Understanding Acid-Base Disorders

The accurate interpretation of laboratory tests in patients with acid-base disorders is critical for understanding pathophysiology, making a diagnosis, planning effective treatment and monitoring progress. This is an important topic particularly for junior medical staff who may encounter acid-base problems outside normal working hours when patients become acutely unwell. These clinical situations may be a source of confusion particularly because of the variety of terms used to describe and classify acid-base disorders. In this article, we aim to provide the reader with an overview of the key concepts necessary for developing a good working understanding of acid-base disorders that commonly present in clinical medicine. We start with some acid-base disorder definitions and then provide a series of case vignettes to illustrate the key points. Case for consideration A 25 year old mechanic is admitted in a very confused and drowsy state. His initial laboratory results are listed below. What diagnoses should be considered? Reference Range Reference Range pH 7.1 7.35-7.45 Na+ 136 mmol/L 136-145 PaO2 13.4 kPa 11.0-14.0 K+ 3.8 mmol/L 3.5-5.3 PaCO2 2.5 kPa 4.5-6.0 Cl- 95 mmol/L 95-108 Bicarbonate 5.3 mmol/L 22-29 Total CO2 6 mmol/L 22-29 Lactate 1.2 mmol/L 0.6-2.4 Urea 4.2 mmol/L 2.5-7.8 Ketones <1 mmol/L Depends on context Creatinine 92 μmol/L 40-110 eGFR >60 mL/min/1.73m2 >60 Open in a separate window DEFINITIONSAcidaemia An arterial pH below the normal range (pH<7.35). Alkalaemia An arterial pH above the normal range (pH>7.45). Acidosis A process lowering pH. This may be caused by a fall in serum bicarbonate and/or a rise in the partial pressure of carbon dioxide (PaCO2). Alkalosis A process raising pH. This may be caused by a rise in serum bicarbonate and/or a fall in PaCO2. ACID-BASE HOMEOSTASIS Like temperature, blood pressure, osmolality and many other physiological parameters, the human body strives to keep its acid-base balance within tightly controlled limits. It is not the aim of this article to review in detail the physiology of acid-base homeostasis, but to provide a working knowledge of some key concepts that will help in the interpretation of results encountered commonly in clinical practice. More detailed free text reviews of acid-base homeostasis are available1-5. A buffer is a solution that resists a change in pH. There are many different buffer systems in the body, but the key one for understanding most acid-base disorders is the bicarbonate system present in the extracellular fluid. Like any buffer, this system comprises a weak acid (in this case carbonic acid, H2CO3) and its conjugate base (the bicarbonate ion, HCO3-), which exist in a dynamic equilibrium as shown in Equation 16: Equation 1 The acidity of a solution is governed by the concentration of hydrogen ions (H+) present. If a disease process results in an increase in the concentration of hydrogen ions, one would expect the body to become more acidic. However, the bicarbonate buffer system resists this change because the excess of hydrogen ions drives the reaction in Equation 1 to the right: hydrogen ions react with and “consume” bicarbonate ions and any change in acidity is minimised. This process requires an adequate supply of bicarbonate ions. The kidneys are vital organs in acid-base balance as they can both generate “new” bicarbonate buffer and reclaim filtered bicarbonate in the proximal tubules (Figure 1). Open in a separate window Fig 1. Acid-base balance is maintained by effective renal and respiratory homeostatic mechanisms