this document gives details on water and biomolecues

1. MEDICAL BIOCHEMISTRY AND
MEDICINAL NUTRITION
(MBC 110):
-water
-Biomolecules
MS. SIANKANGA T.M

2. Water, Weak Acids and Bases
 Water is the most abundant substance in living systems, making up to 70% or more of the weight of most
organisms.
 The attractive forces between the water molecules and the slight tendency of water to ionize are of crucial importance to the
structure and function of biomolecules
Weak interactions in aqueous systems
 Hydrogen bonds between water molecules provide the cohesive forces that make water a liquid at room temperature and favor
formation of crystalline water at low temperature
 Polar biomolecules dissolve readily in water because they can replace water-water interactions with more energetically
favorable water-solute interactions
 Nonpolar biomolecules interfere with water-water interactions but are unable to form water-solute interactions- hence poorly
soluble in water; in aqueous solutions nonpolar molecules tend to cluster together
 Hydrogen bonds and ionic, hydrophobic , and van der Waal’s interactions are individually weak but collectively have a
significant influence on the three dimensional structures of proteins, nucleic acids, polysaccharides, and membrane lipids
2
2

3. 3

4. 1. Hydrogen bonding and polarity
 Water molecule is slightly skewed tetrahedron with oxygen at its centre. Two hydrogens and the two unshared electrons of the
remaining two sp3-hybridized orbitals occupy the corners of the tetrahedron
 The H-O-H bond angle is 104.5o, slightly less than the 109.5o of a perfect tetrahedron because of crowding by the nonbonding
orbitals of the oxygen atom
 The strongly electronegative oxygen atom pulls electrons away from the hydrogen nuclei, leaving them with partial positivity
while its two unshared electron constitute a region of local negativity
4
H
H
O
δ–
δ+
δ+
δ–
Electropositive region
Electronegative region
4

5.  The result of this unequal electron sharing is two electric dipoles in the water molecule, one along each of the H-O bonds; the
oxygen atom bears a partial negative charge (2δ−) and each hydrogen a partial positive charge (δ+)
 As a result there is an electrostatic attraction between the oxygen atom of one water molecule and the hydrogen of another,
called a hydrogen bond
5
5
H
H
O
104.5o
Hydrogen bond 0.177nm
Covalent bond 0.0965 nm
2δ−
2δ−
δ+
δ+
δ+
δ+

6. 6
 Hydrogen bonds are weaker than covalent bonds
 Hydrogen bonds in liquid water have a bond dissociation energy of about 20 kJ/mol, compared with 348 kJ/mol for the
covalent C-C bonds
 Although most of the molecules in liquid water are engaged in hydrogen bonding, the life time of each hydrogen bond is less
than 1 X 10−9 s. The sum of all hydrogen bonds in liquid water however confer great internal cohesion
 The nearly tetrahedral arrangement of the orbitals about the oxygen atom allows each water molecule to form hydrogen bonds
with as many as four neighboring water molecules
 In liquid water at room temperature each water molecule forms hydrogen bonds with an average of 3.5 molecules due to
continuous motion of water molecules
 In ice each water molecule is fixed in space and forms hydrogen bonds with 4 other water molecules to yield a regular
lattice structure
 Hydrogen bonds give water a high melting point.
 Hydrogen bonds contribute to water’s high specific heat (amount of heat needed to raise the temperature of 1 gm of a
substance 1oC) - due to the fact that hydrogen bonds must be broken to increase the kinetic energy (motion of molecules)
and temperature of a substance --> temperature fluctuation is minimal.
 Water has a high heat of vaporization - large amount of heat is needed to evaporate water because hydrogen bonds must
be broken to change water from liquid to gaseous state.
6

7. 2. Universal solvent. 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 with a lone pair of electrons) and a hydrogen covalently bonded to another electronegative atom ( the
hydrogen donor)
 Hydrogen atoms covalently bonded to carbon atoms (which are not electronegative) do not participate in hydrogen bonding
 This explains why butanol (CH3(CH2)2CH2OH) has a relatively high boiling point of 117 oC whereas butane
(CH3(CH2)2CH3) has a boiling point of -0.5 oC. Butanol has polar hydroxyl group and thus can form intermolecular
hydrogen bonds
 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
7

