The document discusses the concepts of acid-base balance, explaining that acids donate protons and bases accept them. It details the strength of acids, the pH scale, and the importance of maintaining the pH of body fluids within a narrow range to avoid serious consequences. The mechanisms for regulation include chemical buffers, respiratory regulation, and renal regulation, which work to manage hydrogen ion concentrations and maintain metabolic function.

1. Acid-base balance
R.C.Gupta
Professor and Head
Dept. of Biochemistry
National Institute of Medical Sciences
Jaipur, India

2. Acids and bases
Acids and bases are defined
according to whether they donate or
accept hydrogen ions i.e. protons
An acid is a proton (H+) donor and
a base is a proton (H+) acceptor

3. H2CO3 H+ + HCO3
–
HNO3 H+ + NO3
–
H3PO4 H+ + H2PO4
–
H2PO4
– H+ + HPO4
– –
Acid Proton Base

4. HCl → H+ + Cl–
H2CO3 → H+ + HCO3
However, all acids do not dissociate to
the same extent
Strength of acids
Acids, e.g. hydrochloric acid and
carbonic acid, dissociate into their
component ions:

5. In a solution
containing HCl:
The acid is
almost completely
dissociated
The hydrogen ion
concentration is
very high
In a solution
containing H2CO3:
Majority of acid
molecules are
undissociated
The hydrogen ion
concentration is
low

6. Mineral acids:
Are highly
dissociated
Are strong acids
Organic acids:
Are poorly
dissociated
Are weak acids

7. Hydrogen ion concentrations in body
fluids are extremely low
Hydrogen ion concentration in blood is
about 0.00004 mmol/litre
The pH scale was devised by Sorensen
to express such low concentrations
pH scale

8. The pH scale can express the minute
hydrogen ion concentrations conveniently
pH of a solution is the negative log of the
hydrogen ion concentration in it
Hydrogen ion concentration is expressed
in mol/litre
pH = –log [H+]

9. Water dissociates into hydrogen ions and
hydroxyl ions
In pure water, number of hydrogen ions
is equal to the number of hydroxyl ions
Ion product of water (Kw) is:
Kw = [H+] [OH–] = 10-14

10. Hence, pH of water = – log [H+]
= – log 10‒7
= – (–7) = 7
Concentration of hydrogen ions as well
as hydroxyl ions in pure water is 10‒7
mol/litre

11. If the pH of a solution is
less than 7:
The hydrogen ion concentration
would be more than that of water
The solution would be acidic

12. If the pH of a solution is
more than 7:
The hydrogen ion concentration
would be less than that of water
The solution would be basic

13. If the pH of a solution
is exactly 7:
The hydrogen ion concentration
would be same as in water
The solution would be neutral

14. Normal pH of arterial blood is 7.35 -7.45
The average pH is 7.4
This means that the reaction is slightly
basic

15. The reaction can also be described
in terms of hydrogen ion concentration
The hydrogen ion concentrations in
body fluids are very low
1 mol/litre = 109 nanomol/litre
These are expressed in nanomol/litre

16. Increase in pH from 7.4 to 7.5,
means decrease in H+ concentration
by 9 nanomol/litre
pH H+ concentration
7.4 40 nanomol/litre
7.5 31 nanomol/litre
7.6 25 nanomol/litre
Increase in pH from 7.5 to 7.6,
means decrease in H+ concentration
by 6 nanomol/litre

17. The pH scale is not linear
It is preferable to describe changes in H+
concentration in nanomol/litre
However, due to prolonged usage, the pH
scale has come to stay

18. Acids and bases
Are constantly formed during
metabolic reactions
Can also enter from outside
May be lost in abnormal quantities
in pathological conditions
EMB-RCG

19. Acids may be:
Volatile e.g. carbonic acid
Non-volatile e.g. lactic acid
EMB-RCG

20. Daily production of H+ is:
15,000 mmol/day as volatile acids
75 mmol/day as non-volatile acids
EMB-RCG
If these are not removed, pH of body
fluids will be seriously disturbed

21. The pH of body fluids has to be
maintained within a narrow range
Departures from the normal range
can cause serious consequences
Conformation of proteins is extremely
sensitive to changes in pH
Regulation of pH of body fluids

22. Changes in conformation of enzymes
can impair the functioning of the
metabolic machinery
Therefore, mechanisms are required for
maintaining the pH of body fluids within
the normal range

