Blood Gases

Acid-Base Physiology

Blood Gases Details

  Blood Gases............................... $49


  • Acids and pH

          The pH is a method of expressing extremely small concentrations of acid in solution. The scale is exponential — a change of one unit is actually a ten-fold change. A solution with pH of 1 has ten times as much acid as a solution with pH of 2.
          The pH scale allows the description of concentrations from 1 to 1/100,000,000,000,000 (10 to the minus 14) moles*/liter. pH is defined as the negative logarithm of the hydrogen ion (H+) concentration or activity.

               pH = - log [H+]

          Water has a pH of 7, with a hydrogen ion concentration of 0.0000001 M/L (-log[10 to the minus 7]=7).
          Acids are substances that can provide a hydrogen ion. Strong acids hold their H+ ion weakly, so it’s free to dissociate and act on other substances. Weak acids hold the hydrogen more tightly, so it doesn’t contribute as greatly to free hydrogen ion concentration — and thus the solution has a higher (less acidic) pH.

          *A mole is a unit of measure based on number of molecules rather than on weight or volume. It’s 6.02246 x 10 to the 23rd power, the number of molecules required to make up a molecular weight in grams. For example, a mole of hydrogen (atomic weight 1.00797, two atoms per molecule) would weigh 2.01594 grams. An “equivalent” — from which the unit milliquivalent (mEq) comes — is the same thing in terms of ionic activity. An equivalent of a substance with a charge of +1 is a mole, while an equivalent of something with +2 charge requires only one-half a mole.

    Back to Top

    Acid Buffering

          A buffer is a substance that resists changes in pH (acid concentration) by undergoing a reversible reaction — in other words, you add a bunch of acid, the buffer undergoes a reaction so that only a small change in pH occurs. Add a bunch of alkali, the buffer changes back to its original state.
          A buffer system consists of a weak acid combined with its salt. An example of a buffer is the carbonic acid (H2CO3) and bicarbonate (H+ HCO3-) system. When the H2CO3 / HCO3- system is at a pH that allows existence of significant amounts of both molecular (undissociated) and dissociated (hydrogen ion has split off) forms, it resists a change in pH by undergoing a change in relative concentrations. When acid is added to a buffer solution, the resulting change in pH is less than it would have been if the buffer were not present.
          When hydrogen ion (H+) is added, much of the hydrogen is taken up by the salt of the buffering acid. With bicarbonate, H+ bonds to HCO3- to form H2CO3, which is a weak acid. With less hydrogen ion available, the solution is less acidic than it would have been without the buffer.
          The second characteristic of a buffer is that the reaction is reversible — the hydrogen ion can be given back. When alkali is added to the solution, carbonic acid provides the proton to neutralize the alkali, again resisting the normally large change in pH that would occur with the addition of the hydroxide (OH-) ion. Back to Top

    Acid-Base Balance in the Body

          The pH of the body must be maintained within a narrow range. Most body systems function optimally at a pH of near 7.4. As the pH changes (either higher or lower), enzymes may cease to function, nerve and muscle activity weakens, and finally all metabolic activity becomes deranged.
          Hemoglobin is one of the most important buffering agents. Its buffering capacity is due to the imidazole chain of the hemoglobin molecule, which contains histadines. Imidazole can accept an extra proton or donate it back at the normal body pH.
          Hemoglobin has an additional property that enables it to maintain pH within the capillaries, known as the Bohr effect. When combined with oxygen, hemoglobin tends to release hydrogen ions that have attached to the imidazole chain (it becomes a stronger acid). When hemoglobin is exposed to acid and lower oxygen concentrations in the capillaries, it gives up the oxygen. It then becomes a weaker acid, taking up extra hydrogen ion. This change maintains the pH in the capillaries essentially the same despite the higher CO2 concentration.
          An opposite change occurs when hemoglobin is exposed to the higher oxygen concentration in the lung. As it takes up oxygen, it becomes more acidic (more prone to release the hydrogen ion). The hydrogen ions react with bicarbonate to form carbonic acid, which in turn is converted to carbon dioxide and released into the alveoli. Hemoglobin is therefore not only an oxygen-transporting molecule, but is also an acid-transporting system. Back to Top

    Carbonic Acid-Bicarbonate Buffering

          The carbonic acid - bicarbonate system is a classic chemical buffer. In addition, the body has the ability to eliminate chemicals from either end of the chemical reaction to maintain the pH. In the case of bicarbonate, the chemical reaction is:
          H+       +       HCO3-       =       H2CO3       =       H2O       +       CO2

          This buffering system is very effective because of the ability to convert carbonic acid to carbon dioxide (through the enzyme carbonic anhydrase) then remove CO2 from the body through respiration. For example, adding enough acid to lower the serum bicarbonate by half would normally drop the pH from 7.4 to 6.0 — but instead all the extra H2CO3 is removed by conversion to CO2. The drop in pH stimulates extra respirations so CO2 (and subsequently more H2CO3) is removed. The pH therefore falls only to 7.3 or 7.2.
          On the other side of the equation, excess acid or excess alkali can be removed through the kidneys.
          Changes in carbonic acid concentration occur rapidly (seconds) in response to hypo- or hyperventilation. On the other hand, changes in bicarbonate require hours or days through the relatively slow process of elimination by the kidney.
          The ratio of bicarbonate to carbonic acid determines the pH of the blood. Normally the ratio is about 20:1 bicarbonate to carbonic acid.
          This relationship is described in the Henderson-Hasselbach equation:

               pH = pK + log (HCO3-/H2CO3)

          (pK is the dissociation constant of the buffer, 6.10 at body temperature. The change in pK with temperature is the reason pH determinations must be adjusted for patients with abnormal temperatures.)

