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Acids and Bases – their definitions and meanings
Molecules containing hydrogen atoms that can release hydrogen ions
in solutions are referred to as acids. (HCl – H+ Cl-) (H2CO3
H+ HCO3-)
A base is an ion or a molecule that can accept a hydrogen ion.
(HPO42- is base because it can accept hydrogen ion to form H2PO4-)
The proteins in the body also as bases because some of the amino acids that make up
proteins have negative charges that readily accept hydrogen ions.
Alkalosis refers to excess removal of hydrogen ions from the body fluids.
Acidosis refers to the excess addition of hydrogen ions in the body fluids.
A strong acid is one that rapidly dissociates and releases large amounts of H+ in solution (HCl)
A week acid have less tendency to dissociate its ions and, therefore release H + (H2CO3)
Control of Acid-Base Balance
1.
2.
3.
4.
There must be a balance between the production of H+ and the net
removal of H+ from the body.
Precise H+ regulation is essential because the activities of almost
all enzyme systems in the body are influenced by H+ concentration.
Na+ = 142 mmol/L, H+ = 0.00004 mmol/L (40 nmol/L)
pH = -log [H+] = -log[0.00004] = 7.4
(The lower limit of pH at which a person can live more than a few hours is
about 6.8 and the upper limit is about 8.0)
5.
There are three primary systems that regulate the H+ concentration
in body fluids to prevent acidosis:
A/ Chemical acid-base buffer systems of the body fluids (seconds)
B/ Lungs (few minutes)
C/ Kidneys (hours to days)
Metabolic Sources of Acids and Bases
A. Reactions producing CO2 (Merely a Potential Acid)
1. Complete oxidation of neutral carbohydrated and fat
2. Oxidation of most neutral amino acids
CO2 + H2O
Urea + CO2 + H2O
B. Reactions producing nonvolatile acids
1. Oxidation of sulfur-containing amino acids
Urea + CO2 + H2O + H2SO4
(examples: methionine, cysteine)
2. Metabolism of phosphorous-containing compounds
H3PO4
3. Oxidation of cationic amino acids
Urea + CO2 + H2O + H+
(examples: lysine+, arginin+)
4. Production of nonmetabolizable organic acids
HA
H+ + A(examples: uric acid, oxalic acid)
5. Incomplete oxidation of carohydrate and fat
HA
H+ + A(examples: lactic acid, ketoacidosis)
C. Reactions producing nonvaletile bases
1. Oxidation of anionic amino acids
(examples: glutamate-, aspartate-)
2. Oxidation of organic anions
(examples: lactate-, acetate-)
Urea + CO2 + H2O + HCO3CO2 + H2O + HCO3-
2H+ + SO42H+ + H2PO42-
Buffering of Hydrogen Ions in the Body Fluids
Daily production of H+ = 80 mmol,
Body fluid concentration = 0.00004 mmol/L
Buffer + H+
H Buffer
In this example, a free H+ combines with the buffer to form a weak acid (H Buffer)
Bicarbonate Buffer System
H+ + HCO3-
H2CO3
CO2 + H2O
From these reactions, one can see that the hydrogen ions from the strong acid react with
HCO3- to form the very weak acid (H2CO3), which in turn forms CO2 and H2O.
The excess of CO2 stimulates respiration
CO2 + H2O
H2CO3
+
NaOH
HCO3+
Na
+ H+
The weak base NaHCO3- replaces the strong base NaOH. At the same time the concentration
of H2CO3 decreases (because it reacts with NaOH), causing more CO2 to combine with H2O,
in order to replace the H2CO3.
The net result is a tendency for the CO2 levels in the blood to decrease, but it is prevented
by the decreased ventilation.
The rise in blood HCO3- is compensated by increased renal excretion of HCO3-.
Phosphate Buffer System
It plays a major role in buffering renal tubular fluid and intracellular fluid
HPO42- + H+
H2PO4-
Proteins: Important Intracellular Buffer
H+ + Hemoblobin
HHemoglobin
Approximately 60 to 70 percent of the total chemical buffering
of the body fluids is inside the cells, and most of this results from
the intracellular proteins.
H+
H2CO3
+
HCO-3
For any acid, the concentration of the acid relative to its dissociated ions is defined by the
dissociation constant K´.
H+ x HCO3K´
=
H2CO3
This equation indicates that in an H2CO3 solution, the amount of free H+ is equal to:
H2CO3
H+
=
K´
x
HCO3-
The concentration of undissociated H2CO3 cannot be measured in solution because is rapidly
dissociated. However, the amount of CO2 dissolved in the blood is directly proportional
to the amount of undissociated H2CO3.
CO2
H+
=
K
x
HCO3-
The dissociation constant (K) is only about 1/400 of K´ because the proportionality ratio between
H2CO3 and CO2 is 1:400. In addition, we don´t measure the total amount of CO2 in blood. Fortunately,
the amount of CO2 in the blood is linear function of pCO2 multiplied by the solubility coefficient.
