buffers Ph acid base

3. Aaj ki Achee Baat

4. PH AND BUFFER
Acid Base Balance
BY
DR SHAMIM AKRAM
Assoc. Prof.Biochem
AIMEC

5. Syllabus For MBBS:
Acid Base Balance and Body Fluids
1. Ionization of water, week acids and bases
2. pH and pH scale
3. pK values, dissociation constant and titration curve of
week acids
4. Body buffers and their mechanism of action
5. Henderson – Hesselbach equation
6. Acid base regulation in human body
7. Biochemical mechanisms for control of water and
electrolyte balance
8. Types of particles in solution
9. Importance of selectively permeable membranes,
osmosis and osmotic pressure, surface tension, viscosity
also in relation to body fluids

6. Syllabus For BDS:

7. Reference Books For PH & Buffers
Biochemistry vol 1 by Mushtaq Ahmad ch.2,p.15-30
Lippincott ch.1,p.6-9 (Applications of H-H Equation)
Harper CH.2, P.8-12(Great Physiologic significance of
functional gps.that are weak acids)

8. Water-Functions
60% of body weight—42 L---2/3 IC & 1/3 EC
Solvent of life —due to dipolar property
Act as Reactant & Product in many metabolic reactions
Important for transport of Molecules & Heat
Many chemical groups (acid/base) produced in the
body,dissolve in water releasing/accepting H+ to maintain
homeostasis.
Disturbances in acid base & water electrolyte balance may
threat life by producing acidosis and dehydration.
Important role in Thermal Regulation.

9. Water and Thermal Regulation
1-Water has high Heat of Fusion: Large decrease in temp.
required to convert water to ice.
2-High Thermal Conduction: Fascilitate heat discipation
e.g.from brain to blood.
3-High capacity & Heat of Vaporization: Evaporization
from skin has cooling effect and total body water pool.

10. Water Movement-Osmosis
-Water can move freely through semipermeable cell
membrane via ion channels/aquaporins.
-Water moves from a compartment of low solute conc. to
high solute conc.to achieve an equal osmolality on both
sides.
(Osmolality is the conc. of solute/Kg water measured in
m osmol/Kg water)
-The force keeping this equilibrium is the Osmotic
Pressure exerted by plasma proteins.
(e.g. ↓ albumin lead to oedema in nephrotic syndrom)
-Similarly loss of cellular water through urine (polyurea) in
uncontrolled DM is due to hyperglycemia.

11. Dissociation of water
Dissociation of water is very slow,one molecules out of 107
(10 lacs) molecules dissociates.
H2O H+ + OH-
Ionic product H+ X OH- is dependent upon temperature,it
increases 8% by raising temp. by 10C.
At 250C, it has been found to be equal to 10-14 gram mole/lit
It is a constant value expressed by Kw = 10-14 gram mole/lit.
Pure water is neutral,so conc. of H+ = OH-
Because H+ X OH- = 10-14 gram mole/lit
So OH- is equal to 10-7 gram mole/lit
And H+ is equal to 10-7 gram mole/lit

12. Dissociation of water
SORENSEN suggested that exponent of H+ conc. 10-7
expressed as negative sign should be changed to positive
sign,be termed pH (P denoted’’ -ve log of’’ H+ conc.)
By taking –ve log of H+ conc. & changing sign
PH = -(log H+)
For water PH = - (log 10-7) (bcz log 10-7 = -7)
= - (-7)
= 7
So pH of water is = 7

13. ACID BASE BALANCE
PH
It is the negative log of the hydrogen ion concentration.
pH = -log [H+]
 pH is a unit of measure which describes the degree of
acidity or alkalinity of a solution.
 It is measured on a scale of 0 to 14.
 Low pH values correspond to high concentrations of H+
and high pH values correspond to low concentrations of
H+.

14. PH VALUE
 The pH value of a substance is directly related to the
ratio of the hydrogen ion and hydroxyl ion
concentrations.
 If the H+ concentration is higher thanOH- the material
is acidic.
 If theOH- concentration is higher than H+ the material
is basic.
 7 is neutral, <7 is acidic, >7 is basic

15. THE PH SCALE
 The pH scale corresponds to the concentration of
hydrogen ions.
 For example pure water has H+ ion concentration 10-7
gram mole/litre at 250C, therefore the pH would then be 7.

16. pH of Different Fluids

17. Acid
Any compound which forms H⁺ ions in solution
(proton donors)
eg: Carbonic acid releases H⁺ ions
Base
Any compound which combines with H⁺ ions in
solution (proton acceptors)
eg:Bicarbonate(HCO3⁻) accepts H+ ions

18. ACID–BASE BALANCE
 Normal pH : 7.35-7.45
 Acidosis
 Physiological state resulting from abnormally low plasma
pH
 Alkalosis
Physiological state resulting from abnormally high plasma
pH
 Acidemia: plasma pH < 7.35
 Alkalemia: plasma pH > 7.45

19. MEASUREMENT OF PH
The pH can be measured by:
ph strips
Ph indicators
Ph meter

20. ph strips
A pH test strip is a strip of litmus paper with which you can
measure the pH value of a liquid. The substance in the
paper causes the paper to show a different colour at
different acidities.
How to use a pH strips
Dip the pH test strip for two seconds in the fluid, wait for
ten seconds.
Because the strip contacts an acidic or alkaline substance,
the strip discolours.
The more acidic the liquid, the more red the strip becomes,
and the more alkaline the liquid, the more blue the strip
becomes.
Using the pH scale with the different colours on the box,
you can determine how acidic or alkaline the liquid is.

