Water, pH, buffer system

1. Water, pH and Buffers
MBBS Biochemistry (1st year)

2. WATER IS AN IDEAL BIOLOGIC
SOLVENT
• Water is the predominant chemical component of living organisms. Its unique physical properties,
which include the ability to solvate a wide range of organic and inorganic molecules, derive from
water’s dipolar structure and exceptional capacity for forming hydrogen bonds
• Water Molecules Form Dipoles
A molecule with electrical charge distributed asymmetrically about its structure is referred to as a
dipole
Water’s strong dipole is responsible for its high dielectric constant
Its strong dipole and high dielectric constant enable water to dissolve large quantities of charged
compounds such as salts
• Water Molecules Form Hydrogen Bonds
Hydrogen bonding profoundly influences the physical properties of water and accounts for its
relatively high viscosity, surface tension, and boiling point
Hydrogen bonding enables water to dissolve many organic biomolecules that contain functional
groups which can participate in hydrogen bonding

3. The oxygen atoms of aldehydes, ketones, and amides, for
example, provide lone pairs of electrons that can serve as
hydrogen acceptors. Alcohols, carboxylic acids, and amines can
serve both as hydrogen acceptors and as donors of unshielded
hydrogen atoms for formation of hydrogen bonds

4. INTERACTION WITH WATER
INFLUENCES THE STRUCTURE
OF BIOMOLECULES
• Covalent and Noncovalent Bonds Stabilize Biologic Molecules
The covalent bond is the strongest force that holds molecules together. Noncovalent forces, while of
lesser magnitude, make significant contributions to the structure, stability, and functional
competence of macromolecules in living cells
• Biomolecules Fold to Position Polar & Charged Groups on Their Surfaces
Most biomolecules are amphipathic; that is, they possess regions rich in charged or polar
functional groups as well as regions with hydrophobic character
Proteins tend to fold with the R-groups of amino acids with hydrophobic sidechains in the interior.
Amino acids with charged or polar amino acid side chains
• Hydrophobic Interactions
Hydrophobic interaction refers to the tendency of nonpolar compounds to self-associate in an
aqueous environment
Self-association minimizes the disruption of energetically favorable interactions between the
surrounding water molecules

5. • Electrostatic Interactions
Interactions between charged groups help shape biomolecular structure
Electrostatic interactions between oppositely charged groups within or between biomolecules are
termed salt bridges
Salt bridges are comparable in strength to hydrogen bonds but act over larger distances. They
therefore often facilitate the binding of charged molecules and ions to proteins and nucleic acids
• van der Waals Forces
van der Waals forces arise from attractions between transient dipoles generated by the rapid
movement of electrons of all neutral atoms
Significantly weaker than hydrogen bonds but potentially extremely numerous, van der Waals forces
decrease as the sixth power of the distance separating atoms

6. Multiple Forces Stabilize Biomolecules
• The DNA double helix illustrates the contribution of multiple forces to the structure of biomolecules.
• While each individual DNA strand is held together by covalent bonds,
• the two strands of the helix are held together exclusively by noncovalent interactions such as
hydrogen bonds between nucleotide bases (Watson-Crick base pairing) and van der Waals
interactions between the stacked purine and pyrimidine bases.
• The double helix presents the charged phosphate groups and polar hydroxyl groups from the ribose
sugars of the DNA backbone to water while burying the relatively hydrophobic nucleotide bases
inside.
• The extended backbone maximizes the distance between negatively charged phosphates,
minimizing unfavorable electrostatic interactions

7. WATER IS AN EXCELLENT
NUCLEOPHILE
• Metabolic reactions often involve the attack by lone pairs of electrons residing on electron-rich
molecules termed nucleophiles upon electron-poor atoms called electrophiles
• Other nucleophiles of biologic importance include the oxygen atoms of phosphates, alcohols, and
carboxylic acids; the sulfur of thiols; and the nitrogen atom of amines and of the imidazole ring of
histidine
• Common electrophiles include the carbonyl carbons in amides, esters, aldehydes, and ketones and
the phosphorus atoms of phosphoesters
• Nucleophilic attack by water typically results in the cleavage of the amide, glycoside, or ester bonds
that hold biopolymers together. This process is termed hydrolysis
• In the cell, protein catalysts called enzymes accelerate the rate of hydrolytic reactions when
needed. Proteases catalyze the hydrolysis of proteins into their component amino acids, while
nucleases catalyze the hydrolysis of the phosphoester bonds in DNA and RNA

