The history of acid-base balance has left us with challenging terminology. This is of special importance to medical students and residents because examiners rely on this confusing material to write test questions. The beneficiary would appear to be the examiner - not the student.
Some of our legendary forbears are to blame for getting us into this mess. Hindsite throws a harsh light on their decisions. They would, almost certainly, have made other decisions if they had enjoyed some of our contemporary knowledge. They didn't. We can't go back. But, we can understand what they did and how it affects us.
|The History = The Problem|
In addition to publishing newspapers, drafting constitutions, and flying kites in thunderstorms, Benjamin Franklin (1706 - 1790) found time to make an unfortunate guess about polarity. After exploring which of his charged substances attracted and repelled each other, he decided to call "vitreous" charges "positive". This decision led, later, to assigning a "negative" charge to electrons. As a result the "flow" of electricity from the positive terminal to the negative is accomplished by electrons travelling "backwards". His decision also affects electrolyte chemistry. In acid-base balance when a hydrogen atom loses its electron (its charge) it should, intuitively become H–. But, because of Franklin's guess, we are stuck with assigning a positive sign. The first awkward inverse: when a hydrogen atom loses its electron it becomes H+.
In 1909 Sorensen introduced the pH terminology to measure hydrogen ion concentration. This logarithmic notation was wonderful for chemists who deal with a vast range of concentrations. However, clinicians only deal with a clinical range in the 20 - 80 nMol/L range. The use of the logarithm explains the fiction that the body maintains remarkably tight control of extracellular acidity. It doesn't! The compressed, non-linear, dimensionless scale, also created the second awkward inverse: a "decrease" in pH means an "increase" in acidity.
Base Excess was introduced by Astrup and Siggaard-Andersen (1958) to measure the metabolic component of acid-base disturbances. This brilliant concept allows us to predict the treatment required to correct metabolic disturbances. Because acidosis is the commonest metabolic disturbance, however, we deal with the third awkward inverse: metabolic acidosis is described as a "negative" base excess. How much easier it would be if we could retain their concept but print reports like: "There is a metabolic acidosis of 10 mEq/L"; or "There is a metabolic alkalosis of 5 mEq/L". No ambiguity and no added terminology!
The technique introduced by Astrup and Siggaard-Andersen was one of a series proposed to measure Metabolic Acidosis. Remarkably, all four of these measurements share the same shortcoming: an increased acidemia causes a decrease in the numerical value - four awkward inverses.
|1916||Standard pH||(Hasselbalch 1916)|
|1957||Standard Bicarbonate||(Jorgensen and Astrup)|
|1958||Base Excess (BE)||(Astrup and Siggaard-Andersen)|
|1960||Standard Base Excess (SBE)||(Siggaard-Andersen)|
In 1916 Hasselbalch complicated Henderson's simple equation by adopting Sorensen's pH notation. Carbonic Acid (H2CO3) is a mixture of ionization and dissociation products:
Henderson (1908) did pioneer work by modifying this to create the familiar equilibrium equation:
Unfortunately, eight years later Hasselbalch ruined Henderson's efforts by adopting Sorensen's unneeded logarithms to produce the dreaded:
Calling this equation an awkward inverse may be too polite. It is a major source of confusion and provides no extra information. However, you can read more about it if you wish.
|- Log (Base 10) of [ H+ ] in Mol/L|
The pH is the negative logarithm of the hydrogen ion concentration. A complete definition requires that the logarithm is defined as being to the base ten and the concentration be measured as activity in moles per liter. Confusion arises because, as the acidity increases, the pH decreases. To avoid mistakes when discussing acid-base balance, it is often safer to avoid "increase" and "decrease" and use "more acid" and "more alkaline" instead.
To learn more about pH, play in the pH playground.
|Log (1000) = 3|
Logarithm is another source of confusion in acid base balance and is responsible for the mistaken impression that the body maintains remarkably tight control over its hydrogen ion concentration - it doesn't. (Blood pressure or pulse measured with a logarithmic notation would appear equally stable).
To understand logarithm, think of "power." Thus 103 = 1000 and log (1000) = 3. When the pH changes by 0.3 units, e.g., from 7.4 to 7.1 the hydrogen ion concentration doubles (from 40 to 80 nMol/1). Wouldn't [H+] be so much easier to understand?
|Partial Pressure = 40 mmHg|
PCO2 is the partial pressure of carbon dioxide. The normal value in arterial blood is 40 mmHg (or 5.33 kPa). The end-exhaled value is usually very similar. Under anesthesia the end-exhaled value is often lower than the arterial value due to several contributing factors. The mixed venous PCO2 is approximately 46 mmHg (6.13 kPa)
|Neutral = pH 6.8 at 37oC|
Neutral is the pH at which there are equal numbers of [H+] ions and [OH-] ions. However, when water is warmed, it ionizes more. Thus, at body temperature water is more ionized than at room temperature and neutral is pH 6.8 rather than 7.0.