8. 8
 Uncharged but polar groups such as sugars dissolve readily in water because of the stabilizing effect of hydrogen bonds
between hydroxyl groups or carbonyl oxygen of the sugar and the polar water molecules
 Compounds that dissolve easily in water are hydrophilic and those that dissolve poorly are hydrophobic
 Nonpolar solvents such as chloroform and benzene are poor solvents for polar biomolecules but easily dissolve those that are
hydrophobic such as lipids and waxes
 Amphipathic biomolecules have both polar and nonpolar groups

9. 9
 Water dissolves salts such as NaCl by hydrating and stabilizing the Na+ and Cl- ions, weakening the electrostatic interactions
between them and counteracting their tendency to associate in a crystalline lattice
 Same factors apply to charged biomolecules, compounds with charged functional groups such as COO-, NH3+ and
phosphate esters and anhydrides by replacing the solute-solute hydrogen bonds with solute-water hydrogen bonds, thus
screening the electrostatic interactions between solute molecules
 Water is effective in screening the electrostatic interactions because of its high dielectric constant (78.5)(diminishes the
strengths of electrostatic attractions by a factor of 78.5, compared with the same interactions in a vacuum)
F, force of ionic interactions in a solution
Q, magnitude of the charges
R, the distance between the charged groups
Ε, dielectric constant of the solvent in which the interactions occur
F= Q1Q2/εr2
 As salt dissolves the Na+ and Cl- ions leaving the crystalline lattice acquire greater freedom of movement; the resulting increase
in entropy (randomness) is largely responsible for the easy of dissolving salts such as NaCl in water

10. 10

11. N N Nitrogen nonpolar
O O Oxygen nonpolar
O C O Carbon dioxide nonpolar
- -
H H
H
N
-
Ammonia polar
H H
S
Hydrogen sulfide polar
-
11
 Molecules of biologically important gases CO2, O2 and N2 are nonpolar and poorly soluble
 The movement of molecules from the disordered gas phase into the aqueous solution restricts their motion and that of water and the decrease in
entropy combine to make them poorly soluble in water
Electrons are shared equally
Electrons are shared equally
Two dipoles are oppositely
directed and cancel each
other

12.  Water soluble carrier proteins (Hemoglobin and myoglobin) facilitate the transport of O2
 Cabon dioxide forms carbonic acid and is transported as the bicarbonate ion in solution or bound to hemoglobin
12

13. 3. Hydrophobic interactions. The effect of nonpolar compounds in water
 When a nonpolar compound like benzene is mixed with water, two phases form; neither liquid is soluble in the other
 unable to undergo energetically favorable interactions with water and interfere with hydrogen bonding among water
molecules
 Water molecules in the immediate vicinity of nonpolar molecules are constrained in their orientation 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 reduction in entropy is proportional to the surface area of the hydrophobic
solute enclosed with in the cage of water
 Amphipathic molecules contain regions that are polar (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
 the nonpolar regions tend to cluster together to present the smallest hydrophobic area to the aqueous solvent and the
polar regions are arranged to maximize their interaction with water
 These stable structures of amphipathic compounds are called micelles
 The forces that hold nonpolar regions of the molecules are the hydrophobic interactions
13

14. 14

15.  Hydrophobic interactions result from the system’s achieving the greatest thermodynamic stability by minimizing the
number of ordered water molecules required to surround the hydrophobic portion of the solute molecule; hydrophobic
interactions are not due to any intrinsic attraction between nonpolar moieties
 Many biomolecules are amphipathic
 Structures composed of amphipathic molecules are stabilized by hydrophobic interactions among the nonpolar regions
 Hydrophobic interactions between among lipids and between lipids and proteins are the most important determinants
of structure of biological membranes
 Hydrophobic interactions between nonpolar amino acids also stabilize the three dimensional structure of proteins
 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
 Hydrogen bonding between water and polar solutes also causes some ordering of water molecules but the effect is small
15

