hydrogen bonding profoundly influences the physical
properties of water and accounts for its exceptionally high
viscosity, surface tension and boiling point
• On average, each molecule in liquid water associates through
hydrogen bonds with 3.4 others; these bonds are both
relatively weak and transient, with a half-life of pico seconds
• In ice, each water molecule forms a hydrogen bond with four
other water molecules, giving rise to a crystalline tetrahedral
arrangement
• Rupture of a hydrogen bond in liquid water requires only
about 4.5 kcal/mol, less than 5% of the energy required to
rupture a covalent O—H bond
• However, The cumulative effect of many hydrogen bonds is
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equivalent to the stabilizing effect of covalent bonds
• Hydrogen bonding enables water to dissolve many organic
biomolecules that contain functional groups which can
participate in hydrogen bonding
• The oxygen atoms of aldehydes, ketones, and amides, for
example, provide lone pairs of electrons that can serve as
hydrogen acceptors; alcohols and amines can serve both as
hydrogen acceptors and as donors of unshielded hydrogen
atoms for formation of hydrogen bonds
Polar groups participating in
hydrogen bonding
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The interaction of water with charged solutes
Water has a high dielectric constant; it greatly decreases the
force of attraction between charged and polar species relative
to water-free environments with lower dielectric constants
Water’s strong dipole and high dielectric constant enable water
to dissolve large quantities of charged compounds such as salts
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
As a salt dissolves, the ions leaving the crystal lattice acquire far
greater freedom of motion
The resulting increase in entropy of the system is largely
responsible for the ease of dissolving salts such as NaCl in
water
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• In thermodynamic terms, formation of the solution occurs with
a favorable free-energy change: ΔG =Δ H - T Δ S, where Δ H
has a small positive value and T Δ S a large positive value; thus
ΔG is negative
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Non-polar gases and water
The molecules of the biologically important gases O2, CO2 and
N2 are non-polar
In O2 and N2, electrons are shared equally by both atoms. In
CO2, each C=O bond is polar, but the two dipoles are oppositely
directed and cancel each other out
The movement of molecules from the disordered gas phase
into aqueous solution constrains their motion and the motion
of water molecules and represents a decrease in entropy
The non-polar nature of these gases and the decrease in
entropy when they enter solution combine to make them very
poorly soluble in water
O2 is carried by the water soluble proteins hemoglobin and
myoglobin ; CO2 is either carried as it is by hemoglobin or is
changed to the soluble form –bicarbonate, HCO310
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Non-polar solutes in water
• Non-polar compounds such as benzene and hexane are
hydrophobic—they are unable to undergo energetically
favorable interactions with water molecules, and they interfere
with the hydrogen bonding between water molecules
• All molecules or ions in aqueous solution interfere with the
hydrogen bonding of some water molecules in their immediate
vicinity, but polar or charged solutes (such as NaCl)
compensate for lost water-water hydrogen bonds by forming
new solute-water interactions; the net change in enthalpy (ΔH)
for dissolving these solutes is generally small
• Hydrophobic solutes, however, offer no such compensation,
and their addition to water may therefore result in a small gain
of enthalpy; the breaking of hydrogen bonds between water
molecules takes up energy from the system
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• Furthermore, dissolving hydrophobic compounds in water
produces a measurable decrease in entropy. Water molecules
in the immediate vicinity of a non-polar solute are constrained
in their possible orientations as they form a highly ordered
cage-like shell around each solute molecule
• The ordering of water molecules reduces entropy. The number
of ordered water molecules, and therefore the magnitude of
the entropy decrease, is proportional to the surface area of
the hydrophobic solute enclosed within the cage of water
molecules
• The free energy change for dissolving a non-polar solute in
water is thus unfavorable: ΔG=ΔH - TΔS, where ΔH has a
positive value, TΔS has a negative value, and ΔG is positive
• This unfavorable state is relieved when the non-polar solutes
coalesce to form droplets
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• This coalition is called the hydrophobic effect /interaction
• Hydrophobic interaction refers to the tendency of non-polar
compounds to self-associate in an aqueous environment
• This self-association is driven neither by mutual attraction nor
by what are sometimes incorrectly referred to as “hydrophobic
bonds.” Self-association minimizes energetically unfavorable
interactions between non-polar groups and water
• A solvation sphere of hydrogen-bonded water molecules forms
around the hydrophobic molecules
• Although non-polar molecules, when in close proximity, are
attracted to each other by van der Waals forces, the driving
force for in the formation of the solvation spheres is the strong
tendency of water molecules to form hydrogen bonds among
themselves; non-polar molecules are excluded because they
cannot form hydrogen bonds
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Formation of an oil-droplet in an aqueous solution
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Amphipathic molecules in water
Amphipathic compounds contain regions that are polar (or
charged) and regions that are non-polar
