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Acid-Base Balance

by "Grog" (Alan W. Grogono), Professor Emeritus, Tulane University Department of Anesthesiology

Acid-Base BalanceTerminology

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Terminology Icon"Language is the only instrument of science, and words are but the signs of ideas." – Samuel Johnson ( Preface to his English Dictionary).

Page Index

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Unfortunate Historical Decisions

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 does not appear to be the student.

Hindsight throws a critical light on some of our legendary forbears and their ideas. With contemporary knowledge they would, surely, have made other decisions. We can't go back. But, we can try to understand what they did and how it affects us.

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Awkward Inverses

The History = The Problem
Plus "Taking away an electron should have left us with H¯."  

1.  When a hydrogen atom loses its electron it becomes H+.

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. When a hydrogen atom loses its electron (its charge) it should, intuitively, become H. But, because of Franklin's guess, we are stuck with H+ – the first awkward inverse.

Up "One goes up; the other goes down!"  

2.  A "decrease" in pH means an "increase" in acidity.

In 1909 Sorensen introduced the pH terminology to measure hydrogen ion concentration. This logarithmic notation is wonderful for chemists who deal with a vast range of concentrations. However, clinicians only deal with an [H+] range in the 20 - 80 nMol/L range. The use of the compressed, non-linear, dimensionless scale, underlies this second awkward inverse.

BE "Base Excess predicts treatment required to correct Metabolic Acidosis."  

3.  "Metabolic Acidosis" is described as a "Negative" Base Excess.

Base Excess was introduced by Astrup and Siggaard-Andersen (1958) to measure the metabolic acidosis. This brilliant concept predicts the treatment required to correct metabolic disturbances. However, the name is unnecessary. It would have been easier to write "A metabolic acidosis of 10 mEq/L"; or "A metabolic alkalosis of 5 mEq/L". The name Base Excess underlies this third awkward inverse.

The technique introduced by Astrup and Siggaard-Andersen was one of a series proposed to measure Metabolic Acidosis:

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)

For all four, increasing acidosis decreases the numerical value – four more awkward inverses.

Hasselbalch "Calling this an 'Awkward Inverse' may too polite."  

4.  Hasselbalch complicated Henderson's simple equation.

Carbonic Acid (H2CO3) is a mixture of ionization and dissociation products:

[H+] x [HCO3-] [H2CO3] [CO2] x [H2O]

Henderson (1908) modified this to create the simple Henderson (equilibrium) Equation:

[H+] x [HCO3-] = K x [CO2] x [H2O]

Eight years later Hasselbalch complicated Henderson's work by introducing Sorensen's logarithmic notation to produce the dreaded Henderson-Hasselbalch Equation see below:

pH = pK + log ( [HCO3-] / [CO2] )

Calling this equation an awkward inverse may be too polite. It is a major source of confusion and provides no extra information.

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- 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 is measured as activity in moles per liter. Because pH falls as the acidity increases it is safer to avoid "increase" and "decrease" and use "more acid" and "more alkaline" instead.

To learn more about pH, experiment with the pH playground.

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Log (1000) = 3

The Logarithm is responsible for the mistaken impression that the body maintains remarkably tight control over its hydrogen ion concentration - it doesn't. (Measure blood pressure or pulse using the logarithmic notation and they 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). [H+] would be so much easier to understand!

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Partial Pressure = 40 mmHg (5.33 kPa)

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)

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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 varies from one part of a cell to another but is fairly close to neutral and tightly maintained – it is where most of the body's chemistry occurs. The minute drop of neutral seawater, trapped inside a cell wall millions of years ago, has undergone major structural changes but the pH has changed very little.

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.

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" .....more acid than normal "

Acidemia means that the blood's pH is acid compared to the normal 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 the respiratory and metabolic components (see Acidosis and Alkalosis).

Because, at body temperature, True Neutral is pH 6.8, acid-base clinical values are almost always on the Alkaline side of neutral (see Acidosis and Alkalosis). "Acidemia" is a useful shorthand for "...more acid than the normal blood pH of 7.4".

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" 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, Metabolic Alkalosis. 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.

Example: a patient with a pH of 7.33 has blood on the Alkaline side of neutral but, being on the acid side of the normal 7.4, is Acidemic. If his PCO2 is 60 he has a Respiratory Acidosis with a partially compensating Metabolic Alkalosis of 6 mEq/L.