23. Principal mechanisms for
regulation of pH are:
Chemical buffers
Respiratory regulation
Renal regulation
EMB-RCG

24. A chemical buffer
Can instantly neutralize acids and
bases
Is a system which resists a change in
pH on addition of an acid or a base
Is usually made up of a weak acid
and its alkali salt

25. The pH of a buffer can be calculated from
Henderson-Hasselbalch equation
Henderson-Hasselbalch equation is:
pH = pKa + log
[Salt]
[Acid]

26. In the equation, pKa is negative log of
dissociation constant (Ka) of the acid
component of buffer
Since pKa is constant, the pH depends
upon the ratio of the salt and the acid
component

27. Major buffers in body fluids
Bicarbonate-carbonic acid buffer
Phosphate buffer
Proteins
Haemoglobin
EMB-RCG

28. Made up of carbonic acid, a weak acid,
and its salt, bicarbonate
Quantitatively, the major buffer of extra-
cellular fluids especially plasma
Bicarbonate-carbonic acid buffer

29. Carbonic acid is formed from carbon
dioxide and water:
H2O + CO2 → H2CO3
Carbonic acid dissociates to form
bicarbonate:
H2CO3 → H+ + HCO3
-

30. pKa of carbonic acid (in equilibrium with
dissolved CO2) is 6.1
Average concentration of bicarbonate in
plasma is 24 mEq/L
Average concentration of carbonic acid in
plasma is 1.2 mEq/L

31. pH = pKa + log
or pH = 6.1 + log
or pH = 6.1 + log 20
or pH = 6.1 + 1.3 = 7.4
24
1.2
[HCO3
-]
[H2CO3]
Therefore, the pH of plasma would be:
As long as bicarbonate: carbonic acid
ratio is 20:1, the pH would remain 7.4

32. A buffer is most effective near its pKa
pKa of H2CO3 is rather distant from 7.4
Still bicarbonate-carbonic acid is an
important buffer in plasma because of its
high concentration

33. Since H2CO3 is a much weaker acid than
HCl, the change in pH would be minimal
HCl + NaHCO3  H2CO3 + NaCl
The salt component of the buffer can
convert strong acids into weak acids:

34. The acid component of the buffer can
convert strong bases into weak bases:
NaOH + H2CO3  NaHCO3 + H2O
As NaHCO3 is a much weaker base than
NaOH, the change in pH would be minimal
Thus, the buffer resists a change in pH on
addition of acids as well as bases

35. Measuring pCO2 is easier than measuring
H2CO3
Concentration of H2CO3 can be calculated
by multiplying pCO2 by a constant

36. The constant depends upon the solvent
and the temperature
For plasma at 37°C, the constant is
0.0301 or approximately 0.03
In the equation for calculating pH, [H2CO3]
can be replaced by pCO2 x 0.03

37. Phosphate buffer
Phosphate buffer is formed from inorganic
phosphate
Phosphate ions are present in two forms:
Dihydrogen phosphate (H2PO4
–)
Monohydrogen phosphate (HPO4– 2)

38. H2PO4
– is a weak acid as it can donate a
proton
HPO4
– 2 is a base as it can accept a
proton
In ECF, these exist as NaH2PO4 and
Na2HPO4, and constitute a buffer

39. Na2HPO4 can neutralize acids:
HCl + Na2HPO4  NaCl + NaH2PO4
Thus a strong acid is converted into
a weak acid

40. NaOH + NaH2PO4  H2O + Na2HPO4
Thus, the change in pH on addition of an
acid or a base is minimal
In this reaction, a strong base is
converted into a weak base
NaH2PO4 can neutralize bases:

41. The pH of a fluid containing phosphate
buffer depends upon:
Ratio of HPO4
– 2 to H2PO4
– which
is 4:1 in plasma
pKa of H2PO4
– which is 6.8

42. In the presence of phosphate buffer, the
pH will be:
pH = pKa + log
or pH = 6.8 + log 4
or pH = 6.8 + 0.6 = 7.4
[HPO4
– 2 ]
[H2PO4
–]

43. Concentration of inorganic phosphate in
extra-cellular fluids is low
Yet phosphate buffer is an effective buffer
as pKa of H2PO4
– is close to 7.4

44. Proteins act as buffers because of their
amphoteric nature
In acidic medium, they act as bases and
neutralize acids
In basic medium, they act as acids and
neutralize bases
Proteins

