Mechanisms for regulation of pH and disorders of acid-base balance

1. Acid-base Balance
Regulation of pH of body fluids
R.C. Gupta
M.D. (Biochemistry)
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; a base
is a proton (H+) acceptor

3. H2CO3 H+ + HCO3
–
HNO3 H+ + NO3
–
H3PO4 H+ + H2PO4
–
H2PO4
– H+ + HPO4
–2
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. Organic
acids:
Are poorly
dissociated
Are weak
acids
Mineral
acids:
Are highly
dissociated
Are strong
acids

7. Hydrogen ion concentrations in the 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. In pure water, number of hydrogen ions
is equal to the number of hydroxyl ions
Water dissociates into hydrogen ions
and 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. 1 mol/litre = 109 nanomol/litre
These are expressed in nanomol/litre
The hydrogen ion concentrations in
body fluids are very low
The reaction can also be described in
terms of hydrogen ion concentration

16. If pH goes from 7.4 to 7.5, H+ concen-
tration decreases by 9 nanomol/litre
pH H+ concentration
7.4 40 nanomol/litre
7.5 31 nanomol/litre
7.6 25 nanomol/litre
If pH goes from 7.5 to 7.6, H+ concen-
tration decreases 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 is still in common use

18. Acids and bases:
Are formed continuously during
metabolic reactions
Can also enter from outside
May be lost in abnormal quantities
in pathological conditions
Acid-base balance

19. Acids formed in
the body may be:
Volatile e.g.
carbonic acid
Non-volatile e.g.
lactic acid

20. Daily production of H+ is:
15,000 mmol/day as volatile acids
75 mmol/day as non-volatile acids
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 the conformation of enzymes
can impair the functioning of 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 buffering
Respiratory regulation
Renal regulation

24. Chemical buffers are the first line of defense
against a threat to acid-base balance
They act within seconds to minutes
Chemical buffering

25. 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

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

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

28. Major buffers in body fluids:
Bicarbonate-carbonic acid buffer
Phosphate buffer
Proteins
Haemoglobin

29. 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

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

31. pKa of carbonic acid is 6.1
Average concentration of bicarbonate in
plasma is 24 mEq/L
Average concentration of carbonic acid in
plasma is1.2 mEq/L

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

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

34. 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:

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

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

37. 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

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

39. 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

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

41. NaOH + NaH2PO4  H2O + Na2HPO4
Thus, the buffer minimizes the change in
pH on addition of an acid or a base
This reaction converts a strong base
into a weak base
NaH2PO4 can neutralize bases:

42. 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

43. 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
–]

44. 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

45. 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

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

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

48. 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

49. 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

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

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

53. Respiratory centre in the medulla is
sensitive to changes in:
pH of
blood
pCO2 of
blood
pO2 of
blood

54. Respiratory centre regulates the rate and
depth of respiration accordingly
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

55. 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

56. 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

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

58. Large amounts of volatile and non-volatile
acids are produced in the body every day
Volatile acids are eliminated by respiratory
mechanism
Renal mechanism takes care of the non-
volatile acids
Renal regulation

59. On an average diet, about 75 mEq of non-
volatile acids are produced every day
These include organic acids, phosphate,
sulphate etc
If these acids are not eliminated, the pH of
blood will become acidic

60. Kidneys prevent a change
in the pH of blood by:
Excreting hydrogen ions in urine
Returning bicarbonate to blood

61. While pH of blood (and glomerular filtrate)
is basic, the pH of urine is usually acidic
This is due to renal tubular secretion of H+
Secretion of H+ acidifies the urine

62. The renal mechanism
excretes the acids by:
Reabsorption of
bicarbonate
Acidification of
monohydrogen phosphate
Secretion of ammonia

63. 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

64. Loss of bicarbonate is prevented by its
tubular reabsorption
All the bicarbonate filtered in glomeruli is
reabsorbed in proximal convoluted tubules
This is also known as tubular reclamation
of bicarbonate

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

66. The H+ reacts with NaHCO3 in the lumen to
form H2CO3 and Na+
The Na+ freed from NaHCO3 enters the
tubular cell
Sodium-hydrogen exchanger facilitates the
trans-membrane movement of Na+ and H+

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

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

69. After all the bicarbonate is reabsorbed, H+
secretion proceeds against Na2HPO4
This occurs in distal convoluted tubules
Hydrogen ions secreted by the cells react
with Na2HPO4 in the lumen
Acidification of monohydrogen phosphate

