Blood is buffered by carbonic acid and the bicarbonate ion.Normal blood plasma is 0.024 M in HCO3- and 0.0012 M H2CO3(pKa1 for H2CO3 at body temperature is 6.1). a. What is the pH of blood plasma? b. If the volume of blood in a normal adult is 5.0 L, what mass of HCl can be neutralized by the buffering system in blood before the pH falls below 7.0 (which would result in death)? c. Given the volume from part b, what mass of NaOH can be neutralized before the pH rises above 7.8?

The Henderson-Hasselbalch equation is a mathematical expression used in chemistry to estimate the pH of a buffer solution. A buffer solution is a special kind of solution that resists changes in pH upon the addition of small amounts of acid or base. Its formulation is as follows:
\[ \text{pH} = \text{p}K_a + \log\left(\frac{[\text{A}^-]}{[\text{HA}]}\right) \]
where \( \text{p}K_a \) is the negative log of the acid dissociation constant (Ka), \( [\text{A}^-] \) represents the concentration of the conjugate base, and \( [\text{HA}] \) represents the concentration of the acid.
In the context of blood buffering systems, carbonic acid (\( H_2CO_3 \)) and bicarbonate (\( HCO_3^- \)) form such a pair. In the exercise provided,To calculate the pH of blood plasma, we leverage the concentration values for bicarbonate and carbonic acid, alongside the specified \( \text{p}K_a \) for carbonic acid, and apply them to the Henderson-Hasselbalch equation. This application allows us to comprehend the buffering capability of blood and determine the pH under various scenarios.

pH calculation is fundamental in understanding the acidity or basicity of a solution. The pH scale ranges from 0 to 14 with lower values indicating acidity, higher values indicating alkalinity, and 7 representing neutrality.
To calculate the pH, one must first understand the relationship between hydrogen ion concentration (\( [H^+] \)) and pH, which can be represented by:
\[ \text{pH} = -\log([H^+]) \]
In the exercise, we bypass direct hydrogen ion concentration calculation by using the Henderson-Hasselbalch equation since it provides an indirect method to calculate pH already considering the acid-base equilibrium.
A deep understanding of pH is crucial not only in chemistry but also in biological systems, such as maintaining the precise pH range of blood plasma, which is essential for proper physiological function.

Acid-base neutralization is a chemical reaction where acid and base properties are eliminated, producing salt and usually water. It's an essential process in the context of buffer systems, such as the blood buffer system in our exercise.
For instance, when the blood buffer system encounters a strong acid like HCl, bicarbonate ions neutralize it, which can be represented by the chemical reaction:
\[ HCl + HCO_3^- \rightarrow Cl^- + H_2CO_3 \]
The exercise explores the capacity of the blood buffer system by calculating how much HCl can be neutralized before a critical pH level is reached. Conversely, when a strong base like NaOH is added, the following reaction occurs:
\[ NaOH + H_2CO_3 \rightarrow Na^+ + HCO_3^- + H_2O \]
Once more, the exercise calculates the mass of NaOH that can be neutralized by the blood before reaching a pH that compromises human health. These neutralization reactions underpin the essential role of buffer systems in maintaining pH stability in biological systems.