1. This document provides an introduction to biochemistry and topics related to acids and bases. It defines acids, bases, and ampholytes. 2. Key concepts discussed include the pH scale, factors that influence pH, and the importance of maintaining pH homeostasis in the body. Several buffer systems in the body are described, with bicarbonate identified as the most important buffer. 3. The Henderson-Hasselbalch equation is introduced, which defines the relationship between pH, pK, and the concentrations of a weak acid and its conjugate base. This equation is important for understanding acid-base balance.

1. Welcome
to
The House of BIOCHEMISTRY
1

2. Biophysics
Dr. Mohammad Shiblee Zaman
MBBS, MD (Biochemistry)
Lecturer
Dhaka Medical College

3. 3
Topics
1 5
3 4
2 6
Introduction to
Biochemistry
Law of mass
action
pH, pK and pH
scale
Buffer
Acid, Base
Handerson-
Hasselbach
equation

4. Acid and Base
❑ Thomas Lowry (England)
or JN Bronsted (Denmark)
(1923)
❑ Arrhenius (traditional)
❑ GN Lewis (1923)
4

5. Acid
❑ Acid…..proton donor in aqueous solution
❑ Conjugate base of acid…..remaining anionic part
after proton donation
❑ Strong acid…..quick & complete dissociation
❑ Weak acid…..slow & incomplete dissociation
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6. Base
❑ Base…..proton acceptor in aqueous solution
❑ Conjugate acid of a base…..formed acid after
accepting proton by that base
❑ Strong base…..quick & complete dissociation
❑ Weak base…..slow & incomplete or no
dissociation
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7. Acid, Base contd.
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Fig: Acid & its conjugate base, Base & its conjugate acid

8. Acid, Base contd.
❑ Sources of acids:
1. Protein metabolism specially
sulpher containing AAs
2. Lipid metabolism -
phospholipids
3. Nucleic acids metabolism
4. Glucose metabolism
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❑ Sources of bases:
1. Metabolism of citrus
foods and vegetables
Net production of
acid is more than
base in human body

9. Ampholyte
Metallic hydroxides of alkaline metals like Na and K
which in solution ionize to OH- ions that can bind H+
ion to form H2O
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Alkali
Can act both as an acid and a base
eg. H2O → H+ + OH–
H2O + H+ → H3O+

10. pH
❑ Term introduced by Sorenson in 1909
❑ Negative logarithm of hydrogen ion
concentration of a solution when H+
conc. is expressed in terms of mol/L
pH = – log [H+]
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11. pH contd.
❑ Relationship between pH and [H+].....inverse
❑ ↓ pH of 1 unit.....10 fold ↑ [H+].....vice versa
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12. pH contd.
❑ Normal H+ conc. of blood: 35 – 45 nmol/L
❑ Normal pH of blood: 7.35 - 7.45 (Ave. 7.4)
❑ This pH is maintained by:
 Body fluid buffer system
 Respiratory buffer system
 Renal buffer system
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13. Measurement of pH
❑ Non specific methods:
1. Body fluid buffers
2. Using indicators
3. pH paper
4. Litmus paper
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❑ Specific methods:
1. PH meter
2. pH gas electrode
3. Hydrogen electrode
4. Calomel electrode

14. Importance of pH
❑ Maintenance of optimum pH is essential for life as
most enzymatic reactions occurs at that pH
❑ Maintenance of optimum pH is essential for body
homeostasis
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Optimum pH
Level of pH at which maximum rate of enzymatic
reaction takes place

15. pH scale
❑ Range of pH that covers practical range
of acidity and alkalinity of commonly
used solution expressed as scale
❑ Range: 0 -14
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Fig: pH scale

16. pH scale contd.
❑ Shows relationship between H+, OH- and pH of
aqueous solution at 25 0C when Kw = 10-14
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Fig: Conceptual diagram of pH scale

17. pK
 Negative logarithm of ionization or dissociation
constant (K)
 pH at which an acid is half dissociated, existing as
equal proportions of acid and conjugate base
pK = – log K
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18. Buffer
 Mixture of weak acid & its conjugate base usually
in form of salt
 Has ability to resist change of pH of a solution when
moderate amount of acid or base is added
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Buffer =
Conjugate base
or,
Salt
Weak acid Weak acid