The solubility coefficient for CO2 is 0.03 mmol/mmHg. This means that 0.03 millimole of H2CO3 is present
in the blood for each mmHg pCO2 measured
(0.03 x pCO2)
H+
=
Kx
HCO3-
It is customary to express H+ concentration in pH units rather than in actual concentration.
pH = -log H+. The dissociation constant can be expressed in a simila manner, pK = -log K.
-log H+ = -log pK – log
(0.03 x pCO2)
HCO3(0.03 x pCO2)
pH
= pK
- log
HCO3-
Rather than work with a negative logarithm, we can change the sign of the logarithm and invert the
numerator and denominator.
HCO3pH = pK +
log
0.03 x pCO2
Isohydric Principle:
all buffers in a common solution are in equilibrium with same H+ concentration
HA1
H+
= K1 x
HA2
= K2 x
A1
HA3
= K3 x
A2
A3
All buffer systems work together, because H+ is common to the reactions of all these systems.
Any condition that changes the balance of one of the buffer system also changes the balance of al the
others.
Henderson-Hasselbalch Equation:
HCO3pH = 6.1 + log
0.03 x pCO2
Kidney
pH = Constant +
Lung
Henderson-Hasselbalch Equation:
HCO3pH = 6.1 + log
0.03 x pCO2
1. Increase in bicarbonate ion concentration causes the pH to rise.
2. Increase in pCO2 causes the pH to decrease.
1. Bicarbonate concentration is regulated mainly by the kidneys.
2. pCO2 concentration is regulated by the rate of respiration.
1. When disturbances of acid-base balance results from a primary changes
in extracellular fluid bicarbonate concentrations are referred to as
metabolic acid-base disorders.
2. When disturbances of acid base balance results from a primary changes in pCO2
are referred as respiratory acid-base disorders.
Respiratory regulation of acid-base balance
Respiratory regulation of acid-base balance is a physiological type of buffer system, because
it acts rapidly and keeps the hydrogen ion concentration from changing too much until
the much slowly responding kidneys can eliminate the imbalance.
H+
Alveolar ventilation
pCO2
If hydrogen ion concentration is suddenly increased by adding acid to extracellular fluid
and pH falls from 7.4 to 7.0, the respiratory systém can retrun the pH to a value of 7.2 to 7.3
within 3 to 12 minutes.
Respiratory system has an efectivness between 50 to 75 per cent.
The kidneys regulate extracellular fluid H+ concentrations thought
three fundamental mechanisms:
1. Reabsorption of filtered HCO32. Secretion of H+
3. Production of new HCO3-
Ad. 1.
180 L/day x 24 mmol/L = 4320 mmol of HCO3-
Proximal tubule, thick ascending loop of Henle, early distal tubule
Thus, each time a hydrogen ion is formed in the tubular epithelial cells,
a bicarbonate ion is also formed and released back into the blood.
The net effect of these reactions is a „reabsorption“ of bicarbonate, although
the bicarbonate ions that actually enter the extracellular fluid are not the same.
The transport of HCO3- accros the basolateral membrane is facilitated by:
1. Na+-HCO3- co-transporter
2. Cl—HCO3- exchange
Although the secretion of hydrogen ions in the late distal tubule and collecting duct
accounts for only percent of the total hydrogen secreted, this mechanism is important
in forming a maximally acidic urine.
In the proximal tubules, hydrogen ion concentration can increase only about threefold
(compared to the filtered load), in the collecting tubule the hydrogen concentration can
be increased as 900-fold.
Late distal tubule and collecting tubules (intercalated cells)
Phosphate and Ammonia Buffers
Minimal urine pH is 4.5, corresponding to an H+ concentration 0.03 mmol/L.
In order, to excrete the 80 mmol of nonvolatile acid formed each day, about
2667 liters of urine would have to be excreted if the H+ remained free
in solution.
500 mmol/day of H+ must be sometimes excreted.
Therefore, whenever an H+ secreted into the tubular lumen combines with
a buffer other than, HCO3- the net effect is addition of a new HCO3to the blood.
A second buffer system in the tubular fluid that is even more important
quantitatively than the phosphate buffer system is composed of ammonia (NH3)
and the ammonium ion (NH4+).
Collecting duct
Proximal tubule, thick ascending limb of the loop of Henle, distal tubule
Glutarate-
Glutamine
α-Ketoglutarate2-
2NH4+ + 2HCO3-
2NH4+
LIVER
(Metabolism)
2NH4+ + CO2
Urea + H20 + 2H+
Loss of 2HCO3- by buffering of 2H+
KIDNEY
(Excretion)
Save 2HCO3- by excretion of 2NH4+
Under conditions of chronic acidosis, the rate of NH4+ excretion can increase to as much
500 mmol/day. Therefore, with chronic acidosis, the dominant mechanism by which
acid is eliminated from the body is excretion of NH4+.
Quantifying Renal Acid-Base Excretion
Net acid excretion = NH+4 excretion + Urinary titratable acid – bicarbonate excretion
Titratable acid represents the nonbicarbonate, non-NH4+ buffer excreted in the urine
(phosphate and other organic buffers)
The most important stimuli for increasing H+ secretion by the tubules are:
1. An increase in pCO2 of extracellular fluid.
2. An in H+ concentration in extracellular fluid.
HCO3- from 24 to 14 mM
pH from 7.4 to 7.17
ΔH+ = (10-7.17 – 10-7.40)
= 0.000068 – 0.000040 mM
= 0.000028 mM
Even though we have added 10 mM,
H+ increased by only 28 nanomols.