When an amphipathic compound is mixed with water, the
polar, hydrophilic region interacts favorably with the
solvent and tends to dissolve, but the non-polar,
hydrophobic region tends to avoid contact with the water
The non-polar 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
These stable structures of amphipathic compounds in
water, called micelles, may contain hundreds or thousands
of molecules
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• Many biomolecules are
amphipathic: proteins tend to
fold with the R-groups of
amino acids with hydrophobic
side chains in the interior;
amino acids with charged or
polar amino acid side chains
generally are present on the
surface in contact with water
• A similar pattern prevails in a
phospholipid bilayer, where
the charged head groups
contact water while their
hydrophobic fatty acyl side
chains cluster together,
excluding water
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Liposomes are
formed through
the sonication of a
solution of
amphipathic
molecules. They
have a potential for
drug delivery
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Water as a Participant in Chemical Reactions
Metabolic reactions often involve the attack by lone pairs of
electrons residing on electron-rich molecules termed
nucleophiles upon electron-poor atoms called electrophiles
Nucleophiles and electrophiles do not necessarily possess a
formal negative or positive charge; water, whose two lone pairs
of electrons bear a partial negative charge, is an excellent
nucleophile
Other nucleophiles of biologic importance include the oxygen
atoms of phosphates, alcohols and carboxylic acids; the sulfur
of thiols; the nitrogen of amines; and the imidazole ring of
histidine
Common electrophiles include the carbonyl carbons in amides,
esters, aldehydes, and ketones and the phosphorus atoms of
phosphoesters
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• Nucleophilic attack by water generally results in the cleavage of
the amide, glycoside, or ester bonds that hold biopolymers
together; this process is termed hydrolysis
• Conversely, when monomer units are joined together to form
biopolymers such as proteins or glycogen, water is a product
The Thermal Properties of Water
• If water followed the pattern of compounds such as hydrogen
sulfide, it would melt at -100 0C and boil at -910C
• Under these conditions, most of the earth’s water would be
steam, making life unlikely
• However, water actually melts at 0 0C and boils at +100 0C;
consequently, it is a liquid over most of the wide range of
temperatures found on the earth’s surface
• Hydrogen bonding is responsible for this behavior of water
• Energy is required to break hydrogen bonds
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• When ice is warmed to its melting point, approximately 15% of
the hydrogen bonds break
• Liquid water consists of ice-like clusters of molecules whose
hydrogen bonds are continuously breaking and forming
• As the temperature rises, the movement and vibrations of the
water molecules accelerate and additional hydrogen bonds are
broken
• When the boiling point is reached, the water molecules break
free from one another and vaporize
• The energy required to raise water’s temperature is
substantially higher than expected
• One consequence of water’s high heat of vaporization (the
energy required to vaporize 1 mole of a substance at 1 atm) and
high heat capacity (the energy that must be added or removed
to change the temperature by one degree Celsius) is that water
acts as an effective modulator of climatic (and body) temp. 22
• Water can absorb solar heat and release it slowly
• Water’s high heat capacity, coupled with the high water
content found in most organisms helps maintain an
organism’s internal temperature
• The evaporation of water is used as a cooling mechanism,
because it permits large losses of heat
• For example, an adult human may eliminate as much as
1200g of water daily in expired air, sweat and urine
• The associated heat loss may amount to approximately
20% of the total heat generated by metabolic processes
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Colligative Properties
Solutes of all kinds alter certain physical properties of the
solvent, water: its vapor pressure, boiling point, melting point
(freezing point), and osmotic pressure
These are called colligative (“tied together”) properties,
because the effect of solutes on all four properties has the
same basis: the concentration of water is lower in solutions
than in pure water
The effect of solute concentration on the colligative properties
of water is independent of the chemical properties of the
solute; it depends only on the number of solute particles
(molecules, ions) in a given amount of water
A compound such as NaCl, which dissociates in solution, has
twice the effect on osmotic pressure, for example, as does an
equal number of moles of a non-dissociating solute such as
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glucose
• Water molecules tend to move from a region of higher water
concentration to one of lower water concentration –osmosis
• When two different aqueous solutions are separated by a
semipermeable membrane (one that allows the passage of
water but not solute molecules), water molecules diffusing
from the region of higher water concentration to that of lower
water concentration produce osmotic pressure
• A solution containing 1 mol of solute particles in 1 kg of water is
a 1-osmolal solution
• When 1 mol of a solute (such as NaCl) that dissociates into two
ions (Na + and Cl-) is dissolved in 1 kg of water, the solution is 2osmolal
• Measurement of colligative properties is useful in estimating
solute concentrations in biological fluids. For example, in blood
plasma, the normal total concentration of solutes is remarkably
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constant (275-295 milliosmolal).