22. SOME IMPORTANT INDICATORS USED IN A CLINICAL
BIOCHEMISTRY LABORATORY ARE LISTED BELOW:
sr,.
No.
INDICATOR Ph range Colour in
acidic ph
Colour in
basic ph
1 Phenophthalein 9.3-10.5 colourless pink
2 Methyl orange 3.1-4.6 red yellow
3 Bromophenol blue 3.0-4.6 yellow blue
4 Methyl red 4.4-6.2 Red yellow
5 Phenol red 6.8 – 8.4 yellow red
6 Litmus 4.5-8.3 red Blue

23. PH METER
 The pH meter is a laboratory equipment which used to measure
acidity or alkalinity of a solution
 The pH meter measures the concentration of hydrogen ions [H+]
using an ion-sensitive electrode.
 It is the most reliable and convenient method for measuring
ph,even of coloured solutions.
 But it is costly and require large sample volume.

25. BUFFER
DEFINITION
A buffer solution is a solution which resists changes in pH when
a small amount of acid or base is added.
COMPOSITION
Typically a mixture of a weak acid and a salt of its conjugate
base or weak base and a salt of its conjugate acid.

26. TYPES OF BUFFERS
Two types :
Acidic buffers are solutions that have a pH below 7 and
contain a weak acid and one of its salts.
For example, a mixture of acetic acid and sodium acetate
acts as a buffer solution with a pH of about 4.75.
Alkaline buffers, on the other hand, have a pH above 7
and contain a weak base and one of its salts.
For example, a mixture of amonium hydrooxide and
amonium chloride act as buffer solution with a pH of about
10.

27. TYPES OF BUFFERS
ACIDIC BUFFERS
Solution of a mixture of a weak acid and a salt of this weak
acid with a strong base.
( weak acid )
E.g. CH3COOH + CH3COONa
( Salt )
BASIC BUFFERS
Solution of a mixture of a weak base and a salt of this weak
base with a strong acid.
e.g. NH4OH +
(Weak base)
NH4Cl
(Salt)

28. HOW BUFFERS WORK
 Equilibrium between acid and base.
 Example:ACETATE BUFFER
 CH3COOH  CH3COO- + H+
 If more H+ is added to this solution, it simply shifts the
equilibrium to the left, absorbing H+, so the [H+]
remains unchanged.
 If H+ is removed (e.g. by addingOH-) then the
equilibrium shifts to the right, releasing H+ to keep the
pH constant

29. Bicarbonate Buffer-NaHCO3/H2CO3
Suppose HCl is added to a buffer system containing
NaHCO3 and H2CO3
NaHCO3+HCl------- H2CO3+NaCl
So HCl, a strong acid,reacts with base to give a weak acid
& a neutral salt,so negligible change in the pH.
Similarly,If NaOH is added to this buffer soln.
NaOH +H2CO3----- NaHCO3+H2O
Since NaHCO3 is much weaker base than NaOH,resulting
rise in pH will be quite small.
Buffer solutions act best at a certain pH,where their
buffering power is the greatest.For example aceticacid/Na
acetate has max. buffering power at pH 4.76

30. HANDERSON HASSELBALCH
EQUATION
LawrenceJoseph Henderson wrote an equation, in 1908,
describing the use of carbonic acid as a buffer solution.
KarlAlbert Hasselbalch later re-expressed that formula in
logarithmic terms, resulting in the Henderson–
Hasselbalch equation.

31. Ka =
[H+] [A-]
[HA]
take the -log on both sides
The Henderson-Hasselbalch Equation derivation for buffer solutions
a
-log K = -log [H+] -log
[A-]
[HA]
a
pH = pK + log [A-]
[HA] a
= pK + log
[Proton acceptor]
[Proton donor]
HA H+ + A-
a
pK = pH -log [A-]
[HA]
apply p(x) = -log(x)
and finally solve for pH…

32. Demonstration of Buffering action of a weak
acid &its Conjugate base-Aceticacid/Naacetate
10 ml of 0.1 M soln.of Acetic acid is taken in a flask
& its pH noted.
NaOH soln. of same strength(0.1M) is added in small
measured amounts at a time & rise in pH is noted after each
addition till 10 ml consumed.The results are plotted as shown
by titration curve.Note that:
The rise in pH is faster in the beginning,but from 3.8 it
slows,till 5.8,where it become fast again.
These values show that buffering power is less at pH value
below3.8 & above 5.8.
The region b/w pH 3.8 &5.8 represent strong buffering activity
At pH value of 4.8,one half of the acetic acid had been
dissociated & half undissociated,is called Dissociation
Constant or pKa.So max buffering seen at pKa value ±1.

33. How to prepare 0.1 molar solution of acetic
acid
The molecular weight of CH3COOH is 60.05 g/mole.
To prepare 1 lit of 1 Molar soln. you will add
60.05g (57.2 ML) to water to make 1 Lit total volume.
To prepare 1 L of 0.1M solution, you need 0.1 mol acetic
acid, which is 60.05/10= 6.005g.
1.05 grams of acetic acid is equivalent 1 milliliter.
6.005g of acetic acid is =6.005/1.05 = 5.72 Ml.
So you need 5.72 mL glacial acetic acid which you need to
fill up to 1L.
Similarly for NaOH(40g/mole),.1 M soln is prepared.