Therefore, open system has neutralized
9.999972 mmol of the added 10 mmol H+.
In this system only depletion of HCO3limits the neutralization of H+.
The buildup of CO2 is not a limiting
factor because the atmosphere is an
infinite sink for newly produced CO2.
Henderson-Hasselbalch Equation:
HCO3pH = 6.1 + log
0.03 x pCO2
[HCO3-] (mM)
PCO2 = 20 mmHg
7.1
6 mM
12 mM
24 mM
7.2
8
15
30
7.3
10
19
38
7.4
12
PCO2 =
40 mmHg
pH
24
PCO2 = 80 mm Hg
48
7.5
15
30
60
7.6
19
38
76
7.7
24
48
96
Characteristics of Primary Acid-Base Disturbances
pH
Normal
7.4
H+
40 nmol/L
pCO2
HCO3-
40 mmHg
24 mmol/l
In respiratory acidosis is the initial cause an increase in pCO2 . The compensatory
response is an increase in plasma HCO3-, caused by addition of new bicarbonate
to the extracellular fluid by the kidney. The rise in HCO3- helps to offset the increase
in pCO2, thereby returning the plasma pH toward normal.
Isohydric hypercapnia
(i.e. the same pH at a higher PCO2)
Characteristics of Primary Acid-Base Disturbances
pH
Normal
Respiratory Acidosis
7.4
H+
40 nmol/L
pCO2
HCO3-
40 mmHg
24 mmol/l
In respiratory alkalosis the initial cause is a decrease in plasma pCO2 caused by
hyperventilation.
The compensatory response is a reduction in plasma HCO3- caused by increased
renal excretion of bicarbonate.
Characteristics of Primary Acid-Base Disturbances
pH
Normal
Respiratory Acidosis
Respiratory Alkalosis
7.4
H+
40 nmol/L
pCO2
HCO3-
40 mmHg
24 mmol/l
In metabolic acidosis, is also a decrease in pH (as in respiratory acidosis) and rise in
extracellular fluid hydrogen ion concentration, however in this case the primary abnormality
is a decrease in plasma HCO3-. The primary compensations include increased ventilation,
which reduces pCO2 and renal compensation, which by adding new bicarbonate to the
extracellular fluid helps to minimize the fall in extracellular fluid HCO3- concentration.
1. Diabetes mellitus
2. Diarrhea (loss of large amounts of sodium bicarbonate into the feces)
3. Vomiting of intestinal contents
4. Renal tubular acidosis (a/ impairment bicarbonate reabsorption, b/ inability to secrete H +)
5. Chronic renal failure
Characteristics of Primary Acid-Base Disturbances
pH
Normal
Respiratory Acidosis
Respiratory Alkalosis
Metabolic Acidosis
7.4
H+
40 nmol/L
pCO2
HCO3-
40 mmHg
24 mmol/l
In metabolic alkalosis the primary cause is a rise in HCO3- in extracellular fluid.
The primary compensation is decreased ventilation, which raises pCO 2 and the secondary
compensation is increased renal bicarbonate excretion.
1. Vomiting of gastric content
2. Administration of diuretics (except the carbonic anhydrase inhibitor)
(diuretics cause increases in tubular fluid volume and delivery of sodium to distal parts
of nephron. This in turn leads in increased reabsorption in these parts of the nephron.
Reabsorption is coupled with H+ secretion)
Characteristics of Primary Acid-Base Disturbances
pH
Normal
Respiratory Acidosis
Respiratory Alkalosis
Metabolic Acidosis
Metabolic Alkalosis
7.4
H+
40 nmol/L
pCO2
HCO3-
40 mmHg
24 mmol/l
<
Buffer Base (BB) equal the sum of all the conjugate bases in 1 L of arterial whole blood:
1. HCO3- = 24 mmol/L
2. Protein = 15 mmol/L
3. Hemoglobin = 9 mmol/L
BB = 48 mmol/L
Base Excess (BE) = the observed BB minus the normal BB (BE = - 2 to + 2 mmol/L)
Anion gap (AG) = (Na+ + K+) – (HCO3- + Cl-)
= (142 mmol/L + 4 mmol/L) – (25 mmol/L+ 102 mmol/L)
= 16 mmol/L
Organic acids increase AG, because the lost of HCO3- is not replaced by routinely unmeasured anions.
(Metabolic acidosis)
Respiratory acidosis does not increase AG, because excess H+ is derived from H2CO3 pool, not the
noncarbonic acid pool.
Osmolar gap refers to the disparity between the measured and calculated serum osmolarity.
Increased osmolar gap provides a reasonable good screening procedure for toxins.
Other causes of metabolic acidosis do not affect osmolar gap, since the metabolic acid simply
replaces HCO3-
Poms = 2 [Na+] + [BUN] + [glucose]