Terminology

Language is the only instrument of science, and words are but the signs of ideas.

Samuel Johnson ( Preface to his English Dictionary)

Acidemia

Acidemia is a useful shorthand meaning that the pH of the Blood is acid compared to the normal pH of 7.4. Thus a pH of 7.2 is an Acidemia and a pH of 7.6 is an Alkalemia. Because True Neutral is pH 6.8 at body temperature, clinical acid-base values are almost inevitably Alkaline compared to neutral. The pH tells us nothing about the respiratory and metabolic components (see Acidosis and Alkalosis).

Acidosis

An Acidosis, e.g., Respiratory tends to make the pH more acid than usual unless there is a dominating, opposing, Metabolic Alkalosis. The two components frequently occur together, e.g., a Metabolic Acidosis and a Respiratory Alkalosis. One dominates and the other is partially compensating.

Example: a patient with a blood pH of 7.33 is slightlyAcidemic, i.e., on the acid side of the normal 7.4. If his PCO2 is 60 he has a strong Respiratory Acidosis accompanied by a partially compensating Metabolic Alkalosis of 6 mEq/L.

Anion Gap

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.

Base Excess

In 1958 Astrup and Siggaard-Andersen introduced Base Excess to measure metabolic acidosis. This brilliant concept predicts the treatment required to correct metabolic disturbances. However, it would have been easier to understand, e.g., “A metabolic acidosis of 10 mEq/L” or “A metabolic alkalosis of 5 mEq/L”. This would have avoided the frequent use of “Negative Base Excess”.

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. Standard Base Excess (SBE) is also calculated and has the advantage of being the best measure of the metabolic disturbance.

Calculated Bicarbonate Dose

To fully correct a metabolic acidosis, the dose of bicarbonate is calculated from the Standard Base Excess (SBE) 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 Treatable Volume).

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

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.

Dilutional Acidosis

This is the reverse of Contraction Alkalosis.  Diluting the normal 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.

Henderson’s Equation

Carbonic Acid (H2CO3is 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]
Hasselbalch (1916) complicated Henderson’s work using Sorensen’s logarithmic notation see below:

Henderson-Hasselbalch Equation

Hasselbalch modified Henderson’s elegant idea because 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 would be to start with the modified Henderson’s equation.
[H+] = K x [CO2] / [HCO3]
The second step would be 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?

The reason lies partly outside medicine; chemists use the negative logarithm of “K” (pK) as a useful shorthand way of writing a long number. In addition, the same logarithmic version is in widespread use, although known by other names elsewhere. In Copenhagen at the Royal Veterinary and Agricultural University 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. In medicine the Modified Henderson Equation is ideal.

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 test us. The beneficiary is neither the student – nor the patient!

Hindsight throws a critical light on some of the ideas of our legendary forbears. 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:

      1. When a hydrogen atom loses its electron it becomes H+.
      2. A “decrease” in pH means an “increase” in acidity.
      3. “Metabolic Acidosis” is measured as a “Negative” Base Excess.
      4. Hasselbalch added logarithms to complicate Henderson’s simple equation.

Logarithms

Logarithms underlie the belief 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 – [H+] – doubles (from 40 to 80 nMol/1). Using [H+] would make acid-base balance much easier to understand!

Metabolic Acids and Metabolic Acidosis

The best definition is: Metabolic acidosis is a pH too acid for the PCO2. This definition emphasizes the importance of the respiratory component to the overall pH. Except for carbon dioxide, “Metabolic Acids” embraces all the body’s acids. They are not respirable and have to be neutralized, metabolized, or excreted via the kidney.

The pH is always determined by the respiratory and metabolic components, and the latter 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.

Measuring Metabolic Acidosis

The technique introduced by Astrup and Siggaard-Andersen was one of a series proposed between 1916 and 1960.  For all four of these measurements, increased acidosis is accompanied by a decrease in the numerical value – four awkward inverses –al of which contribute to the challenge of understanding Acid-Base Balance.

Neutral

Neutral is the pH at which there are equal numbers of [H+] ions and [OH] ions. 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 0.6 pH units more Alkaline than the intracellular pH (Reeves and Rahn, 1979). This is equivalent to about a 4-fold difference in [H+] concentration: extracellular is 40 nMol/L; intracellular is 160 nMol/L.

PCO2

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).

pH

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 using the words “increase” and “decrease” and use “more acid” and “more alkaline” instead.  To learn more about pH, experiment with the pH playground.

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.

Standard Base Excess

The Standard Base Excess estimates the quantity of Acid or Alkali required, in-vivo, to correct the pH of the entire extracellular fluid including the plasma – like anemic blood (Hb = 5 g/dl). The rationale is that in practice, hemoglobin is effectively buffering all the extracellular fluid – not just the plasma,.

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.

Standard pH

One of the earliest measurements of the metabolic component was Standard pH 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.

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 gradually enters 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 indicated only about half the calculated dose is given before reassessment. This caution is based on several concerns

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