16. 16

17. 17

18. 18

19. Van der Waal’s interactions
 When two uncharged atoms are brought very close together, their surrounding electron clouds influence each other. As the
nuclei draw close together, their electron clouds begin to repel each other
 Induced transient opposite electric dipoles attract each other; these interactions are the van der Waal’s interactions
Four types of noncovalent interactions
 Hydrogen bonds
 Ionic interaction
 Attraction
 Repulsion
 Hydrophobic interactions
 Va der waal’s interactions
 Although the four types of interactions are individually weak relative to covalent bonds the cumulative effect is of many such
interactions is significant
19

20. 5. Colligative properties of water
 Effect of solute on water is reflected in characteristic changes in behaviour called colligative properties
 Alterations in solvent properties depend only on ther number of particles and not on the chemical nature of the solute
 The effects are
 Freezing point depression
 Boiling point elevation
 Vapour pressure lowering
 Osmotic effects
20
4. Nucleophilic nature of water. Water as a reactant
Chemicals that are electron-rich (nucleophiles) seek electron-deficient chemicals (electrophiles).
• Nucleophiles are negatively charged or have unshared pairs of electrons --> attack electrophiles during substitution or addition
reactions.
Examples of nucleophiles: oxygen, nitrogen, sulfur, carbon, water (weak).
Important in condensation reactions, where hydrolysis reactions are favored.
e.g. protein ------> amino acids
In the cell, these reactions actually only occur in the presence of hydrolases.
•Condensation reactions usually use ATP and exclude water to make the reactions more favorable.

21.  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 glucose.
 Solutes alter the colligative properties of aqueous solutions by lowering the
effective concentration of water. For example, when a significant fraction of
the molecules at the surface of an aqueous solution are not water but solute,
the tendency of water molecules to escape into the vapor phase—that is, the
vapor pressure—is lowered.

22.  Physical properties
 Water is a liquid at standard temperature and pressure.
 It is tasteless, odourless and transparent in the visible electromagnetic spectrum.
 Pure water has three phases: ice, water and steam (or solid, liquid and gas for other
materials) and the phase change temperatures are influenced by the presence of solutes and
polymers present in the liquid
22

23. 6. Ionization of water, weak acids, and weak bases
 Water molecules have a slight tendency to undergo reversible ionization
H2O H+ + OH-
 Reversible ionization is important in cellular function
 The position of equilibrium for any chemical reaction is given by its equilibrium constant, Keq
A + B C + D
The equilibrium constant can be defined in terms of the concentrations of reactants (A and B) and products (C + D) at equilibrium
Keq =
 The equilibrium constant is fixed and characteristic for any given chemical reaction at a specified temperature
23
]
[
]
[
]
[
]
[
B
D
A
C

24.  The degree of ionization of water at equilibrium is small (about 2 out of every 109 molecules in pure water are ionized at any instant)
 The equilibrium constant for the reversible ionization of water is
K eq=
 In pure water at 25oC, the concentration of water is 55.5 M (grams of H2O in 1L divided by its gram molecular weight) and is essentially
constant because of the low concentrations of H+ and OH-
 Substituting 55.5 in the equilibrium constant expression gives;
Keq=
and on rearrangement becomes (55.5M)(Keq)=[H+][OH-] = Kw, the ion product of water
 The value of Keq of pure water at 25oc is 1.8 10-16 M. Substituting this value for in the above equation gives the ion product of water
Kw =[H+][OH-] = (55.5M)(1.8 10-16 M) = 1.0 10-14 M2
 Thus the ionic product of) [H+][OH-] in aqueous solution at 25oC
equals 1.0 X 10-14 M2
24
]
[
]
][
[
2O
H
OH
H 




]
5
.
55
[
]
][
[
M
OH
H 


25.  When there are exactly equal concentrations of both H+ and OH-, as in pure water the solution is at neutral pH
At this pH the concentration of H+ and OH- can be calculated from the ion product of water
Kw =[H+][OH-] = [H+]2
[H+] = =
[H+] = [OH-] =10-7
pH scale
 pH is a convenient way of designating the concentration of H+( and thus of OH-) in an aqueous solution between 1M H+ and
1M OH-
 pH= log = - log [H+] and value in neutral solution is 7
 Alkaline solutions have pH values greater than 7 and acidic solutions pH values less than 7
25
Kw 2
14
10
1 M


