Question

In the blood plasma, \(\mathrm{pH}\) is maintained by the carbonic acid- bicarbonate buffer system. a. How is pH maintained when acid is added to the buffer system? b. How is pH maintained when base is added to the buffer system?

Short answer

When acid is added, \text{HCO}_3^- neutralizes \text{H}^+. When a base is added, \text{H}_2\text{CO}_3 dissociates to restore \text{H}^+.

Step-by-step solution

Understanding the Carbonic Acid-Bicarbonate Buffer System

The carbonic acid-bicarbonate buffer system maintains the \text{pH} of blood plasma. It involves the equilibrium between carbonic acid (\(\text{H}_2\text{CO}_3\)) and bicarbonate ions (\( \text{HCO}_3^- \)). The chemical equation for this equilibrium is: \(\text{H}_2\text{CO}_3 \rightleftharpoons \text{H}^+ + \text{HCO}_3^-\).

Process When Acid is Added

When an acid is added to the buffer system, the \text{H}^+ ions from the acid increase the concentration of hydrogen ions in the blood. The added \text{H}^+ ions react with bicarbonate ions \(\text{HCO}_3^-\) to form carbonic acid \(\text{H}_2\text{CO}_3\), shifting the equilibrium to the left. Thus, less free \text{H}^+ remains in the solution, maintaining the \text{pH} within the narrow range.

Process When Base is Added

When a base is added, it reacts with \text{H}^+ ions to form water, reducing the concentration of \text{H}^+ ions. To counteract this, carbonic acid \(\text{H}_2\text{CO}_3\) dissociates into bicarbonate ions \(\text{HCO}_3^-\) and \text{H}^+ ions, shifting the equilibrium to the right. This release of \text{H}^+ ions helps maintain the \text{pH} by replenishing the lost hydrogen ions.

Key concepts

blood plasma pH balance

Blood plasma pH balance is crucial for maintaining the body's physiological functions. The pH scale measures how acidic or basic a solution is, ranging from 0 (very acidic) to 14 (very basic), with 7 being neutral. Blood plasma typically has a pH of around 7.4. A deviation from this range can be dangerous and disrupt processes such as oxygen transport and enzyme activity.
The carbonic acid-bicarbonate buffer system plays a critical role in keeping the blood plasma's pH within this narrow range. It is a balancing act involving chemical reactions that neutralize changes in hydrogen ion (H⁺) concentrations. This system ensures that the body can function properly, even when metabolic activities produce acids and bases.
Understanding how the body maintains this balance enables us to appreciate the complexity of our body's internal environment.
It serves as a foundation for studying more detailed topics, like acid-base equilibrium and other buffer systems.

acid-base equilibrium

Acid-base equilibrium describes the balance between acidic and basic (alkaline) substances in the body. This balance is vital for many biochemical processes. It involves a series of chemical reactions that either release or absorb H⁺ ions, helping maintain the pH of bodily fluids.
In the carbonic acid-bicarbonate system, the equilibrium reaction is: \(\text{H}_2\text{CO}_3 \rightleftharpoons \text{H}^+ + \text{HCO}_3^-\). This means that carbonic acid (H₂CO₃) can dissociate into hydrogen ions (H⁺) and bicarbonate ions (HCO₃^-), and vice versa.
When there is an excess of H⁺ ions (more acidic conditions), the reaction shifts to the left, forming more carbonic acid and reducing free H⁺ ions to stabilize pH. Conversely, when there are fewer H⁺ ions (more basic conditions), carbonic acid dissociates, releasing H⁺ ions and shifting the equilibrium to the right, thereby stabilizing pH again. This delicate balance ensures that the body's internal environment remains stable and can adapt to changes from metabolism or external factors.

buffer systems

Buffer systems are combinations of weak acids and their conjugate bases (or weak bases and their conjugate acids) that resist changes in pH when small amounts of acid or base are added. They help to maintain the pH of a solution within a narrow range.
In biological systems, the carbonic acid-bicarbonate buffer system is one of the most important. Other buffer systems include the phosphate buffer system and protein buffer system, all of which work together to keep the body's pH stable.
The overall effectiveness of a buffer system depends on the concentration of its components and the pKa (acid dissociation constant) value of the acid involved. Ideally, a buffer works best when the pKa is close to the desired pH of the solution.
Understanding buffer systems is essential for comprehending how the body controls pH under varying conditions. This knowledge applies not only to physiology but also to fields like chemistry and environmental science, showing the widespread importance of pH regulation.