Summarised lecture for undergraduates (beginners) in the course of Biochemistry

1. Water, pH scale
Acid–Base Balance
Buffers systems
Dr. Male Keneth

2. Water
2

3. Water
•65% of our bodies are water; cell are
70-95% water.
•70% of our earth is ocean water.
•3 different forms of water:
• Solid (ice), liquid (water), and gas (steam)

4. Water
1. Electrons of each covalent bond are not
shared equally between the H & O atoms.
2. Oxygen has a stronger attraction.
3. The unequal sharing causes the Oxygen to
have a slightly negative (-) charge, while
the hydrogen ends have a slightly positive
(+) charge.

5. Polarity
•A molecule in which opposite ends have
opposite electrical charges is called a
POLAR molecule for example water.

6. Water
•Formula for water:
•H2O
•H atoms are joined to O atoms
by a single covalent bond

7. Unusual and unique
properties of water
• Boiling Point and Freezing Point
• Surface Tension, Heat of
Vaporization, and Vapor Pressure
• Viscosity and Cohesion
• Solid State
• Liquid State
• Gas State
• Solvent

8. Boiling Point and Freezing
Point
• Bp and Fp decrease as molecular size decreases,
but not for water:
• Water requires more energy to break its hydrogen bonds
before it boils/ freezes:
• Vital in living organisms living in water; as they could die
instantly
• Also vital for cooling body temperature via sweating.

9. Surface Tension, Heat of
Vaporization, and Vapor
Pressure
•S.T: high, due to the hydrogen bonds.
•HoV: very high.
•V.P: inversely proportional to
intermolecular forces: water has strong
intermolecular forces = low V.P

10. Viscosity and Cohesion and
Adhesion
• Viscosity: the property of fluid having high
resistance to flow: water is viscous due to
stronger intermolecular forces
• Cohesion: intermolecular forces between
like molecules; this is why water molecules
are able to hold themselves together in a
drop.
• Due to its polarity.
• Importance : ????

12. Solid State and Liquid state
• Water is more dense than ice: why?
• As water cools below 40C, hydrogen bonds
rearrange themselves into an open crystalline,
hexagonal structure, making water molecules to
be held further apart
• Molecules are held tightly packed in water’s
liquid state: being held by hydrogen bonds,
moving freely.
• Why does this imply? Why is it important
for living organisms?

13. Gas state
• As water boils, the hydrogen bonds are
broken as particles move faster.
• Lack of hydrogen bonds explains why steam
is causes worse burns water.
• All the energy to break the hydrogen bonds is
contained in steam.
• And as stem is converted to liquid water, more
heat energy is released.

14. Universal Solvent
• Because of it’s polarity, it dissociates many
particles.
• Oxygen has a slightly negative charge, while
the two hydrogens have a slightly positive
charge.
• The slightly negative particles of a compound
will be attracted to water's hydrogen atoms
• the slightly positive particles will be attracted to
water's oxygen molecule; this causes the
compound to dissociate

15. Water dissolves polar
compounds
15
solvation shell
or
hydration shell

16. Hydrogen Bonding of Water
16
Crystal
lattice
of
ice
One H2O molecule can
associate with 4
other H20 molecules

17. Non-polar substances are insoluble in water
17
Many lipids are amphipathic

18. Micelles
• An aggregate of molecules in a colloidal
solution, such as those formed by
detergents.
• A typical micelle in aqueous solution forms
an aggregate with the hydrophilic "head"
regions in contact with surrounding solvent,
sequestering the hydrophobic single-tail
regions in the micelle center.