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• If a cell is put in In a hypotonic solution, with lower osmolality
than the cytosol, the cell swells as water enters
• In their natural environments, cells generally contain higher
concentrations of biomolecules and ions than their
surroundings, so osmotic pressure tends to drive water into
cells
• If not somehow counterbalanced, this inward movement of
water would distend the plasma membrane and eventually
cause bursting of the cell (osmotic lysis)
• In multicellular animals, blood plasma and interstitial fluid are
maintained at an osmolality close to that of the cytosol; the
high concentration of albumin and other proteins in blood
plasma contributes to its osmolality
• Cells also actively pump out ions such as Na+ into the
interstitial fluid to stay in osmotic balance with their
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surroundings
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• Because the effect of solutes on osmolality depends on the
number of dissolved particles, not their mass, macromolecules
(proteins, nucleic acids, polysaccharides) have far less effect on
the osmolality of a solution than would an equal mass of their
monomeric components
• One effect of storing fuel as polysaccharides (starch or
glycogen) rather than as glucose or other simple sugars is
prevention of an enormous increase in osmotic pressure within
the storage cell
The Gibbs-Donnan Equilibrium
• The three fluid compartments, that is, the intracellular fluid,
interstitial fluid and blood plasma each contain diffusible ions
such as Na+, K+, Cl- and HCO3• In addition, the intracellular fluid and the plasma contain nondiffusible proteins
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• The negatively charged, non-diffusible proteins present
predominantly in the plasma space will attract positively charged
ions and repel negatively charged ions
• Despite the high permeability of small ions across membranes, a
similar concentration of ionic species is not seen
 The passive distribution of cations and anions is altered to
preserve electroneutrality in the compartments
• The normal difference in concentrations of diffusible ions
between the plasma and interstitial compartments is due to the
presence of non-diffusible proteins in plasma
• The diffusible cation concentration is higher in the compartment
containing non-diffusible, anionic proteins, whereas diffusible
anion concentration is lower in the protein-containing
compartment
• Gibbs-Donnan equilibrium is established when the altered
distribution of cations and anions results in electrochemical 30
equilibrium
Semi-permeable membrane
Distribution of inorganic ions in the absence of
non-diffusible ions
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More Cl- leaves I to balance charges
Distribution of inorganic ions in the presence of
non-diffusible ions
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• The existence of ionic asymmetry on the surfaces on the
surface of cell membrane results in the establishment of the
electrochemical gradient or membrane potential which
provides the means for electrical conduction and active and
passive transport
• A related outcome is that water tends to move from the
interstitial space to the plasma (maintaining blood volume) and
the intercellular space (causing a constant threat of cellular
swelling)
• Cells must, therefore constantly regulate their osmolality;
many animal and bacterial cells pump out inorganic ions such
as Na+ thereby regulating cell volume
• About 1/3 of ATP in an animal cell is used to power Na+-K+
pumps; in nerve cells, which use Na+ and K+ gradients to
propagate electrical signals, up to 2/3 of the ATP is used to
power these pumps
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Dissociation of Water and the pH Scale
• Acids are compounds that donate a hydrogen ion (H+) to a
solution, and bases are compounds (such as the OH- ion) that
accept hydrogen ions
• Water itself dissociates to a slight extent, generating hydrogen
ions , which are also called protons, and hydroxide ions
H2O <---> H+ + OH• The hydrogen ions are extensively hydrated in water to form
species such as H3O+ (hydronium), but nevertheless are usually
represented simply as H+.Water itself is neutral, neither acidic
nor basic
• For the dissociation of water:
where the brackets represent molar
concentrations and K is the
dissociation constant
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• Since 1 mole (mol) of water weighs 18 g, 1 liter (L) (1000 g) of
water contains 1000/18 = 55.56 mol. Pure water thus is 55.56
molar
• K can be determined by measurement of the electrical
conductivity of pure water, which has the value of 1.8 x 10 -16
M at 25 ℃ indicative of a very small ion concentration, where
M (molar) is the unit of moles per liter
• Therefore, the concentration of undissociated water is
essentially unchanged by the dissociation reaction
• Substituting for the values of K and [H2O]:
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[H+] [OH ] = 1.8 x10 -16 M x 55.56 M = 1 x 10-14 M2 =KW
• KW is known as the ion product of water
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+
• Since the concentrations of [H ] [OH ] in pure water are
equal:
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[H+]= [OH ] = 10-7 M
• pH is employed to express proton concentrations in a
convenient form; it is the negative log (to the base ten) of
the hydrogen ion concentration:
pH=-log[H+]
• For pure water, pH=-log [10-7 ]=7; and pOH =-log [10-7 ]=7
• A pH of 7 is termed neutral because [H+] and [OH-] are equal.
