The normal \(\mathrm{pH}\) of blood is \(7.40 \pm 0.05\) and is controlled in part by the \(\mathrm{H}_{2} \mathrm{CO}_{3} / \mathrm{HCO}_{3}^{-}\) buffer system. (a) Assuming that the \(K_{\mathrm{a}}\) value for carbonic acid at \(25^{\circ} \mathrm{C}\) applies to blood, what is the \(\left[\mathrm{H}_{2} \mathrm{CO}_{3}\right] /\left[\mathrm{HCO}_{3}^{-}\right]\) ratio in normal blood? (b) In a condition called acidosis, the blood is too acidic. What is the \(\left[\mathrm{H}_{2} \mathrm{CO}_{3}\right] /\left[\mathrm{HCO}_{3}^{-}\right]\) ratio in a patient whose blood \(\mathrm{pH}\) is \(7.20 ?\)

Acid-base equilibrium refers to the balance between acidic and basic (or alkaline) components in a solution. In the human body, this balance is essential for maintaining a stable pH, which is critical for various biochemical processes and overall health. When discussing acid-base equilibrium, it is important to understand:

1. **Acids and Bases:** Acids release hydrogen ions (H⁺), while bases accept hydrogen ions. Carbonic acid (H₂CO₃) is an example of an acid, and bicarbonate (HCO₃⁻) is a base.
2. **Dissociation Constants (K_a):** The acid dissociation constant measures the strength of an acid in a solution. For carbonic acid, this is represented as K_a, which has a pK_a value of 6.1 at 25°C.
3. **Henderson-Hasselbalch Equation:** This key equation connects pH, pK_a, and the ratio of an acid and its conjugate base. The equation is: \[\text{pH} = \text{p}K_a + \text{log} \frac{[\text{HCO}_3^-]}{[\text{H}_2\text{CO}_3]}\text{\text{.}}\]
By rearranging the equation based on known values of pH, we can solve for the concentration ratio of acid and base.

Buffer systems are mixtures of weak acids and their conjugate bases, or weak bases and their conjugate acids. They help maintain a stable pH in a solution despite the addition of acidic or basic components. An important buffer system in the human body is the bicarbonate (HCO₃⁻) and carbonic acid (H₂CO₃) system. Here’s how it works:

1. **Mechanisms of Action:** When an excess of H⁺ (protons) is added, bicarbonate ions neutralize them to form carbonic acid, which then converts to CO₂ and water, thus lowering acidity. Conversely, if the environment becomes too basic, carbonic acid dissociates to release H⁺ and neutralize excess OH⁻ (hydroxide ions).
2. **Importance in pH Regulation:** The bicarbonate-carbonic acid buffer system is crucial for maintaining the blood pH close to 7.40, essential for proper cellular function.
3. **Henderson-Hasselbalch Role:** The Henderson-Hasselbalch equation allows us to determine how shifts in the concentrations of H₂CO₃ and HCO₃⁻ affect blood pH.
In scenarios like acidosis (where blood pH falls below the normal range), the buffer systems work hard to restore balance by adjusting the ratio of bicarbonate and carbonic acid.

Blood pH regulation is the process by which the body maintains its pH near the optimal value of 7.40. This is crucial because even slight deviations can severely affect bodily functions. Several systems work together to regulate blood pH:

1. **Buffer Systems:** As mentioned, the bicarbonate-carbonic acid buffer system is pivotal. It handles immediate pH changes by neutralizing excess acids or bases.
2. **Respiratory System:** The lungs help maintain pH by regulating the amount of CO₂ expelled. An increase in CO₂ (which forms carbonic acid in blood) decreases pH, while a decrease in CO₂ increases pH.
3. **Renal System:** The kidneys contribute by excreting excess acids or bases and by reabsorbing bicarbonate from urine. This slower but powerful system ensures long-term pH balance.
Conditions such as acidosis highlight how delicate this balance is. For example, if blood pH drops to 7.20, the body adjusts by altering the H₂CO₃/HCO₃⁻ ratio, as seen from the Henderson-Hasselbalch equation. By knowing pH and pK_a, you can calculate the necessary concentration changes to bring pH back to normal. This integrated approach ensures the body quickly adapts to maintain optimal functionality in varying conditions.