19. Buffer system Conjugate base/ salt Weak acid pK
Bicarbonate** HCO3
– or NaHCO3 H2CO3 6.1
Phosphate HPO4
2– or Na2HPO4 H2PO4- or NaH2PO4 6.8
Plasma protein Pr- HPr 7.3
Hemoglobin Hb– HHb 7.3
Oxyhemoglobin HbO2
– HHbO2 7.3
Ammonia NH3 NH4
+ 9.0
Common buffer systems
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20. Compartment Buffers (in order of importance)
ECF/ Plasma Bicarbonate, Phosphate, Protein
ICF Protein, Bicarbonate, Phosphate
RBC Hemoglobin, Phosphate, Bicarbonate
Blood Bicarbonate, Hemoglobin, Phosphate,
Protein
Urine Ammonia, Phosphate, Bicarbonate
Distribution of body buffers
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21. Mechanism of buffer action
❑ Buffer acts by converting
 strong acid into weak acid
 strong base into weak base or neutral salt
 Thereby minimizes effect of strong acid or strong
base on pH of a solution
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22. Mechanism of buffer action contd.
22
HCl
H2CO3 NaCl
NaHCO3
CO2 H2O

23. Mechanism of buffer action contd.
23
NaOH
NaHCO3
H20
H2CO3

24. Alternatively,
❑ Buffer acts by
 donating proton in proton deficit (alkalosis)
 accepting proton in proton excess (acidosis)
 Thereby keep acid/base tied up to minimize pH change 24
HB
H+
B-
HB
H+
B-

25. Bone Buffer: Hydroxyapatite crystal (HAC)
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Ca10(PO4)6(OH)2 Ca9(PO4)6 + Ca++ + 2H2O
2H+
In persistent acidosis, when body fluid buffer fails

26. Buffering capacity
❑ Amount of acid or base required to produce one unit
change of pH in solution
❑ Smaller the pH change, greater the buffering capacity
❑ Depends on
 pK of buffer – maximum capacity when pH = pK
 Base/acid ratio – ratio closer to 1, more is the capacity
 Total amount – more effective at higher concentration
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27. Bicarbonate Buffer: most effective
❑ High buffering capacity (60% of total)
❑ Wide field of buffering activity (both in ECF & ICF)
❑ Conjugate base (HCO3
-) conc. (24 mmol/L) is 20
times more than its acid (H2CO3) (1.2 mmol/L)
❑ Works synergistically with Hb buffer system
❑ Open end buffer system
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29. Importance of buffer systems
❑ Maintain normal pH of body as optimum pH is
essential for normal enzymatic activity
❑ Regulate narrow limit of body pH, compatible for life
of most cells
❑ Determine pH with indicator
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30. Law of mass action
Rate of any reversible chemical reaction at any
instant, at a given temperature is directly
proportional to the product of molar concentration of
the reactants
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A + B
V1
→
C + D
V = velocity
←
V2
[ ] = molar conc.

31. Explanation
According to Law of mass action:
V1 α [A][B], So, V1= K1 [A][B]
V2 α [C][D], So, V2= K2 [C][D]
At equilibrium, V1= V2
So, K1 [A][B] = K2 [C][D]
K1/ K2= [C][D] / [A][B]
K = [C][D] / [A][B]
K= Equilibrium constant or dissociation constant
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32. Henderson-Hasselbalch Equation (HHE)
❑ Define relationship between pH, pK and
concentration of an acid and its conjugate base
❑ Explain relationship between pH of a weak acid in
solution and its dissociation constant
32
pH = pK + log
Conjugate base or salt
Weak acid

33. Explanation
Dissociation of a weak acid can be represented as
HA (weak acid)↔H+ + A- (Conjugate Base)
According to Law of mass action,
K = [H+][A-] / [HA] (K=dissociation constant of acid)
[H+][A-] = K[HA] (by cross multiply)
[H] = K [HA] / [A-] (divide both side by [A-])
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34. Explanation contd.
log [H+] = log {K[HA] / [A-]} (taking log on both side)
log [H+] = log K + log [HA] / [A-]
- log [H+] = - log K - log [HA] / [A-] (multiply by -1)
pH = pK - log [HA] / [A-]
pH = pK + log [A-] / [HA]
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So, pH = pK + log
Conjugate base
Weak acid

35. Importance of HHE: used to
❑ determine the pH of weak acid
❑ determine the pH of buffer solution
❑ quantify the buffer components needed to prepare a
buffer solution of definite pH
❑ determine blood pH & to evaluate the acid base
status of patients
❑ spell out the concept of acidosis & alkalosis
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