Acidic solutions have a greater hydrogen ion concentration
and a lower hydroxide ion concentration (pH<7) than pure
water and basic solutions have a lower hydrogen ion
concentration and a greater hydroxide ion concentration
(pH >7)
• A decrease in one pH unit reflects a 10-fold increase in H+
concentration
• Strong acids/bases completely dissociate in water; weak
acid/bases dissociate only partially
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• Many biochemicals possess functional groups (carboxyl groups,
amino groups, phosphate esters,…) that are weak acids or
bases
• The relative strengths of weak acids and bases are expressed in
terms of their dissociation constants
• For the reaction HA<---> A- +H+
Where Ka is the dissociation constant, HA is
the conjugate acid and A- is the conjugate base
• Since the numeric values of Ka for weak acids are negative
exponential numbers, pKa is used where
pKa = -log Ka
• The stronger the acid the lower its pKa value
• For any weak acid, its conjugate is a strong base. Similarly, the
conjugate of a strong base is a weak acid. The relative strengths
of bases are expressed in terms of the pKa of their conjugate37
acids
Titration curves reveal the pKa
• Titration is used to determine the amount of an acid in a
given solution
• A measured volume of the acid is titrated with a solution of a
strong base, usually NaOH, of known concentration
• The NaOH is added in small increments until the acid is
consumed (neutralized), as determined with an indicator dye
or a pH meter
• The concentration of the acid in the original solution can be
calculated from the volume and concentration of NaOH
added
• A plot of pH against the amount of NaOH added (a titration
curve) reveals the pKa of the weak acid
• Consider the titration of a 0.1 M solution of acetic acid (for
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simplicity denoted as HAc) with 0.1 M NaOH at 25 0C
• Two reversible equilibria are involved in the process:
H2O <--->H++ OHHAc <---> H++ Ac• The equilibria must simultaneously conform to their
characteristic equilibrium constants, which are, respectively,
• At the beginning of the titration, before any NaOH is added,
the acetic acid is already slightly ionized, to an extent that
can be calculated from its dissociation constant
• As NaOH is gradually introduced, the added OH- combines
with the free H+ in the solution to form H2O, to an extent that
satisfies the equilibrium relationship of water
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• As free H+ is removed, HAc dissociates further to satisfy its own
equilibrium constant
• The net result as the titration proceeds is that more and more
HAc ionizes, forming Ac-, as the NaOH is added
• At the midpoint of the titration, at which exactly 0.5 equivalent
of NaOH has been added, one-half of the original acetic acid has
undergone dissociation that the concentration of the proton
donor, [HAc], now equals that of the proton acceptor, [Ac-]
• At this midpoint, a very important relationship holds: the pH of
the equimolar solution of acetic acid and acetate is exactly
equal to the pKa of acetic acid (4.76)
• As the titration is continued by adding further increments of
NaOH, the remaining non-dissociated acetic acid is gradually
converted into acetate. The end point of the titration occurs at
about pH 7.0: all the acetic acid has lost its protons to OH-, to
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form water and acetate
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What are Buffers?