]
[
1

H

27.  pH affects the structure and activity of biological molecules eg enzymes
 pH measurements in urine and blood are useful in medical diagnosis
 In patients with uncontrolled diabetes the pH of the blood plasma is below the normal value of 7.4 i.e. acidosis
 In alkalosis the blood plasma pH is above 7.4
Weak acids and bases and dissociation constants
 Weak acids and bases are not completely ionized when dissolved in water
 A proton donor and its corresponding proton acceptor make up a conjugate pair eg acetic acid a proton donor and acetate a proton
acceptor
CH3COOH H+ + CH3COO-
 The stronger the acid the greater the tendency to lose a proton
 The tendency of any acid (HA) to lose a proton and form its conjugate base (A-) is defined by the equilibrium constant (dissociation
constant) for the reversible reaction
HA H+ + A- which is Keq= =Ka
27
Acetate
Acetic acid
]
[
]
][
[
HA
A
H 


28.  Stronger acids have larger dissociation constants than weaker acids
 pKa is analogous to pH and is defined by the equation
pKa = log = – log Ka
 The stronger the tendency to dissociate a proton, the stronger is the acid and the lower its pKa
Titration curves
 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, like NaOH of known concentration
 The sodium hydroxide is added in small increments until the acid is neutralized
 The concentration of the acid can be calculated from the volume and concentration of NaOH added
28
Ka
1

29. 29
 Strong acids (eg HCL) completely dissociate into anions and cations even in strongly acidic solutions
 Weak acids dissociate partially in acidic solutions
 Strong bases – but not weak bases are completely dissociated at high pH
 Many biochemicals are weak acids. Exceptions include phosphorylated intermediates, whose phosphoryl
group contains two dissociable protons, the first of which is strongly acidic
Functional groups that are weak acids have great physiological significance
 Many biochemicals possess functional groups that are weak acids or bases
 Carboxyl groups, amino groups, and the second phosphate dissociation of phosphate esters are present in proteins
and nucleic acids, most
coenzymes and intermediary metabolites
 Knowledge of the dissociation of weak acids and bases is vital to understanding the influence of intracellular pH
on structure and biological activity of these compounds

30. 30
-Relative strengths of weak acids and bases are expressed in terms of dissociation constants (Ka)
eg R–CH2–COOH R–CH2–COO– + H+
Ka = [R–CH2–COO–][ H+]
[R–CH2–COOH]
-Since Ka for weak acids are negatively exponential numbers, Ka is expressed as pKa, where
pKa = –log Ka
-Stronger acid groups have lower pKa values
Note that when, R–CH2–COOH = R–CH2–COO– , then Ka = [H+]
i.e when the associated (protonated) and the dissociated (conjugate base) species are present in equal concentrations, [H+] is
numerically equal to Ka
If logs of both sides of the above equation are taken and both sides multiplied by -1, the expression would be as follows
Ka = [H+]
-log Ka = -log [H+]
Or pKa = pH
i.e. the pKa of an acid group is the pH at which the protonated and unprotonated species are present at equal
concentrations
Or the pH at which it is half dissociated

31. 31
The Henderson – Hasselbalch Equation describes the behaviour of weak acids and buffers
A weak acid HA, ionizes as follows
HA H+ + A–
The equilibrium constant Ka =
Cross multiplication gives
[H+][A –] = Ka [HA]
Divide both sides by [A –]
[H+] = Ka
Log on both sides
log [H+] =
= log Ka +
]
[
]
][
[
HA
A
H 

]
[
]
[

A
HA







]
[
]
[
log
A
HA
Ka
]
[
]
[
log 
A
HA

32. 32
Multiply through by – 1:
– log [H+]= – log Ka –log
Substitute pH and pKa for – log [H] and – log Ka then:
pH = pKa – log
Inversion of the last term removes the minus sign and gives the Henderson-Hasselbalch equation
pH = pKa + log
The equation has great predictive value in protonic equilibria
1. When an acid is exactly half –neutralized, [A –] = [HA]
pH = pKa + log = pKa + log 1 = pKa + 0
Therefore at half neutralization, pH = pKa
]
[
]
[