45. The amino acid residues having pKa close
to 7.4 are the most effective in buffering
Among different amino acids, pKa of
histidine is the closest to 7.4

46. Intracellular fluid (ICF) and plasma have
sizeable concentration of proteins
But other extracellular fluids have a low
protein content
Hence, the buffering action of proteins is
exerted mainly in ICF and plasma

47. Haemoglobin (Hb) also acts as a buffer
while transporting O2 and CO2
CO2 is produced continuously in various
metabolic reactions
Hb buffers the large amount of carbonic
acid which is formed from carbon dioxide
Haemoglobin

48. Carbonic acid is present in large amounts
in RBCs
It dissociates into H+ and HCO3
‒
Hb takes up the hydrogen ions and
prevents a change in pH
Haemoglobin is responsible for 60% of
the buffering capacity of blood

49. EMB-RCG

50. Carbonic acid is the major end product of
metabolism in the form of carbon dioxide
Respiratory mechanism regulates the
elimination of carbonic acid
Respiratory regulation

51. The purpose of regulation is to maintain
the ratio of bicarbonate to carbonic acid
Respiratory buffering occurs in minutes to
hours

52. • pH of blood
• pCO2 of blood
• pO2 of the blood
Respiratory
centre in
the medulla
is sensitive
to changes
in:
EMB-RCG

53. Respiratory centre, accordingly, regulates the
rate and depth of respiration
They transmit information to the respiratory
centre
They perceive changes in pH, pCO2 and pO2
Chemoreceptors are located in the aortic arch
and carotid sinus
EMB-RCG

54. A change in pH is the most important
stimulant of respiratory centre
A decrease in pH stimulates the
respiratory centre
This leads to hyperventilation and
increased elimination of CO2
Decreased carbonic acid concentration
raises the pH

55. The respiratory centre is also stimulated
by a rise in pCO2 and marked anoxaemia
But their effect is less than that of a
decrease in pH

56. The respiratory mechanism tries to
maintain the bicarbonate: carbonic acid
ratio in blood
If bicarbonate concentration changes, the
respiratory mechanism alters carbonic
acid concentration accordingly

57. During normal metabolism, the body
produces a large amount of acids
On an average diet, about 75 mEq of non-
volatile acids are produced every day
These include sulphate, phosphate and
organic acids
Renal regulation

58. If the acids are not excreted, the pH of
blood will become acidic
Kidneys prevent a change in pH by:
Excreting hydrogen ions in urine
Returning bicarbonate to blood

59. The pH of urine is usually acidic due to
renal secretion of H+
Renal buffering takes hours to days

60. The renal mechanism excretes the
excess acids by:
Reabsorption of
bicarbonate
Acidification of
monohydrogen phosphate
Secretion of ammonia
EMB-RCG

61. More than 4,000 mEq of bicarbonate is
filtered by glomeruli everyday
If it is lost in urine, it will be a major drain
on alkali reserve
This will deplete the main chemical
buffer of plasma
Reabsorption of bicarbonate

62. Tubular reabsorption of bicarbonate
prevents this loss
All the bicarbonate filtered in glomeruli is
reabsorbed in proximal convoluted tubules
This is also known as tubular reclamation
of bicarbonate

63. Carbonic anhydrase present in tubular
cells converts H2O and CO2 into H2CO3
Carbonic acid dissociates into a hydrogen
ion and a bicarbonate ion
The hydrogen ion is secreted into the
tubular lumen

64. The H+ combines with HCO3
‒ in the
lumen to form H2CO3
The sodium ion freed from bicarbonate
enters the tubular cell
Sodium-hydrogen exchanger facilitates the
trans-membrane movement of Na+ and H+

65. The H+ is pumped into capillaries by
Na+, K+-ATPase
A bicarbonate ion accompanies the
exiting sodium ion

66. CO2
H2O
H
+
H+
CO2
CA
H2O H2CO3 NaHCO3
H2CO3
Blood Tubular fluid
Na
+
Na+
Na
+
Proximal
convoluted
tubule
HCO3
‒
HCO3
‒

67. After reabsorption of all the HCO3
‒, H+
secretion proceeds against Na2HPO4
This occurs in the distal convoluted
tubules
Hydrogen ions secreted by the cells react
with Na2HPO4 in the lumen
Acidification of monohydrogen phosphate

68. Na2HPO4 is converted into NaH2PO4
Sodium ion is reabsorbed in exchange
for hydrogen ion
Sodium and bicarbonate ions are
returned to blood