Buffers are solutions that resist change in pH when small
amounts of proton (acid) or hydroxide (base) are added
They are either a mixture of a weak acid (HA) and its conjugate
base (A-) or a mixture of a weak base (B) and its conjugate acid
(HB+)
The mixture of equal concentrations of acetic acid and acetate
ion, found at the midpoint of the titration curve is a buffer
system
The titration curve of acetic acid has a relatively flat zone
extending about 1 pH unit on either side of its midpoint pH 4.76
In this zone, an amount of H+ or OH- added to the system has
much less effect on pH than the same amount added outside
the buffer range
This relatively flat zone is the buffering region of the acetic
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acid–acetate buffer pair
• At the
midpoint of
the buffering
region, where
the conc. of
the proton
donor exactly
equals that of
the proton
acceptor, the
buffering
power of the
system is
maximal
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The Henderson-Hasselbalch Equation
• The shape of the titration curve of weak acids and bases is
described by the Henderson-Hasselbalch equation
• This equation relates pH, pKa and the concentration of
conjugate acid-base pairs; it is derived as follows:
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• At the midpoint of titration, the concentrations of proton
acceptor and donor are equal; log (1)= 0; pH= pKa
• If the ratio [A-]/[HA] is 100:1, pH= pKa + 2
• If the ratio [A-]/[HA] is 1:10, pH= pKa - 1; …
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Normal pH values in organisms
pH values in the cell and in the extracellular fluids are kept
constant within narrow limits
In the blood, the pH value normally ranges only between 7.35
and 7.45; this corresponds to a maximum change in the H+
concentration of ca. 30%
The pH value of cytoplasm is slightly lower than that of
blood, at 7.0–7.3
In the lumen of the gastrointestinal tract and in the body’s
excretion products, the pH values are more variable
Extreme values are found in the stomach (ca.2) and in the
small intestine (> 8)
Since the kidney can excrete either acids or bases, depending
on the state of the metabolism, the pH of urine has a
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particularly wide range of variation (4.8–7.5)
• If the H + concentration departs significantly from its normal
value, the health and survival of the human body are in
jeopardy
• H + is the smallest ion, and it combines with many
negatively charged and neutral functional groups
• Changes of [H +], therefore, affect the charged regions of
many molecular structures, such as enzymes, cell
membranes and nucleic acids, and dramatically alter
physiological activity
• If the plasma pH reaches either 6.8 or 7.8, death may be
unavoidable
• Despite the fact that large amounts of acidic and basic
metabolites are produced and eliminated from the body,
buffer systems maintain a fairly constant pH in body fluids
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A More Meaningful Way of Stating the Concentration of
Hydrogen Ions
In clinical acid-base problems, the use of the pH scale has some
disadvantages
Since the pH is the logarithm of the reciprocal of [H +],
significant variations of [H +]in a patient may not be fully
appreciated
For example, if the blood pH decreases from 7.4 to 7.1, [H +] is
doubled; or if the pH increases from 7.4 to 7.7, [H +] is halved
Thus, in clinical situations it is preferable to express [H +]
directly as nanomoles per liter in order to better evaluate acidbase changes and interpret laboratory tests
A blood pH of 7.40 corresponds to 40 nM [H +], which is the
mean of the normal range ; the normal range is 7.36-7.44 on
the pH scale, or 44-36 nM [H +]
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Metabolic Acids and Bases
During metabolism, the body produces a number of acids that
increase the hydrogen ion concentration of the blood or other
body fluids and tend to lower the pH
These metabolically important acids can be weak acids or
strong acids
Inorganic acids such as sulfuric acid (H2SO4) and hydrochloric
acid (HCl) are strong acids
Organic acids containing carboxylic acid groups (e.g., the
ketone bodies acetoacetic acid and β-hydroxybutyric acid) are
weak acids
An average rate of metabolic activity produces roughly 22,000
mEq acid per day
If all of this acid were dissolved at one time in unbuffered body
fluids, their pH would be less than 1
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• Until the acid produced from metabolism can be excreted as
CO2 in expired air and as ions (and unmetabolized organic acids)
in the urine, it needs to be buffered in the body fluids
• The major buffer systems in the body are: the bicarbonate–
carbonic acid buffer system, which operates principally in