19. How detergents work?
19

20. IONISATION OF WATER
• Keq, Kw and pH
• H2O is the medium of biological systems.
• Dissociation of ions from biological molecules occurs in
H2O.
• Water is essentially a neutral molecule but will ionize to
a small degree.
20

21. Ionization of Water
21
Keq= [H+] [OH-]
[H2O]
H20 H+ + OH-
Keq=1.8 X 10-16M
[H2O] = 55.5 M
[H2O] Keq = [H+] [OH-]
(1.8 X 10-16M)(55.5 M ) = [H+] [OH-]
1.0 X 10-14 M2 = [H+] [OH-] = Kw
If [H+]=[OH-] then [H+] = 1.0 X 10-7
This equilibrium can be calculated as for any reaction:

22. • Since the concentration of H2O is very high (55.5M)
relative to [H+] and [OH–], it is generally removed from
the equation by multiplying both sides by 55.5 yielding a
new term, Kw (ion product/ dissociation constant):
Kw = [H+][OH–]
• In pure water, to which no acids or bases have been
added:
Kw = 1 x 10–14 M2 at room temperature.
Hence 10–14 = [H+][OH–]
22

23. • From 10–14 = [H+][OH–] and
• taking log; –14 log10 = log[H+] + log[OH–]
• -14 = log[H+] + log[OH–]
• 14 = -log[H+] - log[OH–]
• Therefore, 14 = pH + pOH
• At neutrality, [H+] = [OH-] = 1 x 10–7 M
• Thus pH = 7, and also pOH = 7.
• Note: pH is the negative log of hydrogen ion
concentration of any solution.
23
pH = –log[H+]
pOH = -log[OH–]

24. pH Scale
24
 Devised by Sorenson (1902)
 [H+] can range from 0M and
1 X 10-14M
 using a log scale simplifies
notation
 pH = -log [H+]
Neutral pH = 7.0

25. Biological significance of pH
• H+ is one of the most important ions in biological systems.
Its conc. affects cellular and organismal processes as all
metabolic processes are pH dependent.
1. Structure and function of proteins depends on ionic
interactions which are pH dependent.
2. Movement of ions across membranes depends upon their
net charge as determined by pH.
3. The ionic state of the nucleic acids, lipids,
mucopolysaccharides are determined by pH.
4. All enzymes function best within optimum pH range.
5. O2 & CO2 transport, release/gaseous exchange are pH
dependent.
6. The [H+] also plays a role in energy generation and
endocytosis. 25

28. Acid–Base Balance
28

29. Acid–Base Balance
•pH of body fluids is altered by the
introduction of acids or bases
•Acids and bases may be strong or weak
• Strong acids dissociate completely (only HCl
is relevant physiologically)
• Weak acids do not dissociate completely and
thus affect the pH less compared to strong
acids (e.g. carbonic acid)
29

30. Weak Acids and Bases Equilibria
30
•Strong acids / bases – dissociate completely
•Weak acids / bases – dissociate only partially
•The pH depends on the degree of dissociation
•Enzyme activity is sensitive to pH
•Weak acid/bases play important role in
protein structure/function.

31. Ka
• In biological processes various weak acids and bases are
encountered, e.g. the acidic and basic amino acids,
nucleotides, phospholipids, etc.
• Weak acids and bases in solution do not fully dissociate
and, therefore, there is an equilibrium between the acid
and its conjugate base.
• This equilibrium can be calculated and is termed the
equilibrium constant = Ka.
• Ka is also referred to as the dissociation constant.
31

32. Acid/conjugate base pairs
32
HA + H2O A- + H3O+
HA A- + H+
HA = acid ( donates H+)
A- = Conjugate base (accepts H+)
Ka = [H+][A-]
[HA]
Ka & pKa value describe tendency to
loose H+
large Ka = stronger acid
small Ka = weaker acid
pKa = - log Ka

33. pKa values determined by titration
33

34. Phosphate
has three
ionizable H+
and three
pKas
34

35. Buffers
• Buffers are aqueous systems that resist changes in pH
when small amounts of a strong acid or base are
added.
• A buffered system consist of a weak acid and its
conjugate base.
• Buffers are effective at pHs that are within +/-1 pH unit
of the pKa
• The most effective buffering occurs at the region of
minimum slope on a titration curve
(i.e. around the pKa).
35