A
HA
]
[
]
[

A
HA
]
[
]
[
HA
A
]
[
]
[
HA
A

33. 33
2. When the ratio [A–] / [HA] = 100 :1
pH = pKa +
pH = pKa + = pKa + 2
3. When the ratio [A–] / [HA] =1 : 10
pH= pKa + = pKa + (–1)
If the equation is evaluated at several ratios of [A–] / [HA] ranging from 103 to 10–3 and calculated pH values
plotted, the resulting graph describes the titration curve for weak acids
To calculate the pH of a weak acid you need to know (1) the molar proportion of A– (base) to HA (acid), and pKa of
the acid
]
[
]
[
log
HA
A
1
100
log
10
1
log

34. 34
Calculation of points of a titration curve using the Henderson – Hasselbach equation
Theory When a weak acid is titrated by a strong base, the weak acid dissociates in solution to yield a small amount of
H+ ions. When OH– ions are added, they are neutralized by the H+ ions to form water. The removal of the H+
disturbs the equilibrium between the weak acid and its ions. As a result, more HA dissociates to re-establish the
equilibrium. The newly produced H+ ions can then be neutralized by more OH– and so on until all the H+
originally present is neutralized
The overall result is represented by HA + OH– H2O + A–
Titration of 500 mls of a 0.1M weak acid (Ka =10–5, pKa = 5) with a 0.1M strong base (KOH)
1.What is the pH after you add 100 mls of 0.1M KOH?
How many moles HA are present : 0.5 L X 0.1 M = 0.05 moles
How many moles of OH– are added: 0.1L X 0.1M = 0.01 moles
HA ↔ H+ + A–
Starting 0.05
Change 0.05 -0.01 0.01

35. 35
( adding 0.01 moles OH– will titrate an equal amount of H+ ions and pull the equilibrium to the right, increasing the [A–] and decreasing the [HA] by
0.01)
At equilibrium: pH = pKa + log [A–] /[HA]
= 5 + log 0.01 moles/(500+100 mls)
0.04 moles/(500+100 mls)
= 5 + log (0.01/0.04)
= 5 + (-0.6)
= 4.4
2. What is the pH when you add 250 mls of 0.1M KOH?
(i) 0.05 moles of HA are present and (ii) 0.25 L X 0.1M =0.025 moles OH–
added
HA ↔ H+ + A–
Start 0.05
Change 0.05 – 0.025 0.025
At equilibrium pH = pKa + log [A–] /[HA]
= 5 + log 0.025 moles/(500+250 mls)
0.025 moles/(500+250 mls)
= 5 + log 1
= 5 + 0
= 5
NB. Note that pH can also be determined using a pH meter

36. 36
0
0.2
0.4
0.6
0.8
1.0
pH
2 3 4 5 6 7 8
.
meq
of
alkali
added
per
meq
of
acid
pKa
-The pH change per meq of OH- added varies greatly depending upon the pH
-At pH very close to pKa, the solution resists change in pH most effectively. Because the solutions of weak bases
and acids exhibit the phenomenon of buffering – a tendency to resist more effectively a change in pH following
the addition of a strong base or alkali
-Solutions of weak acids and their conjugate bases buffer effectively in the pH range pKa ± 1
One equivalent of an acid or a base is the
weight that contains 1 mole of replaceable
H+ or OH– respectively

37. 37
Acid strength depends on molecular structure
Many acids of biological interest possess more than one dissociable groups
Dissociation occurs in sequential manner
Each step has its own acid/conjugate base pair and a different pKa.
Each conjugate base serves as the acid in the next reaction and is an increasingly more negatively charged
species so that its more difficult to remove a proton
Therefore pKa increases as protons are removed
Eg
Monoprotic acid Lactic pK 3.86
Diprotic acid Carbonic pK1 6.37
pK2 10.25
Triprotic acid Phosphoric pK1 2.15
pK2 6.82
pK3 12.38

