.
[1] Confusion (often referred to as a paradox) can occur
if only a single factor of say an output or an input is noted with
respect to present status of the internal environment. For example
the "paradoxical" acidic urine which occurs when the blood has an
abnormally high pH due to loss of say acidic gastric juice in pyloric
obstruction. The pH of blood rises because acidic gastric juice is
lost ie.HCl. The "usual" or "normal" acid urine is due to loss of
inorganic acids H2SO4
and H3PO4 which are the end products of
protein and amino-acid metabolism. The explanation is that HCl loss
was the cause of the rise in pH. This would be corrected if HCl or a
Cl containing salt is given. For the urine not to be acidic sulphuric
and phosphoric acids would have to be retained as these are the acids
excreted by the kidneys. This would be theoretically possible but it
not what the renal system happens to do. It becomes a paradox only
when the conventional explanation of urine pH control is used. This
depends on the concept of controlling pH by variable reabsorption of
HCO3.
[2] Measurement of Buffer Base and
[HCO3].
See Appendix 4.3. Note
that Base excess is not measured but derived from pH with or without
PCO2 measurements.
[3] The terms buffer and buffering are used in medical literature with extra meanings to those stated here. Much is made of the conept of 'physiological' buffering as distinct from chemical buffering. These mechanisms are discussed in section 3.3 "Respiratory Control of pH". Using the term "buffer" for these mechanisms of addition or subtraction of acids and bases is confusing. It is fundamentally different from chemical buffering as understood in Chemistry.
[4] Bicarbonate ion can act as an acid at high pH levels. At such levels it can either donate or accept hydrogen ions and therefore act as a buffer.
HCO3H
+ CO22-
[5] pH when PCO2 is corrected to 40mmHg. See Section 4.2.2
[6] Pickering (1978) discusses this problem with hypertension and normotension in relation to measured arterial blood pressure
[7] I think "aetiological" should have been "chemical". The aetiological factor is the disease process, e.g. diabetes, chronic bronchitis etc
[8] See Appendix 3.1 which is on the concept of the universal constancy of "normal values".
[9] The
H2CO2 concentration is
really only a small fraction of the dissolved
CO2. It is convenient to regard dissolved
CO2 as
H2CO2 as it can change
to H2CO2 if
[H]
or [HCO3]
fall. The constant pK in the Henderson-Hasselbalch equation takes the
equilibrium between CO2 and
H2CO2 into account. In
blood some CO2 is carried as
HCO3
derived by the action of buffers on
H2CO2
H
+ HCO3.
H is
removed, therefore more HCO3
is formed. Another quantity of CO2 is combined
directly with protein to form carbamino compound. The true dissolved
CO2 is in equilibrium with
H2CO2,
HCO3
and carbamino.
[10] The usual description is the added acid or base is buffered by the buffer anions.
[11] This might be considered a theorem which is derived from the law of conservation of matter or a self evident trivial statement. Its implications are frequently ignored in clinical and physiological literature.
[12] If the patient's in vivo PCO2 is corrected to 40mmHg, the pH change is slightly different from that which occurs in vitro (Brackett et al. 1965, Prys-Roberts et al 1966). The difference is small and does not cause confusion in interpreting clinical problems. (See Appendix A4.2 and Brackett afternote to table 4.2.2.).
[13] A similar concept was developed by Whitehead (1965)
using nanoequivalents [H]/litre.
[14] It is clear from Brackett et al (1965), that there is a small apparent non-respiratory acidosis if pH-PCO2 titration curve is constructed in vitro in a purely respiratory acidosis. An error of similar magnitude is seen in vitro if a respiratory acidosis is superimposed on a non-respiratory acidosis. (Kappogoda et al 1970 and Stocker et al 1972). (See footnote 3 to table 4.2.2).
[15] Deviations in actual pH are less in a compensated disturbance than in the corresponding disturbance without compensation
[16] It is possible that non-respiratory pH might be 7.4
from the exact neutralisation of, for example, the high pH of HCl
lack, in pyloric obstruction by lactic acid due to hypoxia. If this
were the case the two aetiological factors necessary for such an
occurrence should be apparent. Also chemical analysis would show a
high lactic acid level and a low Cl
level in the blood.
[17] This picture and the next one can occur during recovery from chronic CO2 retention when the PaCO2 is lowered. Clinical evidence will distinguish this from primary non-respiratory alkalosis.
[18] Sometimes intracellular pH changes are postulated as part of the explanations of phenomena eg. acidic urine when blood pH is high.
[19] No other anions are present in the blood in sufficient quantity to make any other acid.
[20] It is axiomatic that if electrical neutrality is
maintained a change in [HCO3]
cannot occur without a change in the concentration of the cations or
another anion.
[21] The terms 'metabolic' and 'non-respiratory' acidosis are synonymous. In practice they are also the same as a substantial lowering of the base excess, non-respiratory pH or any of the blood bicarbonate measurements. The converse applies to alkalosis.
[22] The classical explanation is that
HCO3
is diluted by the intravenous solution, hence the designation
dilutional.
[23] Diuretic alkalosis has been referred to as
'contraction' alkalosis (Cannon et al, 1965). This terminology is
analogous to dilutional alkalosis (6.3.3.1.3.2). The idea is that the
extracellular volume decreases and contracts around the
HCO3
mass. The concentration of HCO3
therefore rises. This explanation does not explain why the other
solutes of the extracellular fluid are excreted but not
HCO3.
Now that it is recognised that Cl
deficiency is invariably associated with diuretic alkalosis the
contraction hypothesis is unnecessary.
[24] As soon as the PaCO2 is lowered slightly, the pH will be greater than 7.4. This level of pH may so diminish respiratory drive that the PaCO2 cannot be lowered further without artificial correction of the "compensation".
[25] Kassirer, 1974, in an article "Serious Acid-Base Disorders" states that a certain formula will enable one to calculate how much NaHCO2 would be required to raise the serum bicarbonate by a certain amount and then says "One should not assume, however, that the change in bicarbonate anticipated from this calculation will always occur "!!!
[26] C.S.F. pH changes are much discussed but there is little objective knowledge about the topic. Plum and Price (1973) discuss some of the problems involved.
[27] Jones, W.J. et al, Systemic and renal acid-base effects of chronic dietary potassium depletion in humans. Kidney International 21, 402-410, 1982. In this study a stable state of potassium deficiency was induced in seven subjects. The mean body deficit of potassium was 4.2mEq/kilogram. There was a mean rise in serum bicarbonate of 2mEq/litre. This change was statistically significant but obviously of no clinical importance. Such a small change in pH or bicarbonate is of no importance and is certainly within the variation to be expected between different patients and in the same patient from time to time. It should be noted that statistical "significanct" as a term is probably undesirable as it implies importance of the result upon which statistical "significance" has been conferred (Model, 1981). A trivial result does not acquire meaning or importance because the result has a low probability of being due to chance
[28] Total CO2 is the "CO2" measured by the S.M.A.12-60 Sequential Multiple Analysis Technicons known to me.
[29] Singer and Hastings' definition of buffer-base was
actually "Base equivalent of the sum of the buffer anion
concentration". This was because they called the Na
a base and the anions, acids