extracellular fluid; the hemoglobin buffer system in red blood
cells; the phosphate buffer system in all types of cells; the
protein buffer system in cells and plasma and phosphate and
ammonia in the urine
The Bicarbonate Buffer System
• The major source of metabolic acid in the body is the gas CO2,
produced principally from fuel oxidation in the TCA cycle
• Under normal metabolic conditions, the body generates
more than 13 moles of CO2 per day (approximately 0.5-1 kg)
• About 95% of the CO2 entering the blood diffuses into the red
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blood cells
• Within the red blood cells, the enzyme carbonic anhydrase I
catalyzes the conversion of most of the CO2 to carbonic acid
(H2CO3 )
• Carbonic anhydrase II is found in most tissues including the
lung, bone and renal tubular cells
• Carbonic acid is a weak acid that
partially dissociates into H+ and
bicarbonate anion, HCO3• Although H2CO3 is a weak acid, its
dissociation is essentially 100%
because of the removal of H+ ions by
the buffering action of hemoglobin
• The remainder of the H+ is buffered
by phosphate and proteins other than
hemoglobin
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• As the concentration of HCO 3- (i.e., of metabolic CO2) in red
blood cells increases, an imbalance occurs between the
bicarbonate ion concentrations in the red blood cell and plasma
• This osmotic imbalance causes a marked efflux of HCO 3- to
plasma and consequent influx of Cl- from plasma in order to
maintain the balance of charges
• This exchange, known as the chloride shift, takes place
through an antiporter known as band 3 protein
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• Once in the extracellular fluid, HCO3- serves as a major buffer
• Most of the CO2 produced in the body reaches the lungs
carried by the plasma in the form of HCO3• In the lungs, the events that took place in the erythrocytes
are reversed and CO2 is exhaled
• Buffering capacity is greatest at or near the pKa of the
conjugate-acid base pair
 The pKa of H2CO3 is 3.8 but it is a good buffer at the blood
pH of 7.4; how could this be?
 The most effective buffers are those that contain equal
concentrations of both components. But at pH 7.4, the
concentration of H2CO3 is a very small fraction of the
concentration of HCO3- and the plasma appears to be
poorly protected against an influx of OH- . How is this
problem solved?
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• Gaseous carbon dioxide from the lungs and tissues is dissolved
in the blood plasma, symbolized as CO2(d), and hydrated to
form H2CO3:
• In mammalian body fluids, the equilibrium for the carbonic
anhydrase reaction lies far to the left, such that about 500 CO2
molecules are present in solution for every molecule of H2CO3
• Because dissolved CO2 and H2CO3 are in equilibrium, the
proper expression for H2CO3 availability is [CO2(d)] + [H2CO3],
the so-called total carbonic acid pool, consisting primarily of
CO2(d)
• The overall equilibrium for the bicarbonate buffer system is:
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• An expression for the ionization of H2CO3 under such
conditions (that is, in the presence of dissolved CO2) can be
obtained from Kh, the equilibrium constant for the hydration of
CO2, and from Ka, the acid dissociation constant for H2CO3:
• Putting this value for [H2CO3] into the expression for the
dissociation of H2CO3 gives:
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• KaKh, the product of two constants, can be defined as a new
equilibrium constant, Koverall
• The value of Kh is 0.003 and Ka is equal to 0.000269. Therefore,
Koverall = 8.07 x 10-7 and pKoverall = 6.1
• This gives a modified Henderson-Hasselbalch equation for the
bicarbonate buffer system:
• The concentration gap that existed between H2CO3 and HCO3has been greatly narrowed by usingCO2(d) in the equation
• But still, 6.1 is more than one unit away from 7.4 and the ratio
of conjugate base (bicarbonate) over conjugate acid (mainly
carbondioxide) is 20:1
• It appears that the conjugate acid, because of its small
concentration would be overwhelmed by small amounts of59
alkali
• However, the acid component is the total carbonic acid pool,
that is, [CO2(d)] + [H2CO3], which is stabilized by its equilibrium
with CO2(g)
• The gaseous CO2 buffers any losses from the total carbonic
acid pool by entering solution as CO2(d), and blood pH is
effectively maintained
• Thus, the bicarbonate buffer system is an open system
• In the equilibrium expression for the bicarbonate-carbonic acid
buffer system at pH 7.4, the carbonic acid term can be replaced
by a pressure term because the carbonic acid concentration is
proportional to the partial pressure of carbondioxide ,PCO2, in
the blood
• For normal plasma at 370C solubility
coefficient, a = 0.03 mmol of