38. 38
Importance of buffers in the body
 Buffers are aqueous systems that tend to resist changes in pH when small amounts of Acid (H+) or base (OH-) are added
 A buffer system consists of a weak acid and its conjugate base (more common) or a weak base and its conjugate acid (less
common)
 Most enzymes that catalyze biochemical reactions are very sensitive to pH
 The intracellular and extracellular fluids of multicellular organisms have a near constant pH
 The body fluids must be protected against changes in pH all the time as acids and bases are continuously being produced by
metabolic processes
 The organism’s first line of defense against changes in internal pH is provided by buffer systems

39. 39
Eg the acetic acid –acetate buffer system

COO
CH3
pH
0
0
9
5
0.5 1
Buffering zone between 10%
and 90% titration of a weak
acid
pH 5.76
pH 3.76
pH= pKa=4.76
COOH
CH3
]
COO
CH
[
]
COOH
CH
[ 3
3


OH- added (equivalents)
0 50 100
Percent titrated
The titration curve of acetic acid has a relatively flat zone extending about 1 pH unit on either side of its midpoint pH of 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
zone. This flat zone is the buffering region of the acetic acid–acetate buffer pair

40.  At the midpoint of the buffering region, where the concentration of the proton donor (acetic acid) equals that of the proton
acceptor (acetate), the buffering power of the system is maximal; the pH changes least on addition of H+ and OH-
 The pH at this point equals PKa
 The pH of the acetate buffer system changes only slightly when a small amount of H+ or OH- is added but this change is small
compared with the pH changes that would result if the same amount of acid or base was added to pure water or to the solution
of a salt of a strong acid or strong base which has no buffering power
40

41.  The system is capable of absorbing either H+ or OH
- through reversible dissociation of acetic acid
 Reserve H+ in acetic acid can be released to neutralize addition of OH
- to the system, forming H2O
 This happens because the product [H+][OH-] transiently exceeds Kw (1X 10-14M2)
 The equilibrium readjusts so that this product equals 1X 10-14 M2 at 25oC, transiently reducing the concentration of H+
 As a result the quotient [H+][Ac
-]/[HAc] is less than Ka so HAc dissociates further to restore equilibrium
 Similarly the conjugate base Ac- can react with H+ ions added to the system and the two ionization reactions simultaneously
come to equilibrium
 Buffering action is simply the consequence of two reversible reactions taking place simultaneously and reaching their points of
equilibrium as governed by their equilibrium constants, Kw and Ka
41
Acetic acid
(CH3COOH)
Acetate (CH3COO-)
HAc Ac-
OH- H2O
H+
]
OH
][
H
[
Kw



]
HAc
[
]
Ac
][
H
[
Ka




42.  Biological buffers include HCO3
–, phosphate and proteins which accept or release protons to resist a change in pH eg
 Proteins may contain amino acids with functional groups which may act as weak acids and bases
 Phosphate groups of nucleotides function as weak acids. Etc.
42

43. The phosphate buffer system
 Acts in the cytoplasm of all cells
 Consists of as proton donor and as proton acceptor
 The phosphate buffer system is maximally effective at a pH close to its pKa of 6.86 and thus tends to resist pH changes in the range
between about 5.9 and 7.9
The bicarbonate system
 Blood plasma is in part buffered by the bicarbonate system consisting of carbonic acid (H2CO3) as proton donor and bicarbonate (HCO3
-
) as proton acceptor
H2CO3 H+ + HCO3
-
 The buffer system is more complex than other conjugate acid-base pairs because one of the components, carbonic acid is formed from
dissolved (d) CO2 and water in a reversible reaction
43