69. HCO
–
3
H
+
H+
CO2
CA
H2O H2CO3 Na2HPO4
NaH2PO4
Blood Tubular fluid
Na
+
Na+
Na
+
Distal
convoluted
tubule
HCO3
‒

70. Conversion of Na2HPO4 into NaH2PO4
causes acidification of urine
The acidity due to NaH2PO4 is known as
titratable acidity
Titratable acidity is measured by titrating
urine with NaOH to a pH of 7.4

71. Reabsorption of sodium also occurs
against ammonium ions in distal
convoluted tubules
Ammonia is formed by deamination of
amino acids, particularly glutamine, in
tubular cells
Secretion of ammonia

72. Ammonia diffuses into tubular lumen
H+ secreted by tubular cell combines
ammonia to form NH4
+
NH4
+ reacts with NaCl forming NH4Cl
Sodium ion released from NaCl is
reabsorbed

73. H+
H2O
Gluta-
minase
Glutamate
Glutamine NH3
NaCl
NH4Cl
Blood
Distal convoluted
tubule Tubular fluid
Na+
CO2
H2CO3
NH+
4
CA
NH3
H+
Na+Na+
HCO3
‒HCO3
‒

74. Renal regulation can respond to changes
in the acid-base balance of blood
If production of acids increases, kidneys
cause more acidification of urine
Any deficiency in chemical and respiratory
buffering is corrected by the kidneys
Renal regulation is slow but is very
thorough

75. Disorders of acid-base
balance occur:
When regulatory mechanisms
fail to maintain the pH
When the bicarbonate: carbonic
acid ratio deviates from 20:1
EMB-RCG

76. Acidosis and alkalosis:
A decrease in pH below
normal is known as acidosis
An increase in pH above
normal is known as alkalosis
EMB-RCG

77. Respiratory acidosis or
respiratory alkalosis
An increase or decrease in
carbonic acid causes:
EMB-RCG

78. Metabolic acidosis or
metabolic alkalosis
A decrease or increase in
bicarbonate causes:
EMB-RCG

79. Renal mechanism tries to compensate
respiratory acidosis or alkalosis
Respiratory mechanism tries to
compensate metabolic acidosis or alkalosis
EMB-RCG

80. Compensation
Is more effective in
chronic disorders
But is never 100%
EMB-RCG

81. Respiratory acidosis
This results from accumulation of carbon
dioxide (and carbonic acid)
Inspiring air having high carbon
dioxide content
Hypoventilation resulting in
decreased elimination of CO2 or
Accumulation can occur due to:

82. Acute respiratory acidosis can occur
due to:
• Collapse of lungs
• Pneumothorax
• Haemothorax
• Head injury depressing respiratory
centre
• Overdose of general anaesthetics,
opiates, alcohol or sedatives that
depress respiratory centre

83. Chronic respiratory acidosis can occur
due to:
• Bronchial asthma
• Emphysema
• Bronchiectasis
• Chronic bronchitis
• Myopathies
• Myasthenia
• Intracranial tumours

84. When respiratory acidosis begins:
pH is decreased
Blood bicarbonate is normal
pCO2 is elevated

85. Kidneys compensate respiratory
acidosis by:
Returning more bicarbonate to
blood
Excreting more hydrogen ions
in urine

86. In acute cases, compensatory increase in
bicarbonate is 1 mmol/L for every 10
mm of Hg rise in pCO2
In chronic cases, compensatory increase
in bicarbonate is 4 mmol/L for every
10 mm of Hg rise in pCO2

87. Least common acid-base disorder
Results from a decrease in CO2 (and
carbonic acid) content of blood
Decrease in CO2 is due to hyperventilation
Respiratory alkalosis

88. Acute respiratory alkalosis can occur
due to:
• Hysterical hyperventilation
• Encephalitis
• Meningitis
• Cerebrovascular accident
• Pneumonia
• Salicylate poisoning (early stage)

89. Chronic respiratory alkalosis can occur
due to:
• Severe anaemia
• Cardiac failure
• Heat exposure
• Overuse of mechanical ventilators

90. When respiratory alkalosis begins:
pH is increased
Blood bicarbonate is normal
pCO2 is decreased

91. Kidneys compensate respiratory
alkalosis by:
Returning less bicarbonate to
blood
Excreting less hydrogen ions
in urine