dissolved CO2 per liter of plasma per
60
mm Hg of CO2 pressure
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•
•
•
•
•
The Phosphate Buffer System
Phosphate is an abundant anion in cells, both in inorganic
form and as an important functional group on organic
molecules that serve as metabolites or macromolecular
precursors
The inorganic phosphate buffer consists of the weak acidconjugate base pair dihydrogen phosphate/hydrogen
phosphate
H2PO4- <---> HPO4-2 + H+
The pka of the system is 7.2 so it would appear that it is an
excellent choice for buffering blood
Although the blood pH of 7.4 is well within the buffer system’s
capability, the concentrations of H2PO4- and HPO4-2 in blood
are too low (4mEq/L) to have a major effect
Instead, the phosphate system is an important buffer in
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intracellular fluids where its concentration is about 75 mEq/L
• Organic phosphate anions, such as glucose 6-phosphate and
ATP, also act as intracellular buffers
• Although cells contain other weak acids these substances are
unimportant as buffers because of their low concentrations
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and pka that is much lower than intracellular pH
•
•
•
•
•
•
Protein Buffers (The Histidine System)
Histidine is one of the 20 naturally occurring amino acids
commonly found in proteins
It possesses as part of its structure an imidazole group, a fivemembered heterocyclic ring possessing two nitrogen atoms.
The pKa for dissociation of the imidazole hydrogen of histidine
is 6.04
In cells, histidine occurs as the free amino acid, as a constituent
of proteins, and as part of dipeptides in combination with other
amino acids
Because the concentration of free histidine is low and its
imidazole pKa is more than 1 pH unit removed from prevailing
intracellular pH, its role in intracellular buffering is minor
However, protein-bound and dipeptide histidine may be the
dominant buffering system in some cells
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• In combination with other amino acids, as in proteins or
dipeptides, the imidazole pka may increase substantially
approach the physiological pH
 The pka of weak acids can be affected by their environments
• The main protein in erythrocytes, hemoglobin, uses its
histidines to buffer the protons released from carbonic acid
and other sources
• Other cells are endowed with other proteins that assist in
intracellular buffering
• Albumin in the blood also serves as a buffer
Histidine
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•
•
•
•
•
Urinary Buffers
The non-volatile acid that is produced from body
metabolism cannot be excreted as expired CO2 and is
excreted in the urine
Most of the non-volatile acid hydrogen ion is excreted as
undissociated acid that generally buffers the urinary pH
between 5.5 and 7.0; a pH of 5.0 is the minimum urinary pH
The acid secretion includes inorganic acids such as
phosphate and ammonium ions, as well as uric acid,
dicarboxylic acids, and tricarboxylic acids such as citric acid
Sulfuric acid is generated from the sulfate-containing
compounds ingested in foods and from metabolism of the
sulfur-containing amino acids, cysteine and methionine
It is a strong acid that is dissociated into H+ and sulfate anion
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(SO4-2) in the blood and urine
• Urinary excretion of phosphate ions helps to remove acid;
to maintain metabolic homeostasis, we must excrete the
same amount of phosphate in the urine that we ingest with
food as phosphate anions or organic phosphates such as
phospholipids
• Whether the phosphate is present in the urine as H2PO4- or
HPO4 -2 depends on the urinary pH and the pH of blood
• Ammonium ions are major contributors to buffering urinary
pH, but not blood pH
• Ammonia (NH3) is a base that combines with protons to
produce ammonium (NH4+) ions a reaction that occurs with
a pKa of 9.25
• Ammonia is produced from the catabolism of nitrogen
containing biomolecules and kept at very low concentrations
in the blood because it is toxic to neural tissues
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• Cells in the kidney generate NH4+ and excrete it into the
urine in proportion to the acidity of the blood
• As the renal tubular cells transport H+ into the urine, they
_
return HCO3 anions to the blood
 Hydrochloric acid (HCl), also called gastric acid, is
secreted by parietal cells of the stomach into the stomach
lumen, where the strong acidity denatures ingested
proteins so they can be degraded by digestive enzymes
• When the stomach contents are released into the lumen
of the small intestine, gastric acid is neutralized by
bicarbonate secreted from pancreatic cells and by cells in
the intestinal lining
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