36. Buffers
• At pKa the pH of a solution does not change
appreciably even when large amounts of acid
or base are added. This phenomenon is known
as buffering.
• Dissolved compounds that stabilize pH by
providing or removing H+
• Weak acids or weak bases that absorb or release
H+
are buffers
36

37. Buffer Systems
• Buffer System: consists of a combination of a weak
acid and the anion released by its dissociation (its
conjugate base)
• The anion functions as a weak base:
H2CO3 (acid)  H+ + HCO3
-
(base)
• In solution, molecules of weak acid exist in
equilibrium with its dissociation products
(meaning all three species exist in plasma)
37

38. Henderson-Hasselbach Equation
38
1) Ka = [H+][A-]
[HA]
2) [H+] = Ka [HA]
[A-]
3) -log[H+] = -log Ka -log [HA]
[A-]
4) -log[H+] = -log Ka +log [A-]
[HA]
HA = weak acid
A- = Conjugate base
* This equation describes
the relationship between
pH, pKa and buffer
concentration
5) pH = pKa + log [A-]/ [HA]
pH values of buffered solutions can be calculated using
Henderson-Hasselbalch equation

39. Case where 10% acetate ion, 90%
acetic acid
39
• pH = pKa + log(90)
(0.1)
[0.9]
• pH = 4.76 + (-0.95)
• pH = 3.81

40. • pH = pKa + log10
[0.5 ]
¯¯¯¯¯¯¯¯¯¯
[0.5]
• pH = 4.76 + 0
• pH = 4.76 = pKa
• Ie. at neutral point pH is equal
to pKa
• Also, there is maximum buffering
Case where 50% acetate ion 50% acetic
acid
40

41. • pH = pKa + log10
[0.9 ]
¯¯¯¯¯¯¯¯¯¯
[0.1]
• pH = 4.76 + 0.95
• pH = 5.71
Case where 90% acetate ion 10%
acetic acid
41

42. Alkaline conditions
• pH = pKa + log10
[0.99 ]
¯¯¯¯¯¯¯¯¯¯
[0.01]
• pH = 4.76 + 2.00
• pH = 6.76
• Acidic conditions
• pH = pKa + log10
[0.01 ]
¯¯¯¯¯¯¯¯¯
[0.99]
• pH = 4.76 - 2.00
• pH = 2.76
Cases when buffering fails
42

43. • Question 1. Calculate the pH of a solution
containing a mixture of 0.25M acetic acid
and 0.1M sodium acetate. The pKa of
acetic acid is 4.76.
• Question 2. Calculate the ratio of lactic
acid to lactate required in a buffer system
of pH 5. The pKa of lactic acid is 3.86.
43

44. Buffer Systems in Body Fluids
44
Figure 27–7

45. 3 Major Physiological Buffer
Systems
1. Protein buffer systems:
• Help regulate pH in extracellular fluid (ECF) and
intracellular fluid (ICF)
• Interact extensively with other buffer systems
2. Carbonic acid–bicarbonate buffer system:
• Most important buffer of blood (ECF)
3. Phosphate buffer system:
• Buffers pH of ICF and urine
45

46. 1. Protein Buffer Systems
• Amino acids in protein buffer systems
• Depend on free and terminal amino acids
• Respond to pH changes by accepting or releasing H+
• If pH rises:
• carboxyl group of amino acid dissociates, acting as weak acid, releasing
a hydrogen ion
• If pH drops:
• carboxylate ion and amino group act as weak bases
• accept H+
• form carboxyl group and amino ion
46

47. Protein buffer (most have pKa =
7.4)
• These include Hemoglobin, serum albumins
and other plasma proteins, proteins in
interstitial fluid and in the intracellular fluid
(ICF)
• Several of these groups have pKa value around 7.4.
• Since proteins are present in significant
concentrations in living organisms, they are
important powerful buffers e.g.
• The Hemoglobin, most abundant molecule in red blood
cells. Because of its structure and cellular concentration,
Hemoglobin plays a major role in maintaining blood pH.
47