35. Buffer Capacity Depends Upon:
 The buffer capacity of a buffer soln. depends
upon conc. of buffer and its pka.
 The greater the conc. of buffer the less the pH
changes upon addition of H+ or OH-
 Choose a buffer whose pKa is closest to the desired
pH (e.g.for blood near to 7.35-7.45).
pH should be within pKa ± 1

36. Applications of The Henderson-
Hasselbalch Equation
The Henderson-Hasselbalch equation is widely used by
many scientists especially chemists, biologists pharmacists
and in bakeries as well.
pH = pKa + log ([A–]/[HA])
In the above equation,unknowns can be calculated i.e.
The pH of the solution- The pKa of a chemical in solution- The
[A-] or salt concentration and the [HA] or acid concentration.

37. Applications of The H-H Equation
Separation of plasma proteins by charge typically is done
at a pH above the isoelectric pH ( pI)of the major proteins.
The charge on the proteins is negative.In an electric field,
the proteins will move toward the positive electrode at a
rate determined by their net negative charge.
Variations in the mobility pattern are suggestive of certain
diseases.

38. Applications of The H-H Equation
The Henderson-Hasselbalch equation can be used to
calculate how the pH of a physiologic solution responds to
changes in the concentration of a weak acid and/or its
corresponding salt form.
For example, in the bicarbonate buffer system, the
Henderson-Hasselbalch equation predicts how shifts in the
bicarbonate ion concentration, [HCO3 −], and the carbon
dioxide concentration [CO2] influence pH (Fig.).

39. Bicarbonate As A Buffer

40. The equation is also useful for calculating the abundance
of ionic forms of acidic and basic drugs.
For example, most drugs are either weak acids or weak
bases (Fig.).
Acidic drugs (HA) release a H+, causing a charged anion
(A−) to form.
Weak bases (BH+) can also release a H+. However, the
protonated form of basic drugs is usually charged, and the
loss of a proton produces the uncharged base (B).

41. A drug passes through membranes more readily if it is
uncharged. Thus, for a weak acid, such as aspirin, the
uncharged HA can permeate through membranes, but A−
cannot.
Likewise, for a weak base, such as morphine, the
uncharged B form permeates through the cell membrane,
but BH+ does not.
Therefore, the effective concentration of the permeable
form of each drug at its absorption site is determined by
the relative concentrations of the charged (impermeant)
and uncharged (permeant) forms.

42. Drug Absorbtion

43. The ratio between the two forms is determined by the pH
at the site of absorption and by the strength of the weak
acid or base, which is represented by the pKa of the
ionizable group.
The Henderson-Hasselbalch equation is useful in
determining how much drug is found on either side of a
membrane that separates two compartments that differ in
pH, for example, the stomach (pH 1.0–1.5) and blood
plasma (pH 7.4).

44. Other Pharmaceutical Applications
Explaination of how excess acidity in stomach can be
reduced by use of different bases.
Explaination of how compounds such as ranitidine(Zantac-
block histamine that activate proton pump.) can be used to
inhibit acid production.
Explaination of how compounds such as omeprazole&
Esoomeprazole can be used to supress acid secretion in
the stomach.

45. TODAYS ASSIGNMENT
An important buffer system of the blood plasma is
NaHCO3/ H2CO3,their ratio under normal circumstances is
20/1.Find out the pH of plasma.The pKa of H2CO3 is 6.1

46. Water,Electrolyte and Acid-Base Balance
Organisms have tremendous capacity to survive against
odds and maintain homeostasis,sp. Imp. For
Water,Electrolyte and Acid-Base Status of the body

47. Water is solvent of life-Functions
.1-Provide aqueous medium for many
biochemical reactions
2-As reactant in many metabolic
reactions.
3-As vehicle for transport of solutes.
4-Regulation of body temperature

48. Distribution of water

49. Water turnover and Balance-Intake/Output

50. Electrolyte Balance
Electrolytes are compounds which readily dissociate in
solution and exist as ions,positively or negatively charged
particles.e.g. NaCl does not exist as such but as cation,Na+
or anion Cl-.
The concentration of electrolytes are expressed as
Milliequivalents per litre (mEq/l)
Electrolytes are well distributed in the body fluids in order
to maintain the osmotic equilibrium and water balance.
The total concentration of cations and anions in each body
compartment (ECF or ICF) is equal to maintain electrical
neutrality.

51. Na+ is the principal extracellular cation while K+ is the
intracellular cation. This difference in the concentration is
essential for the cell survival,which is maintained by Na+ -
K+ pump.
As regards anions, Cl- and HCO3- predominantly occur in
Extracellular fluids, while HPO4- , proteins and organic
Acids are found in the intracellular fluids.

53. ANION GAP
The total concentration of cations and anions (expressed as
mEq/l) is equal in the body fluids to maintain electrical
neutrality.
The commonly measured electrolytes in the plasma are Na+,
K+, Cl- and HCO3-.
Na+ and K+ together constitute about 95% of the plasma
cations. Cl- and HCO3- are the major anions, contributing to
about 80% of the plasma anions.
The remaining 20% of plasma anions (not normally measured
in the laboratory) include proteins, phosphate, sulfate, urate
and organic acids.

54. Anion gap is defined as the difference between the total
concentration of measured cations (Na+ and K+) and that
of measured anion (Cl- and HCO3-).
The anion gap (A-) represents the unmeasured anions in
the plasma which may be calculated as follows, by
substituting the normal concentration of electrolytes
(mEq/l).
Na+ + K+ = Cl- + HCO3- + A-
136 + 4 = 100 + 25 + A-
The anion gap in a healthy individual is around 15 mEq/l
(range 8-18 mEq/l).
Acid-base disorders are often associated with alterations
in the anion gap.