The intracellular pH is about 7.0 at body temperature - fairly close to neutral and tightly maintained; it is where most of the body's chemistry occurs. The minute drop of neutral sea water, trapped inside a cell wall millions of years ago, has undergone major structural changes but the pH has changed very little, becoming just a little more alkaline.
The body maintains the blood at pH 7.4, which is about 0.4 pH units more alkaline than the intracellular pH (Reeves and Rahn, 1979). This is equivalent to a 2.5-fold difference in concentration: extracellular - 40 nMol/L; intracellular - 100 nMol/L.
|" .....more acid than normal "|
Acidemia merely means that the blood's pH is acid compared to the usual pH of 7.4. Thus a pH of 7.2 would be called an acidemia and a pH of 7.6 would be called an alkalemia. It tells us nothing about either of the two components, respiratory and metabolic (see Acidosis and Alkalosis).
In the interests of clear thinking remember that, at body temperature, true neutral is pH 6.8. So, nearly all acid-base results deal with a patient whose plasma is actually on the alkaline side of neutral (see Acidosis and Alkalosis). Nevertheless, "Acidemia"is a useful shorthand for "...more acid than the normal blood pH of 7.4".
|" Tending to make the pH more . . . "|
An Acidosis, e.g., Respiratory tends to make the pH more acid than usual unless there is a dominating, opposing alkalosis, e.g., Metabolic. (Alkalosis does the opposite.) A patient often has two components at the same time, e.g., a metabolic acidosis and a respiratory alkalosis. One dominates and the other is partially compensating.
To summarize the use of these words: a patient with a pH of 7.33 has blood on the alkaline side of neutral but, being on the acid side of normal, is acidemic. If his PCO2 is 60 he has a respiratory acidosis with a compensating metabolic alkalosis of 6 mEq/L.
|"Respiratory Acid = PCO2"|
Carbon dioxide is respiratory acid - it is the only acid which can be exhaled via the lungs. Strictly speaking carbon dioxide is a gas, not an acid. Carbonic acid is only formed when combined with water. Nevertheless, clinicians customarily regard carbon dioxide and respiratory acid as synonymous. If you want to sound like a doctor - a High PCO2 is the same as Respiratory Acidosis and vice versa.
|" . . . too acid for the PCO2."|
The best definition is: Metabolic acidosis is a pH which is too acid for the PCO2. This definition emphasizes the importance of the respiratory component to the overall pH. The term "metabolic acids" includes all of the body's acids except carbon dioxide. Metabolic acids are not respirable; they have to be neutralized, metabolized, or excreted via the kidney.
The pH is always determined by the two components, respiratory and metabolic, and the metabolic component is judged, calculated, or computed by allowing for the effect of the PCO2, i.e., any change in the pH unexplained by the PCO2 indicates a metabolic abnormality. Standard Base Excess is the best overall measurement we have of the level of the metabolic acidosis. An adjunct method sometimes used to help identify the source of a metabolic acidosis is the Anion Gap.
The Anion Gap is the difference between the sum of the major anions and the major cations:
This is discussed in more detail on the page about Clinical Considerations.
|"Further from Neutral"|
The pH of the normal mixture of electrolytes in the extracellular fluid (7.4) is relatively alkaline compared to neutral pH at body temperature (6.8). When this electrolyte mixture is concentrated by dehydration, the relative alkalinity is more marked and the pH is further away from neutral.
Rehydration is the obvious therapy.
|"Closer to Neutral"|
The reverse of contraction alkalosis. Diluting the normal slightly alkaline mixture of extracellular electrolytes, also dilutes the alkalinity. This moves the pH closer to neutral at body temperature (6.8)
Diuresis, physiological or therapeutic, is the required therapy.
|" . . measures neither component . . . "|
In acid-base balance, can bicarbonate measure anything? The belief that it measures metabolic acidosis arises because, in a patient with no respiratory abnormality, the bicarbonate does reflect the metabolic disturbance. However, the bicarbonate level is varied by both components, respiratory and metabolic. It cannot, therefore, be an ideal measure of either. Moreover, the relationship between metabolic acidosis and bicarbonate is neither consistent nor linear. And, finally, in acid-base determinations the concentration (in mEq/L) of the bicarbonate ion (HCO3-) is not measured, it is calculated from the PCO2 and pH. When Standard Base Excess (SBE) is available, which it is in acid-base balance measurements, it is the best measure of the metabolic disturbance.
