The pH scale is a fundamental measure in biology, quantifying the concentration of hydrogen ions in a solution. This logarithmic scale dictates the acidity or alkalinity of an environment, where a small numerical change represents a large shift in ion concentration. Since organisms are fundamentally aqueous systems, the precise regulation of hydrogen ions within and around living cells is required to maintain the chemical reactions that sustain life.
pH and Molecular Function
The three-dimensional structure of large biological molecules, such as proteins, is extremely sensitive to the surrounding pH. Proteins are constructed from amino acids, many of which carry electrical charges that are highly dependent on the hydrogen ion concentration. Changes in pH alter the ionization state of these amino acid side chains, which can disrupt the delicate ionic and hydrogen bonds that hold the molecule in its proper shape.
This structural alteration is known as denaturation, a process that causes the protein to unfold and lose its biological function. Since enzymes are a specialized class of proteins that speed up chemical reactions, they are particularly vulnerable to pH shifts. Each enzyme has an optimal pH range at which it performs its function most effectively.
For instance, the enzyme pepsin operates in the highly acidic environment of the stomach, functioning optimally at a pH of about 2.0. Conversely, the enzyme trypsin, which works in the small intestine, is most active near a neutral pH of 8.0. Deviating from this specific optimum causes the enzyme’s active site to change shape.
The active site is the specialized pocket where the enzyme binds to its target molecule, or substrate. If a change in pH alters the charge distribution on the active site, the substrate can no longer fit correctly. This prevents the chemical reaction from occurring. Because cellular machinery relies on thousands of enzymes working in concert, even slight deviations from the preferred pH can lead to widespread cellular failure.
The Role of Buffers in Maintaining Internal Stability
Living organisms must actively maintain a constant internal environment, a state known as homeostasis, to ensure molecular function remains stable. For humans, the pH of the blood must be kept within a narrow range of 7.35 to 7.45. Any shift outside this range, such as below 7.0 or above 7.7, can result in severe health complications or death.
The body manages this stability through internal chemical systems called buffers, which are solutions that resist changes in pH when an acid or base is added. Buffers consist of a weak acid and a corresponding weak base, working together to neutralize incoming hydrogen ions or hydroxide ions. The most significant of these regulatory mechanisms in the human body is the bicarbonate buffer system.
This system relies on the balance between carbonic acid and the bicarbonate ion within the bloodstream. When metabolic processes generate excess acid, the bicarbonate ions rapidly absorb the surplus hydrogen ions, converting them into carbonic acid. Conversely, if the blood becomes too alkaline, carbonic acid releases hydrogen ions to counteract the increase in base.
This buffering action is tightly regulated by the respiratory and renal systems. The lungs control the carbonic acid component by adjusting the rate of carbon dioxide expiration. If the blood becomes too acidic, breathing speeds up to expel carbon dioxide, which shifts the equilibrium away from acid. The kidneys manage the bicarbonate component, conserving or excreting bicarbonate and hydrogen ions. A failure in this system leads to conditions like acidosis or alkalosis.
pH in Environmental Contexts
The pH of the external environment governs the health of entire ecosystems. The acidity or alkalinity of the soil is a primary factor determining the types of plants that can thrive in a region. Most plants grow best in soil with a slightly acidic to neutral pH, between 6.0 and 7.0.
In acidic soils, a lower pH increases the solubility of toxic elements like aluminum and manganese, which can be absorbed by roots and poison the plant. Lower pH levels also bind up essential nutrients, such as phosphorus, making them unavailable for plant uptake. Alkaline soils, with a pH above 7.5, present problems by reducing the availability of micronutrients like iron, zinc, and manganese, leading to deficiencies.
In aquatic environments, particularly the ocean, small changes in water pH can have extensive ecological consequences. Since the industrial era, the ocean has absorbed a significant portion of atmospheric carbon dioxide, which reacts with seawater to form carbonic acid. This has caused the average ocean surface pH to drop by approximately 0.1 pH units, representing a 30% increase in acidity.
This increase in acidity poses a threat to marine calcifying organisms, including corals, oysters, and certain plankton species. The higher concentration of hydrogen ions reduces the availability of carbonate ions, which these organisms require to build their calcium carbonate shells and skeletons. As a result, shell formation becomes energetically difficult or even impossible, leading to the dissolution of existing structures and disrupting the marine food webs that rely on these organisms.