55. Acid-base disorders- associated with
alterations in the anion gap.
Anion Gap Increase in Dehydration,Renal failure,
Hyperkalemia (due to increase in Na,K+, )
146+4=125+A-
150=125+A-
150-125= A-
25=A- (Wide)
Anion Gap Decrease in Diarrhoea (due to decrease in Na,K+)
130+3=125+A-
133=125+A-
133-125=A-
8=A- (Narrow)
Anion Gap Wide in Metabolic acidosis due to dec.HCO3-
Anion Gap Wide in Meningitis due to dec.Cl-

56. Regulation of electrolyte balance
The regulation is mostly achieved through the hormones .
1-Renin-Angiotensin Aldosterone System,
2-Anti Diuretic Hormone (ADH)
3-Atrial Natriuretaric Peptide (ANP)

57. Renin-Angiotensin Aldosteron System-RAAS
The secretion of aldosterone is controlled by renin-
angiotensin system.
Decrease in the blood pressure (due to a fall in ECF
volume) is sensed by juxtaglomerular apparatus of the
nephron which secrete renin.
Renin acts on angiotensinogen in blood to produce
angiotensin l .
The angiotensin l is then converted to angiotensin ll by
Angiotension Converting Enzyme(ACE) in lungs.
Angiotensin ll stimulates the release of aldosterone - a
mineralocorticoid produced by adrenal cortex (outermost
layer)..

58. Renin-Angiotensin Aldosteron System(RAAS)

59. Angiotension II Functions:
Angiotension II acts in 3 ways to raise the blood pressure
to normal range.
Aldosterone biosynthesis & secretion from Adrenal Gland:
Aldosterone acts on kidneys to increases Na+ reabsorption
by the renal tubules at the expense of K+ and H+ ions.
BLOOD VESSELS:
It causes vasoconstriction,increasing venous return from
periphery to heart,thus increasing stroke volume & B.P.
STIMULATION OF THIRST CENTRE:
Increasing fluid intake,increase blood volume & B.P.
The net effect is the retention of Na+ and water in the
body leading to increase in B.P. to normal.

60. Antidiuretic hormone (ADH)
An increase in the plasma osmolality (mostly due to Na+)
Stimulates hypothalamus to release ADH. ADH effectively
increases water reabsorption by renal tubules.

61. Atrial natriuretic peptide
This is a polypeptide hormone secreted by the right
atrium of the heart.
Atrial natriuretic peptide increases the urinary Na+
excretion and urine output (opp.to RAAS).
It is important to realise that Na+ and its anions (mainly Cl-
) are confined to the extracellular fluid. And the retention
of water in the ECF is directly related to the osmotic effect
of these ions Na+ and Cl-.

62. ACID BASE BALANCE
The normal pH of the blood is maintained in the narrow
range of 7.35--7.45, i.e. slightly alkaline.
Maintenance of blood pH is an important homeostatic
mechanism of the body.
Changes in blood pH will alter the intracellular pH which, in
turn,Influence the metabolism e.g. distortion in
protein structure, enzyme activity etc.
lt is estimated that the blood pH compatible to life is
6.8-7.8.

64. ACIDS IN THE BODY-VOLATILE ACIDS:
-
 Can leave solution and enter atmosphere e.g.CO2
 Produced by oxidative metabolism of CHO,Fat,Protein
 Average 20000 mEq of CO₂ per day
 Excreted through LUNGS as CO₂ gas

65. ACIDS IN THE BODY -FIXED ACIDS
 Acids that do not leave solution ,once
produced they remain in body fluids Until
eliminated by KIDNEYS
 Average 80 mEq/day
 Examples:
 Inorg. Acids:Sulfuric acid ,phosphoric acid
 Are generated during catabolism of:
 Amino acids(oxidation of sulfhydryl gps of cystine,methionine)
 Phospholipids(hydrolysis)
 Nucleic acids

66. ACIDS IN THE BODY-Organic acids
 Organic acids contain carbon produced during
metabolism:
 Pyruvic acid,lactic acid,citric acid,Glutamic
acid,Succinic acid,Oxalic acid,Aspartic acid etc.

67. Production of bases by the body
The formation of basic compounds in the body, in the
normal circumstances is negligible.
Some amount of bicarbonate is generated from the
organic acids such as lactate and citrate.
The ammonia produced in the amino acid metabolism is
converted to urea, hence its contribution as a base in the
body is insignificant.
A diet rich in animal proteins results in more acid
production by the body.
Vegetarian diet has an alkalizing effect on the body.

68. MAINTAINANCE OF BLOOD PH
The body has developed three lines of defense to regulate
the body's acid-base balance and maintain the blood pH
(around 7.4).
I. Blood buffers--1st line of defence
ll. Respiratory mechanism--2nd line of defence
lll. Renal mechanism--3rd line of defence

69. BUFFER
SYSTEM IN
BODY FLUIDS

70. Blood BUFFERS
 First line of defence (> 50 – 100 mEq/day)
 Two most common chemical buffer groups
 Bicarbonate Buffer system
 Phosphate Buffer system
 Protein Buffer system (Plasma proteins,Hemoglobin)

71. Limitations of Blood BUFFERS
Blood buffer systems act instantaneously and
regulate pH by binding or releasing H⁺
The buffer cannot remove H+ ions from the body.
It temporarily acts as a shock absorbant to reduce
the free H+ ions. The H+ ions have to be ultimately
eliminated by the renal mechanism