8. ACIDS AND BASES
• 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 (H), which are also called protons, and
hydroxide ions (OH)
• The hydrogen ions are extensively hydrated in water to form species such as H3O, but nevertheless are
usually represented as simply H. Water itself is neutral, neither acidic nor basic
• Strong and Weak Acids
During metabolism, the body produces several 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 classified as weak acids or strong acids by their degree of
dissociation into a hydrogen ion and a base (the anion component)
Inorganic acids such as sulfuric acid (H2SO4) and hydrochloric acid (HCl) are strong acids that dissociate
completely in solution

9. • Organic acids containing carboxylic acid groups (e.g., the ketone bodies acetoacetic acid and -hydroxybutyric
acid) are weak acids that dissociate to only a limited extent in water.
• In general, a weak acid (HA), called the conjugate acid, dissociates into a hydrogen ion and an anionic
component (A), called the conjugate base
• The tendency of the acid (HA) to dissociate and donate a hydrogen ion to solution is denoted by its Ka, the
equilibrium constant for dissociation of a weak acid
• The higher the Ka, the greater is the tendency to dissociate a proton
• We express the relative strengths of weak acids and bases in terms of their dissociation constants.
• dissociation constant to calculate the concentration of [H+] (or [OH-]) produced by a given
molarity of a weak acid (or base) before calculating total [H+] (or total [OH-]) and subsequently pH

10. Acids in the Blood of a Healthy Individual

11. Water Molecules Exhibit a Slight but
Important Tendency to Dissociate
• K is the dissociation constant. 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. Since the probability that a hydrogen in
pure water will exist as a hydrogen ion is 1.8 × 10−9,
• The molar concentration of H+ ions (or of OH− ions) in pure water is the
product of the probability, 1.8 × 10−9, times the molar concentration of
water, 55.56 mol/L.
• The result is 1.0 × 10−7 mol/L.

12. pH IS THE NEGATIVE LOG OF THE
HYDROGEN ION
CONCENTRATION
To calculate the pH of a solution:
1. Calculate the hydrogen ion concentration [H+]
2. Calculate the base 10 logarithm of [H+].
3. pH is the negative of the value found in step 2.

13. Functional Groups That Are Weak
Acids
Have Great Physiologic Significance
• Many biochemicals possess functional groups that are weak acids or bases
• Carboxyl groups, amino groups, and phosphate esters, whose second dissociation falls within the
physiologic range, are present in proteins and nucleic acids, most coenzymes, and most
intermediary metabolites
• Show blow the expressions for the dissociation constant (Ka) for two representative weak acids,
R—COOH and R—NH3+.
 Since the numeric values of Ka for weak acids are
negative
exponential numbers, we express Ka as pKa
 Note that pKa is related to Ka as pH is to [H+]. The
stronger the
acid, the lower is its pKa value.

14. • pKa is used to express the relative strengths of both acids and bases
• The relative strengths of bases are expressed in terms of the pKa of their conjugate acids.
The pKa of an acid group is the pH at which the protonated
and unprotonated species are present at equal
concentrations.

15. The Henderson-Hasselbalch Equation
Describes the Behavior of Weak Acids &
Buffers

16. pH Measurement
• Indicator dyes
Chemically an indicator (designated as HI) is a weak acid
The indicators most commonly used are phenol red or phenolphthalein
• Indicator papers (pH papers)
pH range, which occurs around its pK’ value, brings about a visible change in colour of the indicator
• pH meter:
It is used for accurate measurement of pH

17. Titration Curve of Weak Acids
• Titration is a procedure that is used to quantitatively determine the amount of acid in a given
solution. In this procedure, an acid reacts with an alkali of known concentration till the point of
neutralization (end point).
• Titration curve is drawn by plotting the pH against the amount of alkali added. It indicates the pH
changes of the acid that occur when alkali is added in small amounts.
• Titration curve of acetic acids vs sodium hydroxide
• 1. At the start of the titration, acetic acid is present predominantly in undissociated state (HA). As
the alkali (sodium hydroxide; 0.1N) is added, it reacts with the acid molecules to form acetate ions
(A). Thus, there is a progressive rise in the concentration of base (i.e. acetate) with a concomitant
fall in the concentration of the acid.
• 2. At the midpoint of titration about half of the acetic acid loses its proton to form acetate. Thus,
concentrations acetic acid and its conjugate base (acetate) are equal at this point
• 3. Towards the end of the titration, most of acetic acid loses its proton, so that the predominant ionic
species present is acetate.

19. Buffers
• Buffer is a system that resists any alteration in its pH when a small amount of acid or alkali is
added to it. It comprises
• two major components: a weak acid (HA) and its conjugate base (A).
• A buffer system is most effective when;
(a) these two components are present in equimolar concentrations
(b) pH (of the medium) equals pK’ (of the acid-base pair)
• Mechanism of Action of a Buffer System
• When a small amount of acid is added, it is taken up by the base component of
the buffer (A), and any pH change is averted.
• Similarly, the acid component of the buffer system (HA) is capable of reacting with
any OH that is added.