|pH at normal temperature and PCO2|
One of the earliest methods of quantitating the metabolic component was proposed in 1916 by Hasselbalch. It was defined as the pH under standard conditions: PCO2 = 40 mmHg, temperature 37oC, and saturated with oxygen.
|bicarbonate at normal temp. and PCO2|
Bicarbonate itself may be a poor measurement of either the respiratory regulator or the metabolic regulator, but standard bicarbonate is an excellent measurement of the metabolic component. It was introduced in 1957 by Jorgensen and Astrup. It was defined as the bicarbonate concentration under standard conditions: PCO2 = 40 mmHg, temperature 37oC, and saturated with oxygen. It is the inverse of Hasselbalch's Standard pH.
|Dose to return plasma to normal (mEq/L)|
The year after introducing Standard Bicarbonate, Astrup and Siggard-Andersen in 1958 introduced Base Excess as a better method of measuring the metabolic component. In essence the method calculated the quantity of Acid or Alkali required to return the plasma in-vitro to a normal pH under standard conditions.
|Dose to return E.C.F. to normal (mEq/L)|
Standard Base Excess is the Base Excess value calculated for anemic blood (Hb = 5 g/dl) on the principle that this closely represents the behavior of the whole human being. The rationale for this is that in the whole body, hemoglobin effectively buffers the plasma as well as the much larger extracellular fluid, i.e., the behavior is that of anemic blood. The method predicts the quantity of Acid or Alkali required to return the plasma in-vivo to a normal pH under standard conditions.
|0.3 x Wt x BE|
To fully correct a metabolic acidosis, the dose of bicarbonate is calculated from the Base Excess (BE) and the volume to be treated (30% of body weight). This allows for the fact that the treatable space is larger than the extracellular fluid (see next section).
|Treatable Volume = 30% of Body Weight|
We "treat" the Extracellular Fluid - about 20% of the body, or 14L. This would be fine if the bicarbonate we give actually respected the cell wall, as it is meant to, and stayed where we put it. In practice, however, some of the bicarbonate administered "leaks" into the cells. It is, therefore, more accurate to assume that the "Treatable Volume" is 30% of the body (21 L). Then, if for example the BE = -10 mEq/L, the dose which would achieve complete correction would be 21 x 10 = 210 mEq. Today, bicarbonate is usually given only when clinically necessitated; it is also customary to give only about half the calculated dose and reassess. This caution is based on several concerns
|pH = pK + log ( [HCO3-] / [CO2] )|
The starting point is the Henderson Equation:
Hasselbalch modified Henderson's elegant idea by regarding the water concentration as constant and taking logarithms of the remaining components (pK is the negative logarithm of "K"). This resulted in the Henderson-Hasselbalch Equation:
The consequence of using negative logarithms is that "everything is upside down" and incomprehensible to most physicians; it contains the same information as Henderson's simple equilibrium equation. It could have been so much easier; the conversion could have been applied to the whole equation at once. The first step is to write Henderson's equation in the right order with the water concentration omitted as a constant.
The "K" is still "K" and the equation is still recognizable. Why, then, have generations of medical students been taught the Henderson-Hasselbalch version? Why, in fact, were we taught it at all? Were our teachers so mathematically naive that they failed to recognize that the two equations were mathematically equivalent. If so, did they succumb to the temptation to teach us - and therefore test us - using the more complex version?
Part of the reason lies outside medicine; chemists find knowing the negative logarithm of "K" (pK) is a useful shorthand way of writing a long number. In addition, the same logarithmic version is in widespread use, although it is known by other names in other places. At the Royal Veterinary and Agricultural University of Copenhagen it is known as the "Bjerrum equation" in honor of Professor Bjerrum who worked there; and, in the chemical world it is generally known as the buffer equation. (Astrup and Severinghaus, 1986), p 194. There is no need in physiology for us to use this equation. It is part of history's legacy. The Modified Henderson Equation is recommended
The following list provides the appoximate pH values for some common chemicals, foods, and other familiar substances:
|0||concentrated hydrochloric acid|
|1||car battery acid|
|4||orange juice or vinegar|
|11||bleach or washing soda|
Alan W. Grogono
|Copyright Oct 2016.|
All Rights Reserved