72. Sodium BICARBONATE and CARBONIC ACID BUFFER SYSTEM
Carbon Dioxide
 Most body cells constantly generate carbon dioxide
 Most carbon dioxide dissolve in water, is converted to carbonic acid,
which dissociates into H+ and a bicarbonate ion
Prevents changes in pH caused by organic acids and fixed acids in ECF
 Cannot protect ECF from changes in pH that result from elevated
or depressed levels of CO2
 Functions only when respiratory system and respiratory control
centers are working normally
 Ability to buffer acids is limited by availability of bicarbonate ions

73. ACID–BASE BALANCE
The CarbonicAcid–Bicarbonate Buffer System

74. THE SODIEM BICARBONATE-CARBONIC ACID-
BUFFER SYSTEM
• It is the most important buffer system.
• Carbonic acid( H2CO3) acts as the weak acid
-
• Sodium bicarbonate(Na HCO3) acts as the conjugate base
• Increase in H+(aq) ions is removed by HCO -(aq)
3
• The equilibrium shifts to the left and most of the H+(aq)
ions are removed.
• NaHCO3+HCl------- H2CO3+NaCl

75.  Any increase inOH-(aq) ions is removed by H2CO3
 H2CO3 dissociates, shifting the equilibrium to the
right, restoring most of the H+(aq) ions
 The small concentration of H+(aq) ions reacts with the
OH-(aq) ions producing H2O
 NaOH +H2CO3----- NaHCO3+H2O

76. The Ratio of HCO3 To H2CO3 And Alkali Reserve
At blood pH 7.4, the ratio of bicarbonate to carbonic acid is
20 : 1
Thus, the bicarbonate concentration is much higher
(20 times) than carbonic acid in the blood. This is referred
to as alkali reserve and is responsible for the effective
buffering of H+ ions, generated in the body.
In normal circumstances, the concentration of bicarbonate
and carbonic acid determines the pH of blood.
The bicarbonate buffer system serves as an index to
understand the disturbances in the acid-base balance of
the body.

77. PHOSPHATE BUFFER SYSTEM
Sodium dihydrogen phosphate and disodium hydrogen
phosphate( NaH2PO4- Na2HPO4) constitute the
phosphate buffer.
It is mostly an intracellular buffer .
It is of less importance in plasma due to its low
concentration with a pKa of 6.8 (close to blood pH 7.4), the
phosphate buffer would have been more effective, had it
been present in high concentration.
It is estimated that the ratio of base to acid for phosphate
buffer is 4:1 compared to 20:1 for bicarbonate buffer.

78. 29
PHOSPHATE BUFFER SYSTEM
 The phosphate buffer system (HPO4
2-/H2PO4
-)
plays a role in plasma and erythrocytes.
 H2PO4
- + H2O ↔ H3O+ + HPO4
2-
 Any acid reacts with monohydrogen phosphate
to form dihydrogen phosphate
dihydrogen phosphate monohydrogen phosphate
 H2PO4
- + H2O ← HPO4
2- + H3O+
 The base is neutralized by dihydrogen phosphate
dihydrogen phosphate monohydrogen phosphate
 H2PO4
- + OH- → HPO4
2- + H3O+

79. PROTEINS BUFFER SYSTEM
The plasma proteins and hemoglobin together constitute
the protein buffer system of the blood.
The buffering capacity of proteins is dependent on the pKa
of ionizable groups of amino acids. The imidazole
group of histidine (pK = 6.7) is the most effective
contributor of protein buffers.
The plasma proteins account for about 2o% of the total
buffering capacity of the plasma.

80. PROTEINS AS A BUFFER
 Proteins contain – COO- groups, which, like acetate ions
(CH3COO-), can act as proton acceptors.
 Proteins also contain – NH3
+ groups, which, like
ammonium ions (NH4
+), can donate protons.
 If acid comes into blood, hydronium ions can be
neutralized by the – COO- groups
- COO- + H3O+ → - COOH + H2O
 If base is added, it can be neutralized by the – NH3
+
groups
- NH3
+ + OH- → - NH2 + H2O
32

81. THE HEMOGLOBIN BUFFER SYSTEM
o From plasma CO2 diffuses into RBC’s
o It combine with H2O to form H2CO3
o As carbonic acid dissociates,Bicarbonate ions
diffuse into plasma,in exchange for chloride
ions (chloride shift/Hamberger phenomenon)
 Hydrogen ions are buffered by
hemoglobin molecules
 Is the only intracellular buffer system with an immediate
effect on ECF pH
Helps prevent major changes in pH when plasma PCO2
is rising or falling

82. RESPIRATORY ACID-BASE CONTROL
MECHANISMS
 When chemical buffers alone cannot prevent
changes in blood pH, the respiratory system is the
second line of defense against changes.
 Eliminate or Retain CO₂
 Change in pH are RAPID
 Occuring within minutes

83. RESPIRATORY MECHANISMS
Respiratory system provides a rapid mechanism for the
maintenance of acid-base balance. This is achieved by
regulating the concentration of carbonic acid (H2CO3) in
the blood i.e. the denominator in the bicarbonate
buffer system.
The large volumes of CO2 produced by the cellular
metabolic activity endanger the acid base equilibrium of
the body.
But in normal circumstances,all of this CO2 is eliminated
from the body in the expired air via the lungs, as
summarized below.
Carbonic anhydrase
H2CO3---------------------- CO2+ H2O.

84. RESPIRATORY MECHANISMS
The rate of respiration (or the rate of removal of CO2) is
controlled by a respiratory centre,located in the medulla of
the brain. This centre is highly sensitive to changes in the
pH of blood.
Any decreasein blood pH causes hyperventilation to blow
off CO2, thereby reducing the H2CO3 concentration.
Simultaneously the H+ ions are eliminated as H20.
Respiratory control of blood pH is rapid but only a short
term regulatory process/ since hyperventilation cannot
proceed for long.