70. Na2HPO4 is converted into NaH2PO4 and
Na+
Na+ released from Na2HPO4 enters the
tubular cell in exchange for H+
The sodium ion and a bicarbonate ion are
released into blood

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

72. 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

73. In distal convoluted tubules, sodium is
reabsorbed against ammonium ions also
Ammonia is formed by deamination of
amino acids in tubular cells
A major source of ammonia is glutamine
Secretion of ammonia

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

75. 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
‒

76. Renal regulation responds to changes in
the pH of blood
If the pH decreases, kidneys increase the
acidification of urine
If the pH increases, kidneys decrease the
acidification of urine

77. Renal buffering occurs over hours to days
Renal regulation is slow but is very
thorough
It can correct any deficiency in chemical
and respiratory buffering

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

79. Disorders of acid-base balance are
acidosis and alkalosis
Acidosis is a decrease in the pH of
blood below the normal range
Alkalosis is an increase in the pH of
blood above the normal range

80. Acidosis
can be:
Respiratory
acidosis
Metabolic
acidosis
Alkalosis
can be:
Respiratory
alkalosis
Metabolic
alkalosis

81. Respiratory acidosis or
respiratory alkalosis
An increase or decrease in
carbonic acid causes:

82. Metabolic acidosis or
metabolic alkalosis
A decrease or increase in
bicarbonate causes:

83. Renal mechanism tries to compensate
respiratory acidosis or alkalosis
Respiratory mechanism tries to comp-
ensate metabolic acidosis or alkalosis

84. Compensation
Is more effective in
chronic disorders
But is never 100%

85. Respiratory acidosis
There is accumulation of carbon dioxide
Carbon dioxide forms carbonic acid
Ratio of bicarbonate to carbonic acid is
decreased

86. Inspiring air having high carbon
dioxide content
Hypoventilation resulting in
decreased elimination of CO2 or
Accumulation of CO2
can occur due to:

87. Acute respiratory acidosis occurs
over a short period of time
Respiratory acidosis can be acute or
chronic
Chronic respiratory acidosis occurs
over a long period of time

88. 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

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

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

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

92. 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

93. This is the least common disorder of acid-
base balance
There is a decrease in CO2 (and hence
carbonic acid) content of blood
Decrease in CO2 is due to hyperventilation
Respiratory alkalosis

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

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

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

97. Kidneys compensate respiratory
alkalosis by:
Returning less bicarbonate to
blood
Decreasing the excretion of
hydrogen ions in urine

98. 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

99. 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
Abnormal loss of bases

100. Patients with metabolic acidosis can be
divided into two groups on the basis of
anion gap
Patients with
normal anion
gap
Patients with
increased
anion gap

101. Commonly measured anions
(Cl- and HCO3
-)
Commonly measured cations
(Na+ and K+)
Anion gap is the difference between the
plasma concentrations of:

102. Normally, the sum of sodium and
potassium exceeds the sum of chloride
and bicarbonate by about 8-15 mEq/L
Anion gap = {[Na+] + [K+]} – {[Cl–] + [HCO3
–]}

103. Anion gap represents the concentration of
unmeasured anions in plasma
The unmeasured anions are pyruvate,
sulphate, phosphate, anionic proteins etc

104. In these patients, plasma bicarbonate is
decreased 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

105. 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

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

107. Metabolic acidosis is compensated by the
respiratory mechanism
Respiratory compensation occurs by way
of hyperventilation (to decrease pCO2)
The decrease in pCO2 is 1.25 mm of Hg
for every 1 mmol/L decrease in HCO3
–

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

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

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

111. The respiratory mechanism compensates
the acidosis
Rate and depth of respiration increase
Compensation occurs by way of
hyperventilation

112. Compensatory decrease in pCO2 is 1.25
mm of Hg for every 1 mmol/L decrease
in bicarbonate
pCO2 decreases due to hyperventilation
Bicarbonate: carbonic acid ratio returns
towards normal

113. 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

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

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

116. pCO2 increases due to hypoventilation
Increase in pCO2 is 0.75 mm of Hg for
every 1 mmol/L increase in bicarbonate
Bicarbonate: carbonic acid ratio is brought
towards normal

117. 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

118. 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

119. Uncontrolled diabetes mellitus can result
in ketoacidosis
If the patient has severe vomiting also, it
can cause alkalosis

120. A patient with chronic obstructive pulmonary
disease may develop respiratory acidosis
Severe vomiting in such a patient may cause
metabolic alkalosis
Compensation may be inadequate or
excessive in mixed acid-base disorders