20. Solutions of Weak Acids & Their
Salts Buffer Changes in pH
• Many metabolic reactions are accompanied by the release or uptake of protons
• Oxidative metabolism produces CO2, the anhydride of carbonic acid, which if not buffered would
produce severe acidosis.
• Biologic maintenance of a constant pH involves buffering by phosphate, bicarbonate, and proteins,
which accept or release protons to resist a change in pH
• A solution of a weak acid and its conjugate base buffers most effectively in the pH range pKa
± 1.0 pH unit.
• Acid Strength Depends on Molecular Structure
Many acids of biologic interest possess more than one dissociating group
• pKa Values Depend on the Properties of the Medium
• The medium may either raise or lower the pKa relative to its value in water, depending on whether
the undissociated acid or its conjugate base is the charged species.

22. METABOLIC ACIDS AND BUFFERS
• An average rate of metabolic activity produces roughly 22,000 milliequivalents (mEq) of acid per day
• If all of this acid were dissolved at one time in unbuffered body fluids, their pH would be 1
• However, the pH of the blood is normally maintained between 7.36 and 7.44, and intracellular pH at
approximately 7.1 (between 6.9 and 7.4)
• The widest range of extracellular pH over which the metabolic functions of the liver, the beating of the heart,
and conduction of neural impulses can be maintained is 6.8 to 7.8
• Acid produced from metabolism can be excreted as CO2 in expired air and as ions in the urine, it needs to be
buffered in the body fluids
• The major buffer systems
Bicarbonate–carbonic acid buffer system, which operates principally in ECF
Hemoglobin buffer system in red blood cells
Phosphate buffer system in all types of cells
Protein buffer system of cells and plasma

23. The Bicarbonate Buffer System
• The major source of metabolic acid in the body is the gas CO2, produced principally from fuel oxidation in
the tricarboxylic acid (TCA) cycle. Under normal metabolic conditions, the body generates more than 13 mol
of CO2 per day (approximately0.5 to 1 kg).
• CO2 dissolves in water and reacts with water to produce carbonic acid, H2CO3, a reaction that is accelerated
by the enzyme carbonic anhydrase.
• Carbonic acid is a weak acid that partially dissociates into H and bicarbonate anion, HCO3
• 1. The base constituent, bicarbonate (HCO-3) is regulated by kidneys.
• 2. The acid component (H2CO3) is regulated by pulmonary ventilation
• Bicarbonate buffer is subject to regulation by kidneys and lungs
• The acid component of this buffer (also called respiratory component) is generated from dissolved
carbon dioxide [CO2(d)] and water
• Carbonic acid can dissociate to yield bicarbonate.

24. Mechanism of Bicarbonate buffer
system
 when an acid is added to blood, concentration of H
rises. The latter is taken up by 3 HCO- resulting in the
rise of concentration of carbonic acid (Step 1)
 This causes the Step 2 to go forward, and the
concentration of carbon dioxide (d) in the blood rises
 This results in an increase in the pressure of carbon
dioxide in the gas phase in the lungs (Step 3)
 The extra carbon dioxide is exhaled through increased
rate of breathing
 Reverse series of reactions occur when an alkali (OH)
is added. It is taken up by carbonic acid to form 3 HCO-
.

25. Effectiveness of Bicarbonate buffer
• Bicarbonate buffer system is an effective physiological buffer because of its equilibration with a
large reserve of gaseous carbon dioxide in the air space of the lungs
• pK’ of carbonic acid is 6.1, the bicarbonate buffer should be most effective at or around pH of 6.1
(i.e. 6.1 1) as a buffer is most effective when pH equals pK
• However, bicarbonate buffer is highly effective at the physiological pH of 7.4 also because of its
equilibration with gaseous carbon dioxide
• Carbonic anhydrase is the principle enzyme that catalyzes generation ofHCO3-
• Decreased activity of this enzyme, therefore, results in decreased plasma bicarbonate
concentration. Consequently, the ratio of bicarbonate to carbonic acid (normally 20) tends to fall,
resulting in a fall of pH