85. Hemoglobin as a buffer
:Hemoglobin of erythrocytes is also important in the
respiratory regulation of pH.
At the tissue level, hemoglobin binds to H+ ions and helps
to transport CO2 as HCO3- with a minimum change in pH
(referred to as isohydric transport).
ln the lungs, as hemoglobin combines with O2, H+ ions are
removed which combine with HCO3 to form H2CO3.T he
latter dissociates to release H2O and CO2 to be exhaled
(Fig.).

86. Transport of CO2 By Hemoglobin

87. Generation of HCO3 by RBC
Due to lack of aerobic metabolic pathways, RBC produce
very little CO2. The plasma CO2 diffuses into the RBC
along the concentration gradient.
Here it combines with water to form H2CO3. This reaction
is catalysed by carbonic anhydrase.
In the RBC, H2CO3 dissociates to produce H+ and HCO3- .
The H+ ions are trapped and buffered by hemoglobin.

88. Chloride Shift /Hamberger Phenomenon
As the concentration of HCO3 increases in the RBC, it
diffuses into plasma along with the concentration
gradient, in exchange for Cl- ions to maintain electrical
neutrality.
This phenomenon, referred to as chloride shift
(Hamberger phenomenon) helps to generate HCO3 (Fig.).

89. Generation of bicarbonate by the erythrocyte

90. RENAL MECHANISMS
The renal mechanism highly significant, tries to provide a
permanent solution to the acid-base disturbances.
This is in contrast to the temporary buffering system and a
short term respiratory mechanism.
The kidneys regulate the blood pH by maintaining the
alkali reserve, besides excreting or reabsorbing the acidic
or basic substances, as the situation demands.

91. RENAL REGULATION OF ACID BASE BALANCE
 Role of kidneys is preservation of body’s bicarbonate
stores.
 Accomplished by:
 Reabsorption of 99.9% of filtered bicarbonate
 Regeneration of titrated bicarbonate by excretion of:
Titratable acidity (mainly phosphate)
Ammonium salts

92. Urine pH normally lower than blood pH
The pH of urine is normally acidic (-6.0). This clearly
indicates that the kidneys have contributed to the
acidification of urine, when it is formed from the blood
plasma (pH 7.4).
In other words, the H+ ions generated in the body
in the normal circumstances are eliminated by acidified
urine.
Hence the pH of urine is normally acidic (-6.0), while that
of blood is alkaline (7.4).
Urine pH, however, is variable and may range between 4.5-
9.5, depending on the concentration of H+ ions.

93. RENAL ACID-BASE CONTROL
MECHANISMS
 The kidneys are the third line of defence against
wide changes in body fluid pH.
 Excretion of H+ ions
 Reabsorbtion of bicarbonate
 Retention/Excretion of acids
 Excretion of Amonium ions
 Long term regulator of ACID – BASE balance
 May take hours to days for correction

96. 1. Excretion of H+ ions :
H+ excretion occurs in the proximal convoluted tubules
and is coupled with the regeneration of HCO3. Fig.
A- Carbonic anhydrase catalyses the production of
carbonic acid (H2CO3)f rom CO2 and H2O in the renal
tubular cell.
B- H2CO3 then dissociates to H+ and HCO3.
C- The H+ ions are secreted into the tubular lumen in
exchange for Na+.
D- The Na+ in association with HCO3 is reabsorbed into
the blood.
This is an effective mechanism to eliminate acids (H+)
from the body with a simultaneous generation of HCO3
which adds up to the alkali reserve.
E- The H+combines with a non-carbonate base(NH3,PO4)
and is excreted in urine.

97. Excretion of H+ ions

98. Renal Reabsorbtion of Bicarbonate

99. RENAL REABSORPTION OF
BICARBONATE
Proximal tubule:
70-90%
Loop of Henle:
10-20%
Distal tubule and
collecting ducts:
4-7%

100. FACTORS AFFECTING RENAL
BICARBONATE REABSORPTION
 Filtered load of
bicarbonate
 Prolonged changes in
pCO2
 Extracellular fluid
volume
 Plasma chloride
concentration
 Plasma potassium
concentration
 Hormones (e.g.,
mineralocorticoids,
glucocorticoids)

101. If secreted H+ ions combine with filtered
bicarbonate, bicarbonate is reabsorbed
If secreted H+ ions combine with
phosphate or ammonia, net acid excretion
and generation of new bicarbonate occur

102. Renal Excretion of Titerable Acid by PO4

103. TITRATABLE ACIDITY
 Occurs when
secreted H+
encounter & titrate
phosphate in tubular
fluid
 Refers to amount of
strong base needed
to titrate urine back to
pH 7.4
 40% (15-30 mEq) of
daily fixed acid load
 Relatively constant
(not highly adaptable)

104. Renal Excretion of Amonium Ions

105. AMMONIUM EXCRETION
 Occurs when
secreted H+
combine with NH3
and are trapped as
NH4
+ salts in
tubular fluid
 60% (25-50 mEq)
of daily fixed acid
load
 Very adaptable (via
glutaminase
induction)

106. NET ACID EXCRETION
 Hydrogen Ions
Are secreted into tubular fluid along
 Proximal convoluted tubule (PCT)
 Distal convoluted tubule (DCT)
 Collecting system

107. AMMONIUM EXCRETION
 Large amounts
of H+ can be
excreted
without
extremely low
urine pH
because pKa
of NH3/NH4
+
system is very
high (9.2)

108. CO2-Central Molecule for Blood pH Regulation

109. Buffers of intracellular fluids
The regulation of pH within the cells is as important as that
discussed above for the extracellular fluid. The H+ ions
generated in the cells are exchanged for Na+ and K+ ions.
This is particularly observed in skeletal muscle which
reduces the potential danger of H+ accumulation in the
cells.