92. In acute cases, compensatory decrease
in bicarbonate is 2 mmol/L for every 10
mm of Hg decrease in pCO2
In chronic cases, compensatory decrease
in bicarbonate is 4 mmol/L for every 10
mm of Hg decrease in pCO2

93. Metabolic acidosis
Commonest disorder of acid-
base balance; can be due to:
Increased production of
endogenous acids
Decreased excretion of
endogenous acids
Entry of exogenous acids
Loss of bases

94. Patients with metabolic acidosis can be
divided into two groups on the basis
of anion gap
Commonly measured anions
(Cl- and HCO3
-)
Commonly measured cations
(Na+ and K+)
Anion gap is the difference between the
plasma concentrations of:

95. Normally, the sum of sodium and
potassium exceeds the sum of chloride
and bicarbonate by about 15 mEq/L
Anion gap = [Na+ + K+ ] – [Cl– + HCO3
– ]

96. The anion gap represents the concen-
tration of unmeasured anions in plasma
The anion gap are pyruvate, phosphate,
sulphate, anionic proteins etc

97. In these patients, plasma bicarbonate is
low but the anion gap is normal due to a
reciprocal increase in chloride
Hence, this condition is also known as
hyperchloraemic metabolic acidosis
Metabolic acidosis with normal
anion gap

98. The causes of metabolic acidosis with
normal anion gap are:
• Diarrhoea
• Gastrointestinal fistula
• Intestinal obstruction
• Renal tubular acidosis
• Administration of ammonium chloride
• Carbonic anhydrase inhibitors

99. Blood bicarbonate is decreased
Chloride is increased
pCO2 is normal
pH is decreased
When the disorder begins:

100. Respiratory compensation occurs by way
of hyperventilation
Compensatory decrease in pCO2 is 1.25
mm of Hg for every 1 mmol/L decrease
in bicarbonate

101. These patients have low blood
bicarbonate and normal chloride
Anion gap is increased due to the
presence of some abnormal and
unmeasured anions
Metabolic acidosis with increased
anion gap

102. Causes of metabolic acidosis with
increased anion gap include:
• Diabetic ketoacidosis
• Ketoacidosis of starvation
• Alcoholic ketoacidosis (sudden
withdrawal)
• Uraemia
• Lactic acidosis
• Salicylate intoxication (in later stages)
• Intoxication with formic acid, oxalic acid,
ethylene glycol, paraldehyde, methanol
etc

103. When the disorder begins:
The respiratory mechanism compensates
the acidosis
pH is decreased
pCO2 is normal
Chloride is normal
Blood bicarbonate is decreased

104. The rate and depth of respiration is
increased
As a result, pCO2 decreases
Bicarbonate: carbonic acid ratio returns
towards normal

105. The compensatory decrease in pCO2 is
1.25 mm of Hg for every 1 mmol/L
decrease in bicarbonate

106. Metabolic alkalosis
Can occur from loss of acids or excess
of bases; common causes are:
Potassium deficit
Excessive use of antacids
Loss of HCl due to severe vomiting
or prolonged gastric aspiration

107. Blood bicarbonate is high
Chloride is reciprocally low
pCO2 is normal
pH is increased
When the disorder begins:

108. Respiratory mechanism compensates
metabolic alkalosis by decreasing:
The depth of respiration
The rate of respiration

109. More carbon dioxide is retained
Bicarbonate: carbonic acid ratio is brought
towards normal
Compensatory increase in pCO2 is 0.75
mm of Hg for every 1 mmol/L increase in
bicarbonate

110. Disorder Blood
pH
Primary
change
Compensatory
change
Respiratory
acidosis
  pCO2  HCO3
–
Respiratory
alkalosis
  pCO2  HCO3
–
Metabolic
acidosis
 HCO3
– pCO2
Metabolic
alkalosis
  HCO3
–  pCO2
Blood chemistry in acid-base disorders

111. Mixed acid-base disorders
Some patients may have two or more
diseases affecting acid-base balance
These can produce independent changes
in acid-base balance

112. A diabetic with renal complications or an
independent renal disease may develop:
The two together may result in severe
acidosis
Metabolic acidosis due to renal disease
Metabolic acidosis due to ketosis

113. A diabetic with chronic obstructive
pulmonary disease may develop:
The two together may result in severe
acidosis as compensation would not occur
Respiratory acidosis due to lung disease
Metabolic acidosis due to diabetes