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"Respiratory Acid = PCO2"

Respiratory Acid and Respiratory Acidosis: Carbon dioxide is respiratory acid because 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.

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" . . . too acid for the PCO2."

Metabolic Acids and Metabolic Acidosis. 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.

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The Anion Gap is the difference between the sum of the major anions and the major cations:

Gap = Na+ + K+ - Cl- - HCO3-

This is discussed in more detail on the page about Clinical Considerations.

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"Further from Neutral"

Contraction Alkalosis: Dehydration ("Contraction") concentrates the mixture of electrolytes in the extracellular fluid (pH 7.4) Because the mixture is normally slightly alkaline, concentration makes the alkalinity more marked and the pH is further away from neutral.

Rehydration is the obvious therapy.

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"Closer to Neutral"

Dilutional Acidosis: 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 appropriate therapy.

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" . . measures neither component . . . "

Bicarbonate: The belief that the bicarbonate level measures Metabolic Acidosis arises because, in a patient with no respiratory abnormality, it does reflect the metabolic disturbance. However, the respiratory and metabolic components both affect the bicarbonate level which cannot, therefore, truly measure 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.

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pH at normal temperature and PCO2

Standard pH: One of the earliest measurements of the metabolic component was proposed in 1916 by Hasselbalch. It was defined as the pH under standard conditions: PCO2 = 40 mmHg (5.33 kPa), temperature = 37oC, and saturated with oxygen.

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bicarbonate at normal temp. and PCO2

Standard Bicarbonate: Bicarbonate itself measures neither the respiratory component nor the metabolic component, but standard bicarbonate does measure the metabolic component. Introduced in 1957 by Jorgensen and Astrup, it was the bicarbonate concentration under standard conditions: PCO2 = 40 mmHg (5.33 kPa), temperature 37oC, and saturated with oxygen. It is the inverse of Hasselbalch's Standard pH.

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Dose to return plasma to normal (mEq/L)

Base Excess: In 1958, a year after introducing Standard Bicarbonate, Astrup and Siggard-Andersen introduced Base Excess as a better method of measuring the metabolic component. It calculated the quantity of Acid or Alkali required, under standard conditions in-vitro, to return the plasma pH to normal.

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Dose to return E.C.F. to normal (mEq/L)

Standard Base Excess is the Base Excess calculated for anemic blood (Hb = 5 g/dl). The rationale for this is that in practice, hemoglobin effectively buffers all the extracellular fluid as well as the plasma, i.e., the behavior is that of anemic blood. It calculates the quantity of Acid or Alkali required, in-vivo, to return the plasma pH to normal.

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0.3 x Wt x BE

Calculated Bicarbonate Dose: 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).

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Treatable Volume = 30% of Body Weight

Treatable Volume: We would "treat" the Extracellular Fluid - about 20% of the body (14L in a 70kg Person) if the administered bicarbonate stayed outside the cell wall. In practice, however, some of the bicarbonate "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. When clinically necessary only about half the calculated dose should be given before reassessment. This caution is based on several concerns

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pH = pK + log ( [HCO3] / [CO2] )

Henderson-Hasselbalch Equation: The starting point is the Henderson Equation:

[H+] x [HCO3] = K x [CO2] x [H2O]

Hasselbalch modified Henderson's elegant idea by regarding the water concentration as constant:

[H+] = K x [CO2] / [HCO3]

He then took logarithms of the remaining components (pK is the negative logarithm of "K"):

pH = pK + log ( [HCO3] / [CO2] )

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.

[H+] = K x [CO2] / [HCO3]

The second is to take the negative logarithm:

pH = -Log ( K x [CO2] / [HCO3] )

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

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Common pH Values: The following list provides the approximate pH values for some common chemicals, foods, and other familiar substances:

 0  concentrated hydrochloric acid
 1  car battery acid, stomach acid
 2  lemon juice
 3  vinegar
 4  orange juice, acid rain
 5  black coffee
 6  milk, saliva
 7  pure water
 8  toothpaste, sea water
 9  baking powder
10  limewater
11  ammonia or washing soda
12  soapy water
13  bleach, oven cleaner
14  sodium hydroxide, drain cleaner
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Acid-Base Tutorial
Alan W. Grogono
Small Logo Copyright Jan. 2020.
All Rights Reserved
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