4
2PO
H

2
4
HPO

4
2PO
H

2
4
HPO
H+ +
]
CO
H
[
]
HCO
][
H
[
K
3
2
3
1




44. CO2(d) + H2O H2CO3
 Dissolved CO2 is in equilibrium with CO2 of the gas phase
CO2(g) CO2(d)
 The pH of a bicarbonate buffer system depends on the concentration of H2CO3 and HCO3
–, the proton donor and acceptor respectively
 The concentration of H2CO3 depends on the concentration of dissolved CO2, which in turn depends on the concentration of CO2 in the gas phase
(partial pressure of CO2)
 Human blood plasma normally has a pH of about 7.4
 In uncontrolled diabetes there is acidosis and pH can fall to dangerously lower levels resulting in cell damage or death
 Other conditions can increase pH to lethal levels
 The catalytic activity of enzymes is particularly sensitive to variations in pH
44
]
O
H
)][
d
(
CO
[
]
CO
H
[
K
2
2
3
2
2 
)]
g
(
CO
[
)]
d
(
CO
[
K
2
2
3 

45. 45
H+ + HCO3
-
H2CO3
CO2(d)
CO2(g)
H2O
H2O
Reaction 1
Reaction 2
Reaction 3
Gas phase (lung air space)
Aqueous phase (blood
in capillaries)
The bicarbonate buffer system
The buffer system involves three reversible equilibria between
gaseous CO2 in the lungs and bicarbonate (HCO3
-) in the
plasma. When H+ (from lactic acid produced in muscle tissue
during vigorous exercise for example) is added to blood as it
passes through the tissues, reaction 1 proceeds toward a new
equilibria in which the concentration of H2CO3 is increases.
This increases the concentration of CO2 (d) in the blood plasma
(reaction 2) and this increases the pressure of CO2(g) in the air
space of the lungs (reaction 3); the extra CO2 is exhaled
When the pH is raised the opposite events
occur: the H+ concentration of blood plasma
is lowered, causing more H2CO3 to dissociate
into H+ and HCO3
-. This in turn causes more
CO2 (g) from the lungs to dissolve in the
blood plasma

46.  If you breathe more deeply and eliminate more CO2, you can drop the levels of carbonic acid and counteract an influx of acid in
the body (metabolic acidosis) by shifting the equilibrium away from the H+ (compensatory respiratory alkalosis).
 In an opposite way, if the blood pH rises (metabolic alkalosis), by breathing more shallowly, CO2 is retained and the blood pH
decreases (H+ increases) (compensatory respiratory acidosis).
Some clinical disorders in the acid base balance
 Respiratory acidosis: Caused by alveolar hypoventilation and accumulation of CO2 in body, occurs when respiration depth
or rate decreases as seen in airway obstruction, neuromuscular disorders, diseases of the central nervous system or in
chronic obstructive lung diseases like emphysema.
 Respiratory alkalosis: Arises from decreased alveolar PCO2 and is most commonly found as the result of hyperventilation
due to anxiety. Also can be caused by salicylate poisoning, fever, artificial ventilation, and high altitude (since there is a
decrease in total atmospheric pressure and therefore alveolar PCO2 ).
 Metabolic acidosis: Caused by disorders in which excess lactic acid, acetoacetic acid or β-hydroxybutyric acid are
produced, or ingestion of salicylates, ethylene glycol, or methyl alcohol all of which produce strong organic acids.
 Metabolic alkalosis: alkaline materials can not be synthesized from neutral starting materials, so metabolic alkalosis must be
caused by intake of excess alkali (sodium bicarbonate) or abnormal loss of acid (prolonged vomiting).
***c
46

47. Suitability of water for the unique role it plays to organisms
 It is a powerful solvent for polar or ionic substances
 It is a poor solvent for nonpolar substances
 Through hydrophobic interactions lipids coalesce, membranes form
 It has a very high dielectric constant
 Medium for ionization. Ions enhance the variety of chemical species. They also modify chemical reactions
 Thermal properties are specially relevant to the environment
 Has a great capacity as a buffer resisting temperature change
 Heat capacity is very high. It takes a substantial amount of energy to change the temperature and especially the state of
water
 These thermal properties
 allow water to buffer climate through processes like evaporation, melting, condensation and freezing.
 allow effective temperature regulation in living organisms
 Water expands as it is cooled, reducing the density. At temperatures below zero ice freezes on the surface of a body of
water insulating the liquid below
 Water has the highest tension of all common liquids
 Surface tension and density determine how high the liquid raises in a capillary tube
 Capillarity is important in plant life
 Osmotic properties.
 Determine the shape and form of living organisms
47