110. ACID–BASE BALANCE DISTURBANCES
Interactions among the Carbonic Acid–Bicarbonate Buffer
System and Compensatory Mechanisms in the Regulation of
Plasma pH.

111. FOUR BASIC TYPES OF IMBALANCE
MetabolicAcidosis
MetabolicAlkalosis
Respiratory Acidosis
Respiratory Alkalosis

112. METABOLIC ACIDOSIS
 Production of strong acids in the body → ↓ pH
 (HCO₃⁻) used up to buffer excess H+ → ↓HCO₃⁻
 ↑ PCO2-Resp.centre stimulated-loss of CO2 → ↓ PCO2

113. CAUSES OF METABOLIC ACIDOSIS
3 ways: ↑Acid production,
 ↓H+ excretion,Loss of base
 LACTIC ACIDOSIS
Gluconeogenesis-Cori’s cycle
Myocardial Infarction,
Pumonary embolism,
 KETOACIDOSIS
 Diabetic
 Alcoholic
 Starvation
 RENAL FAILURE
(acute and chronic)
(H+ accumulates
 TOXINS
 Salicylates
 Diarrhoea,dehydration

114. COMPENSATION OF METABOLIC ACIDOSIS
 Buffers: Absorb H+
 Respiratory compensation: Hyperventilation
 Renal compensation:H+ excretion,HCO3 regeneration

115. ACID–BASE BALANCE DISTURBANCES
.
Responses to MetabolicAcidosis

116. METABOLIC ALKALOSIS
 ↑HCO₃⁻(Exogenous loads)
 ↓acid (Acid loss)
 ↑ pH due to ↑HCO₃⁻ or ↓acid

117. CAUSES OF METABOLIC ALKALOSIS
I. Exogenous HCO3 − loads
A. Acute alkali administration (Eno)
B. Milk-alkali syndrome
I. Gastrointestinal origin
1. Vomiting
2. Gastric content aspiration
III. Renal origin
1. Diuretics (loss of H+)
2. K+ depletion
3. (K+ and H+ compete for excretion by kidney,if K less,H+
is lost-resulting alkalosis)

118. COMPENSATION FOR METABOLIC ALKALOSIS
 Buffers: Donate H+
 Respiratory compensation: Hypoventilation
 Renal compensation:H+ generation,HCO3 secretion

119. ACID–BASE BALANCE DISTURBANCES
.
MetabolicAlkalosis

120. RESPIRATORY ACIDOSIS
 Retention of CO2 (increase H2CO3)
 ↑ PCO₂ → ↓pH

121. Causes of Respiratory Acidosis
Retention of CO2 (increase H2CO3) due to:
1-Any Obstrution in Resp. tract.
,airway obstruction F.B.,diseases of lungs,COPD etc.
2-Depression of resp. centre (drugs,morphine)

122. COMPENSATION IN RESPIRATORY ACIDOSIS
:
 The Renal mechanism comes into action to compensate resp.
acidosis by acidification of urine & bicarbonate retention.
 More HCO3- is generated and retained by kidneys which adds up
to the Alkali reserve of the body.
 The excretion of titerable acidity and NH4+ is elevated in urine.
 Acute respiratory acidosis is emergency.If airway is
patent,O2should be administered and artificial respiration
should be started,Drugs which stimulate the resp.
centre,should be administered.

123. ACID–BASE BALANCE DISTURBANCES
RespiratoryAcidosis-Regulation.

124. RESPIRATORY ALKALOSIS
Decrease in H2CO3 due to Prolonged hyperventilation
after increased exhalation of CO2 in:
1-Hypoxia:pneumonia,asthma,pulmonary oedema
2-Resp. centre stimulation:Hysteria,tension pain
3-Drugs: Slicylats poisoning etc.

125. Compensation
The renal mechanism tries to compensate by increasing
the urinary excretion of HCO3-.

126. ACID–BASE BALANCE DISTURBANCES
RespiratoryAlkalosis-Regulation.

127. ACID BASE DISORDERS
Disorder pH [H+] Primary
disturbance
Secondary
response
Metabolic acidosis    [HCO3
-]  pCO2
Metabolic alkalosis    [HCO3
-]  pCO2
Respiratory acidosis    pCO2  [HCO3
-]
Respiratory alkalosis    pCO2  [HCO3
-]

129. ANION GAP
The total concentrationo f cations and anions (expressed as
mEq/l) is equal in the body fluids to maintain electrical
neutrality.
The commonly measured electrolytes in the plasma are Na+,
K+, Cl- and HCO3-. Na+ and K+ together constitute about 95%
of the plasma cations. Cl- and HCO3- are the major anions,
contributing to about 80% of the plasma anions.
The remaining 20% of plasma anions (not normally measured
in the laboratory) include proteins, phosphate, sulfate, urate
and organic acids.

130. Anion gap is defined as the difference between the total
concentration of measured cations (Na+ and K+) and that
of measured anion (Cl- and HCO3-).
The anion gap (A-) represents the unmeasured anions in
the plasma which may be calculated as follows, by
substituting the normal concentration of electrolytes
(mEq/l).
Na+ + K+ = Cl- + HCO3- + A-
136 + 4 = 100 + 25 + A-
The anion gap in a healthy individual is around 15 mEq/l
(range 8-18 mEq/l).
Acid-base disorders are often associated with alterations
in the anion gap.