48. The Hemoglobin Buffer System
• CO2 diffuses across RBC membrane:
• no transport mechanism required
• As carbonic acid dissociates:
• bicarbonate ions diffuse into plasma
• in exchange for chloride ions (chloride shift)
• Hydrogen ions are buffered by hemoglobin molecules
• The only intracellular buffer system with an
immediate effect on ECF pH
• Helps prevent major changes in pH when plasma PCO
2
is rising or falling
48

52. 2. The Carbonic Acid–
Bicarbonate Buffer System
• Formed by carbonic acid and its dissociation
products .
• Prevents changes in pH caused by organic acids and
fixed acids in ECF/blood
• H+ generated by acid production combines with
bicarbonate in the plasma.
• This forms carbonic acid, which dissociates into CO2
which is breathed out
52

53. 53
• It is a carbonic acid/bicarbonate (H2CO3/HCO3
-)
buffer system.
• It is the most important buffer system in the body
(plasma and ECF) despite the fact that bicarbonate
has pKa of 6.1 far below the physiological pH 7.4
• It is so important because it responds very fast to
changes in plasma pH through loss of carbon
dioxide in the lungs and bicarbonate in the urine
through the kidneys.
• CO2 reacts with water to form carbonic acid

54. • From 7.4 = 6.1 + log [bicarbonate ]
[Carbonic acid]
• The ratio of HCO3
- to CO2 required to maintain the blood
pH of 7.4 is regulated:
• Carbon dioxide conc. is adjusted by changes in the rate of
respiration
• Bicarbonate conc. is regulated by the kidneys; If the [HCO3
-]
decreases, the kidneys remove H+ from the blood
triggering a shift to the right; increasing [HCO3
-].
• When excess HCO3
- ions are produced, they are excreted by
the kidneys triggering a shift to the left.
54

55. Figure 27–9
The Carbonic Acid–
Bicarbonate Buffer System
55

56. Limitations of the Carbonic
Acid Buffer System
1. Hard to protect ECF from changes in pH
that result from elevated or depressed
levels of CO2 (because CO2 is part of it)
2. Functions only when respiratory system
and respiratory control centers are
working normally
3. Ability to buffer acids is limited by
availability of bicarbonate ions
56

57. The Phosphate Buffer
System
•Consists of anion H2PO4
— (a weak
acid)
•Works like the carbonic acid–
bicarbonate buffer system.
•Is important in buffering pH of ICF.
57

58. 3. Phosphate buffer (pKa =
7.2)
• Phosphate buffer consists of weak acid conjugate base pair
(H2PO4
-/ HPO4
2-)
H2PO4
-  H+ + HPO4
2-
• Dihydrogen phosphate Hydrogen phosphate
• It has a pKa of 7.2, the blood pH is 7.4, close to the phosphate
buffer pKa of 7.2: looks a perfect buffer
• but the concentration of H2PO4
- and HPO4
2- in blood are too low to
have a major effect.
• However, the concentration of phosphate buffer in
intracellular fluid is higher (approx. 75mEq/L) than it is in
blood (4mEq/L)
• It is therefore an important buffer in intracellular fluid (ICF).
58

59. Problems with Buffer
Systems
•Provide only temporary solution to
acid–base imbalance.
•Do not eliminate H+ ions.
•Supply of buffer molecules is
limited.
59

60. Maintenance of Acid–Base
Balance
• Requires balancing H+ gains and losses
• For homeostasis to be preserved, captured H+ must
either be:
• permanently tied up in water molecules through CO2
removal at lungs OR
• removed from body fluids through secretion at kidney
• Thus, problems with either of these organs cause
problems with acid/base balance
• Coordinates actions of buffer systems with:
• respiratory mechanisms
• renal mechanisms
60