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INTRODUCTION

The human body needs to regulate free hydrogen ions (H+) within a narrow window in order to maintain proper protein structure and function. In the process of metabolizing carbohydrates, proteins, and fats, 15,000 mmol of volatile acid, carbon dioxide (CO2), is generated and another 50 to 100 mEq of nonvolatile acid is also produced.1 Despite this, free hydrogen ion concentration is maintained at level of 40 nmol/L (10−6 mmol/L).2 When describing hydrogen concentration of physiologic solutions, we refer to the pH of the solution rather than the actual concentration. The pH of a solution is defined by the following equation:

pH = − log [H+]

 

The pH compatible with human life is in the range of 6.8 to 7.8.3 In order to maintain serum pH in this range, the human body requires the following:2

  • an effective buffer system to prevent wide fluctuations of pH in response to small additions or subtractions of acid,

  • the ability to excrete volatile acids (CO2) via the lungs, and

  • the ability to excrete nonvolatile acids (sulfuric acid, phosphoric acid, and ammonium) via the kidneys.

The bicarbonate/carbon dioxide buffer system is how clinicians usually analyze acid–base balance at the bedside. Carbon dioxide and bicarbonate are easily measured; bicarbonate is important simply because of its high concentration in the blood.

Carbon dioxide effectively becomes an acid when dissolved in solution as described in the following reaction:

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The relationship of these reactants can be described by the Henderson-Hasselbalch equation:4

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where Paco2 is the partial pressure of CO2 in arterial blood. The amount of CO2 dissolved in solution is proportional to the partial pressure of CO2(Pco2), which is in equilibrium with Paco2 of alveolar air. The solubility constant of CO2 is 0.03.2

The effectiveness of the buffering system can be appreciated when studying this equation. One measure used to quantify strength of acid in solution is the dissociation constant, or pKa. In general, a weak acid will be the most effective buffer when pH is within 1 log of its pKa. Bicarbonate is not an excellent buffer at physiological pH (its pKa is only 6.1). However, because CO2 and bicarbonate can be independently regulated (by the lungs and kidneys, respectively), bicarbonate can be an effective buffer.2 Of course, other extracellular buffers such as phosphates and plasma proteins do contribute to preserving pH; however, they are not considered to be as important. Intracellular buffers, such as hemoglobin, are very important to the initial buffering of respiratory acidosis and alkalosis, and will be discussed later in the chapter.

Although the values change depending on the laboratory or reference, generally, acidemia refers to ...

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