The Biochemistry of Blood pH

To understand how blood pH can change, one needs to understand the basic biochemistry that most of our cells undergo.  This is particularly true for glucose (sugar) metabolism as this is the primary nutrient for our brain and other tissues.

The body metabolizes glucose sugar (C₆H₁₂O₆) in an oxygenated environment  by the following reaction.

The major byproducts are ATP, water, and carbon dioxide (CO₂). When CO₂ dissolves into water it forms carbonic acid, raising the acidity of the blood. As you may be familiar from current conversations about global climate change which is the reaction largely responsible for the acidification of the world’s oceans.

In the body the enzyme carbonic anhydrase catalyzes this reaction.

When the hydrogen ion (H⁺) dissociates from carbonic acid (H₂CO₃), carbonic acid dissociates into bicarbonate (HCO₃^–)

the following buffer system is responsible for maintaining a pH of 7.35 – 7.45 in the human body ^{[3] [4]}

Shifts in the number of reactants will shift the number of products towards the maintenance of equilibria. Acids typically exist in equilibria in any aqueous solution. The amount of protons (H⁺) that disassociate from a certain concentration of acid dictate the strength of that acid. We describe this relationship with the introduction of the acid disassociation constant K_a.

K_a = [H⁺][Conjugate Base] / [Acid]

If we want to relate K_a to pH we must first remember

pH = -log[H⁺]

The final orientation of this equation is known as the Henderson Hasselbalch equation for pH. In our blood buffer system technically the conjugate base is HCO₃^– and the acid is H₂CO₃, however because H₂CO₃ levels are directly correlated to CO₂ levels and we measure CO₂ in the blood, we can simply use our equilbiria into the following equation.

Therefore the modified Henderson Hasselbalch equation to calculate blood pH is:

The negative log of the acid disassociation constant (pKa) for this system at internal body temperature is

The body has two major compensatory mechanisms for managing this buffer system: ventilation, and renal regulation.

Ventilatory Regulation

Respiratory rate determines how much CO₂ is exhaled, thus removing it from the blood. As mentioned CO₂ + water is easily transformed into carbonic acid H₂CO₃. Thus, excess CO₂ in the blood means there is more ‘acid’ in the blood, leading to a lowering of pH. This occurs with a decreased respiratory rate (hypoventilation) which results in retention of CO₂. Conversely, excessive ventilation (hyperventilation) will remove CO₂ form the body meaning a loss of H⁺. This will result in a raise in pH making it more basic. ^[4]

Renal Regulation

Serum levels of bicarbonate are maintained by kidney function. When serum levels are high, the kidneys excrete additional bicarbonate when serum levels are low the kidneys retain bicarbonate, or produce it.

We refer to an undesirable shift in blood pH due to the management of Bicarbonate as metabolic acidosis, because it is managed by internal metabolic regulatory mechanisms. The kidneys are also responsible for proton excretion, by increasing or decreasing the number of protons excreted the kidneys will shift the blood pH, dysfunction of these mechanisms can also lead to metabolic acidosis/alkalosis. ^[4]

Other Systems of pH Regulation

Hemoglobin can bind hydrogen ions which is an additional buffering mechanism that can cause some undesirable effects discussed in our respiratory alkalosis chapter. The phosphate buffer system contains bases which accept hydrogen ions H⁺ and important for the regulation of urine pH. ^[4]

Review Questions

References

1: https://www.ncbi.nlm.nih.gov/books/NBK526028/
2: Fluid, Electrolyte and Acid-Base Disorders: Clinical Evaluation and Management | SpringerLink
3: https://view.officeapps.live.com/op/view.aspx?src=https%3A%2F%2Fwww.asep.org%2Fasep%2Fasep%2FBloodAcid-BaseBuffering.doc&wdOrigin=BROWSELINK
4: :https://www.ncbi.nlm.nih.gov/books/NBK507807/