26. Bicarbonate and Hemoglobin in Red
Blood Cells
• The bicarbonate buffer system and hemoglobin in red blood cells cooperate in buffering the blood and
transporting CO2 to the lungs.
• Most of the CO2 produced from tissue metabolism in the TCA cycle diffuses into the interstitial fluid and the
blood plasma and then into red blood cell
• red blood cells contain high amounts of this enzyme, and CO2 is rapidly converted to carbonic acid (H2CO3)
within these cells
• carbonic acid dissociates, the H released is also buffered by combination with hemoglobin (Hb)
• The bicarbonate anion is transported out of the red blood cells into the blood in exchange for chloride anions,
and thus bicarbonate is relatively high in the plasma
• As the red blood cells approach the lungs, the direction of the equilibrium reverses. CO2 is released from the
red blood cells, causing more carbonic acid to dissociate into CO2 and water and more hydrogen ions to
combine with bicarbonate
• Hemoglobin loses some of its hydrogen ions, a feature that allows it to bind oxygen more readily. Thus, the
bicarbonate buffer system is intimately linked to the delivery of oxygen to tissues

28. Erythrocyte mechanism for bicarbonate
generation
 The respiratory center within the hypothalamus, which controls the rate
of breathing, is sensitive to changes in pH. As the pH falls, individuals
breathe more rapidly and expire more CO2.
 As the pH rises, they breathe more shallowly. Thus, the rate of
breathing contributes to regulation of pH through its effects on the
dissolved CO2 content of the blood.

29. Role of Kidneys in Bicarbonate
Homeostasis
• Kidneys play an important role in bicarbonate homeostasis through the following mechanisms:
• 1. Reabsorption of the filtered bicarbonate, i.e. bicarbonate reclamation
• 2. Generation of bicarbonate ions, termed new bicarbonate generation
• Both these actions depend on the carbonate dehydratase system.
 The carbonate dehydratase
mechanism may be stimulated
by a rise in pCO2 or a fall of
bicarbonate concentration
within the tubular cells.

30. Intracellular pH
• Phosphate anions and proteins are the major buffers involved in maintaining a constant pH of ICFs
• phosphate anions play a major role as an intracellular buffer in the red blood cell and in other types of cells,
where their concentration is much higher than in blood and interstitial fluid
• Organic phosphate anions such as glucose 6-phosphate and ATP also act as buffers
• The transport of hydrogen ions out of the cell is also important for maintenance of a constant intracellular pH
• Metabolism produces several other acids in addition to CO2. e.g. the metabolic acids acetoacetic acid and -
hydroxybutyric acid are produced from fatty acid oxidation to ketone bodies in the liver, and lactic acid is
produced by glycolysis in muscle and other tissues
• The pKa for most metabolic carboxylic acids is 5, so these acids are completely dissociated at the pH of blood
and cellular fluid
• If the cell becomes too acidic, more H is transported out in exchange for Na ions by a different transporter.
• If the cell becomes too alkaline, more bicarbonate is transported out in exchange for Cl ions.

31. Proteins as Buffer
• They act as buffers because many of them have amino acids which behave like weak acids
• It has a pK’ value of 6.0 which is close to the physiological pH. Therefore, it is very effective in living
systems
• The intracellular fl uid (ICF) proteins serve as major buffers.
• Since ICF volume is around 60% of the total body fluid volume, proteins may be considered as most
abundant buffers in the body

32. Urinary Buffers
• The nonvolatile acid that is produced from body metabolism cannot be excreted as expired CO2 and is
excreted in the urine
• Most of the nonvolatile acid hydrogen ion is excreted as undissociated acid that generally buffers the urinary
pH between 5.5 and 7.0.
• Hydrogen ion excretion requires presence of suitable buffer systems in urine. The H secreted into
the tubular lumen causes acidifi cation of urine
• Phosphate Buffer
• It is the major intracellular buffer. Its pK’ value of 6.86 is near the intracellular pH of 7.0.
Therefore, this buffer is very effective intracellularly. It consists of the following components:
• 1. as the proton donor (i.e. the acid component).
• 2. the proton acceptor (i.e. the base component)
• Ammonia, the other important urinary buffer, is produced by deamination of glutamine in renal
tubular cell

33. Buffering of H+in urine/ Phosphate &
Ammonia buffer

34. The Three-tier Defense
• Normal body metabolism poses a constant threat to pH because it generates various products that
can alter the blood pH
• A three-tier defense system, comprising buffers, lungs, and kidneys therefore, remains constantly in
operation to guard against any changes in pH
• (a) Buffers serve as the first-line of defense against acid load
• (b) Pulmonary ventilation takes from a few minutes to few hours to become operational. It has a
direct bearing on acid-base balance of the body because carbon dioxide is an acidic substance.
Since carbon dioxide is exhaled during expiration, increase in respiratory activity reduces the acidity
of body fluids
• (c) Renal adjustments take from several hours to few days to become effective, but provide a
long-term solution by supplementing buffer action.