131. Acid-base disorders- associated with
alterations in the anion gap.
Anion Gap Increase in Dehydration,Renal failure,
Hyperkalemia (due to increase in Na,K+, )
146+4=125+A-
150=125+A-
150-125= A-
25=A- (Wide)
Anion Gap Decrease in Diarrhoea (due to decrease in Na,K+
130+3=125+A-
133=125+A-
133-125=A-
8=A- (Narrow)
Anion Gap Wide in Metabolic acidosis due to dec.HCO3-
Anion Gap Wide in Meningitis due to dec.Cl-

132. Types of Particles in Solution
1-True solution particles:
Size < 1nm,not visible with electron microscope,never
settle out of solution. Example:water,NaCl,sucrose
molecules
2- Colloidal particles:
Size 1-100 or upto 500 nm,can be seen with electron
microscope do not settle spontanously,but by salting out
methods.Example:plasma proteins,starch
3-Suspension particles:
Size bigger than colloidal,can be seen with naked eye,
make suspensions, settle down
spontaniously.Example:syrups

135. Types of Colloidal Solutions
Lyophobic Colloids or Suspensoids:
Liquid hating-no affinity b/w the particles of solute
(disperse phase) & the solvent(dispersion phase),
e.g.metals(Au,Pt,Ag) in water.these metals carry a charge
of one type only,so repel each other,pevent aggregation n
precipitation.This charge can be neutralized by adding a
dilute electrolyte solution,thus precipitating it,irreversibly.
Lyophilic Colloids or Emulsoids:
These also have charge,but surrounded by a tightly bound
layer of solvent as well.These two factors prevent
precipitation.Strong solutions of (NH4)2SO4 or NaCl.)
remove the charge n solvent thus precipitating
these.These can be resolubelized.Examples are proteins n
agar agar.

136. Separation of Colloidal Particles
1-Electrophoresis:Charged particles move towards
oppositly charged electrodes.
2-Ultracentrifugation:Particles are subjected to
gravitational force,light particles float,heavy sink down-
sedimentation.
3-Ultrafiltration:Filtration under pressure.Special filters
made of unglazed porcelain are used.
Dialysis:Used to remove waste products from the body
that accumulate as in Renal failure.
5-Precipitation with Electrolytes:Examples are separation
of Albumin n Globulin with Half n Full Saturation.
6-Adsorption:In this a layer of ions,molecules is condensed
upon a surface with which they come in contact.
e.g.activated charchoal,,silica gel,kaolin.

137. Protective Colloids
The combination of lyophilic particles with those of the
lyophobic ones causes the resulting particles to have
properties similar to lyophilic colloids.
1- Plasma proteins serve as protective colloids for calcium
phosphate aggregation in blood plasma held as colloidal
suspension.
2- Milk proteins serve as protective colloids for Calcium
phosphate aggregation in milk.
3- Bile salts keep insoluble cholestrol n calcium salts of
billirubin in colloid suspension, preventing stone formation
in gallbladder.

138. SURFACE TENSION
The surface of a liquid behaves as if it were a stretched
elastic membrane which resist external forces without
rupture and exerts a pressure inwards on the rest of the
liquid.Surface tension is the force or tension required to
break this film.
It is defined as the force in dynes acting upon a line one cm
long on the surface of the liquid.

139. Surface Tension Lowering Agents-Importance
Soaps,Detergents: help washing
Bile salts: Fascilitate Emulsification & digestion of Fats in
small intestine.
Lung surfactant (dipalmitoyal phophatydoil choline/
lecithine): prevent alveolar collapse during expiration.
Temperature Increase: decrease surface tension.
Surface Tension is increased by:Salts like NaCl: e.g.
nothing sinks in Dead sea due to saltish water.

140. Emulsions & Emulsifying Agents
- If an oil is shaken vigorously with water,it is broken into
smaller droplets to form an emulsion.This emulsion is
unstable,and droplets of oil colesce and then separate as a
layer of oil on water after a short time.
- If a little soap(e.g. Na stearate) is added, and again
shaken with water,a stable emulsion is formed.
-The ionized soap molecules contain two different
radicals.One is hydrocarbon,hydrophobic,which is soluble
in oil but not in water:the other is carboxyl
group,hydrophilic,soluble in water but not in oil.
-As a result soap ions are oriented at the interface of the
oil and water in such a way that HC are in oil & carboxyl
ions are in water.

141. Mechanism of Emulsification
1-The COO- group on the surface of the droplet give it a
negative charge which is balanced by the Na+ ions in the
surrounding water.Since all the droplets are charged
alike,they repel each other and remain in suspension
forming a stable emulsion.
2-2nd factor is that surface of the droplets becomes
attracted by water because the carboxylic groups are
strongly hydrophilic.
Some substances,detergents,phospholipids(DPPC),and
bile salts act as emulsifying agents by same mechanism.All
these decrease surface tension at the interface b/w oil and
water.

142. VISCOSITY
Fluidity is the characteristic property of liquids.Some
liquids flow more readily than others .
The reciprocal of fluidity is called viscosity,which means
the resistance offered by a liquid to flow.
This resistance depends on the friction of its component
molecules as they flow past one another.It is due to
attraction of the molecules from one to another as well as
asymmetry in their structure.
Thus water flows over a glass plate much faster than
honey.So water is said to have more fluidity but less
viscosity than honey.
The rate of flow of a liquid is inversely proportional to its
viscosity.