          As carbon dioxide is directly proportional to the carbonic acid (H2CO3), and can be directly measured, it will be substituted into the H-H equation.
               PaCO2 = 33 x H2CO3       or       H2CO3 = .03 x PaCO2

          By substituting,

               pH = pK + log (HCO3-/(PaCO2 x 0.03))

          Thus by measuring serum pH and PaCO2, the serum bicarbonate can be calculated:
               log (HCO3-) = pH + log (PaCO2) - 7.604

    Back to Top

    Total CO2

          “Total CO2" is a value often reported on blood gas slips. Total CO2 is defined as the sum of the carbonic acid and the bicarbonate, or

               TCO2 = [H2CO3] + [HCO3-]

          As the normal ratio of bicarbonate to carbonic acid at physiologic pH is around 20:1, total CO2 will therefore be about 5% higher than serum bicarbonate. When you observe a difference between total CO2 and bicarbonate that is larger than 5%, the patient will be acidotic.
          The total CO2 is not particularly informative by itself. However it will be abnormal in cases of chronic (compensated) acid-base disorders, such as when chronicly elevated carbon dioxide levels cause bicarbonate retention.
          For most clinical decisions, the serum bicarbonate, PaCO2, and pH are used to evaluate acid-base status. The typical clinician probably ignores the TCO2. The diagnosis of causes of acid-base disturbance is discussed later in this manual. Back to Top

    Base Excess

          The “Buffer Base” is the total of all the anionic buffer components in the blood — such as bicarbonate, sulfates, and phosphates. The base excess is the amount of deviation of the patient’s buffer base from normal — in other words, how much extra basic (anionic) chemicals the patient has in his blood, expressed in milliequivalents per liter.
          The base excess is also defined as the amount of acid (in mEq/liter) that would have to be added to the patient’s blood to bring it to normal pH of 7.4.
          Base excess can be a negative value. If the patient is acidotic, acid would have to be taken away to bring the pH to normal. In this case, for example, a base excess of -8 would mean that 8 mEq per liter of base would have to be added to bring the patient’s blood to normal pH.
          Although many physicians use the difference between the patient’s bicarbonate and an average bicarbonate of 24 as an indication of the patient’s need for bicarbonate replacement, the base excess is a more accurate measure because it also takes into account other buffers such as phosphate and hemoglobin. It remains accurate in cases where the buffering capacity of hemoglobin is decreased due to anemia. Back to Top

    Base Excess of ECF

          The base excess of blood does not truly indicate the base excess of the total extracellular fluid (ECF). Because of different protein content and the absense of hemoglobin, ECF has a different buffering capacity. What’s more, each extracellular fluid (for example CSF vs interstitial fluid) has a different buffer status.
          The clinical determination of the amount of bicarbonate required for treatment of severe acidosis is usually based on the base excess of the blood. There is an unavoidable inaccuracy, however, due to several factors:

               1) the time course of the acidosis makes the blood acid
                    poorly reflect the total body acid burden in many cases.
               2) depending on the state of hydration, body fluid distribution varies.
               3) ECF as a percent of body weight varies with age and fat content.

          In general, however, recommendations for bicarbonate therapy are in the range of 0.1 to 0.2 mEq times the body weight times the base excess (ignoring the minus sign).

               bicarb = 0.1 x (-B.E.) x wt in kg

          In other words, with the formula of 0.1 times weight, you’re essentially neutralizing the calculated acid excess in a fluid compartment one-tenth of the body size. If you use a value of 0.2 times weight times BE, you’re treating a fluid reservoir 20% body size. This fluid volume includes the blood, plus fluids that quickly equilibrate with it. Of course, more acid will be present elsewhere, especially if the acidosis is of more chronic nature. Back to Top

    Effects of Acid-Base Balance on Oxygen

          Oxygen status is affected by acid-base status. Oxygen affects the buffering capacity of hemoglobin through the Bohr effect (discussed previously), but the opposite is also true. At a given oxygen pressure, oxygen saturation in the blood is lowered by increasing either carbon dioxide or hydrogen ion concentrations.
          The amount of oxygen carried by the blood is related to PaO2, but the relationship is not linear. Rather, the saturation of hemoglobin with oxygen is related to PaO2 by a sigmoid curve. As the oxygen pressure in blood exceeds 70 mm Hg, only small amounts of additional oxygen are added to hemoglobin.
          As the acidity of the serum increases, less oxygen can be bound at a given PaO2. This mechanism assists hemoglobin in unloading oxygen in the capillaries, where acid concentration is higher. Raising the pH, conversely, increases the oxygen binding, allowing more total oxygen to be carried -- a change that occurs in the alveolus as acid is eliminated through CO2. Back to Top

    Back to Blood Gases Manual Index

Copyright 1996 Mad Scientist Software
Argyle, B., Blood Gases Computer Program Manual.
Mad Scientist Software, Alpine UT, 1996.


Copyright© Mad Scientist Software Inc. All rights reserved.