A chemical measure of acidity or alkalinity, \(\text{pH}\), is defined by the concentration of hydrogen ions (\(\text{H}^{+}\)) in a solution. Maintaining a stable internal environment, known as homeostasis, requires the blood \(\text{pH}\) to be kept within the narrow range of 7.35 to 7.45. This slightly alkaline environment is where the body functions optimally. Enzymes, which catalyze most biochemical reactions, rely on a specific three-dimensional shape to function properly. A significant shift in \(\text{pH}\) causes these proteins to change structure (denaturation), rendering them inactive and threatening survival.
Chemical Buffers: The Immediate Defense
The body’s first line of defense against sudden \(\text{pH}\) changes is chemical buffers, which act within seconds to bind or release \(\text{H}^{+}\) ions. These systems use a pairing of a weak acid and a conjugate base to resist rapid shifts in acid-base balance. Although fast-acting, their capacity to absorb large amounts of acid or base is finite.
The most important buffer system in the extracellular fluid is the bicarbonate buffer system (\(\text{HCO}_{3}^{-}\)). It uses carbonic acid (\(\text{H}_{2}\text{CO}_{3}\)) as the weak acid and bicarbonate ions (\(\text{HCO}_{3}^{-}\)) as the conjugate base. If excess acid enters the bloodstream, bicarbonate ions combine with free \(\text{H}^{+}\) ions to form carbonic acid. Conversely, if the blood becomes too alkaline, carbonic acid dissociates to release \(\text{H}^{+}\) ions.
The process is described by the reversible reaction: \(\text{CO}_{2} + \text{H}_{2}\text{O} \rightleftharpoons \text{H}_{2}\text{CO}_{3} \rightleftharpoons \text{H}^{+} + \text{HCO}_{3}^{-}\). Other systems, such as the phosphate buffer system and proteins like hemoglobin, also contribute to this immediate buffering action.
How Respiration Controls Blood pH
The respiratory system is the body’s second line of defense, controlling \(\text{pH}\) by regulating carbon dioxide (\(\text{CO}_{2}\)) levels. Since \(\text{CO}_{2}\) is readily converted into carbonic acid in the blood, its concentration directly impacts the concentration of \(\text{H}^{+}\) ions. The respiratory center in the brainstem is sensitive to changes in \(\text{H}^{+}\) concentration.
When the blood becomes acidic (low \(\text{pH}\)), the brain signals the lungs to increase the rate and depth of breathing (hyperventilation). This rapidly expels \(\text{CO}_{2}\) from the body. Removing \(\text{CO}_{2}\) consumes \(\text{H}^{+}\) ions and raises the blood \(\text{pH}\) toward the normal range.
If the blood becomes too alkaline (high \(\text{pH}\)), the breathing rate decreases (hypoventilation). Slower breathing retains \(\text{CO}_{2}\), which generates more \(\text{H}^{+}\) ions and lowers the blood \(\text{pH}\). This mechanism provides a relatively rapid response, adjusting \(\text{pH}\) within minutes to a few hours.
The Kidney’s Role in Long-Term Balance
The renal system is the third and slowest mechanism for maintaining acid-base balance, responsible for long-term regulation. Unlike the lungs, which manage volatile acid (\(\text{CO}_{2}\)), the kidneys manage non-volatile acids (fixed acids) like sulfuric and phosphoric acids, which are byproducts of protein metabolism. The renal response takes hours or even days to establish, but it is the only way to eliminate these fixed acids from the body.
The kidneys perform two primary functions: reabsorbing filtered bicarbonate (\(\text{HCO}_{3}^{-}\)) and secreting \(\text{H}^{+}\) ions. Approximately 80 to 90 percent of the filtered \(\text{HCO}_{3}^{-}\) is reabsorbed indirectly in the proximal tubules. \(\text{H}^{+}\) is secreted into the tubule lumen, where it combines with \(\text{HCO}_{3}^{-}\) to form \(\text{H}_{2}\text{CO}_{3}\), which dissociates into \(\text{CO}_{2}\) and \(\text{H}_{2}\text{O}\).
The \(\text{CO}_{2}\) diffuses into the tubule cell, where carbonic anhydrase generates a new bicarbonate ion that is transported back into the bloodstream. For every \(\text{H}^{+}\) secreted, a \(\text{HCO}_{3}^{-}\) is returned to the blood, restoring the body’s buffer stores. The kidneys also excrete excess non-volatile \(\text{H}^{+}\) ions into the urine, often buffered by compounds like phosphate or ammonia, allowing the body to eliminate significant acid loads.
Conditions Resulting from pH Imbalance
When the body’s regulatory systems are overwhelmed, the \(\text{pH}\) balance shifts into a state of imbalance, categorized as either acidosis or alkalosis. Acidosis occurs when the blood \(\text{pH}\) falls below 7.35, indicating an excess of acid, while alkalosis is defined by a blood \(\text{pH}\) rising above 7.45, indicating an excess of base. Both conditions have distinct causes, generally classified as either metabolic or respiratory.
Respiratory acidosis results from \(\text{CO}_{2}\) retention due to inadequate ventilation, such as in cases of severe lung disease. Conversely, respiratory alkalosis is caused by hyperventilation, leading to excessive \(\text{CO}_{2}\) loss and a subsequent decrease in blood acid.
Metabolic acidosis arises from an overproduction of non-volatile acids (e.g., lactic acid or ketoacids) or excessive loss of bicarbonate (e.g., severe diarrhea or kidney failure). Metabolic alkalosis results from the excessive loss of acid, typically through prolonged vomiting, or the intake of alkaline substances. Uncorrected \(\text{pH}\) imbalances can lead to serious consequences, including impaired heart function, confusion, seizures, and coma.