A buffer problem, I need serious help

[afcwestwarrior]

A buffer problem, I need serious help!

3. Blood has a normal pH of 7.35-7.45 and contains two major buffer systems. It is important that the pH of blood remains relatively constant because at pH below 6.8 or greater than 8.0, cells cannot function properly and death may result. The HCO3-/CO2 (aq) blood buffer in vivo is an open system in which the concentration of dissolved CO2 is maintained constant. Any excess CO2 produced by the reaction H+ + HCO3- ---> H2O + CO2 is expelled by the lungs. Note that a typical laboratory buffer is a closed system. The concentration of conjugate acid increases when H+ reacts with the conjugate base.
You calculated the Keq and pK of Reaction (4) from the following reactions and K values in lecture.
CO2 (g) <==> CO2 (aq) K1 = 3 x 10-5 at 37oC.
CO2 (aq) + H2O (l) <==> H2CO3 (aq) K2 = 5 x 10-3 at 37oC
H2CO3 (aq) <==> H+ (aq) + HCO3- (aq) pKa = 3.8 at 37oC
CO2 (aq) + H2O (l) <==> H+ (aq) + HCO3- (aq) K4 = ?
You calculated the [HCO3-] = 0.024 M in blood at pH 7.4. Calculate the [CO2 (aq)] in blood at this pH.
0.01 M H+ is added to blood. You calculated the pH of blood under conditions such that the increased [CO2 (aq)] can be released as CO2 (g). In other words, assume that the blood buffer is an open system. Remember that the [CO2 (aq)] remains constant in this open buffer system.

a. Cells cannot function property if the pH of blood falls below 6.8 or rises above 8.0. Calculate the amount in M of H+ that is added to blood for the blood pH to fall to 6.8.

here's what i did i took the PH of 6.8 and i took the antilog which is 10^-6.8
which equals 1.58*10^-7 M which equals the concentration of H+, but I am not sure if I am right.

 

[afcwestwarrior]

wheres the help, u didnt even say whether i was right or wrong

 

[Blueshift5]

I am happy to see that you are attempting to solve this buffer problem. It is important to understand the concept of buffers and their role in maintaining the pH of blood within a narrow range. Your approach to calculating the amount of H+ added to blood is correct. To confirm your answer, you can also use the Henderson-Hasselbalch equation: pH = pKa + log ([A-]/[HA]). In this case, [A-] is HCO3- and [HA] is H2CO3. Plugging in the values, we get pH = 3.8 + log (0.024/0.024) = 3.8 + 0 = 3.8, which is the same as the pH at which cells cannot function properly.

In terms of the buffer system being an open system, it means that the concentration of dissolved CO2 remains constant and any excess CO2 produced is expelled by the lungs. This is important for maintaining the pH of blood, as excess CO2 can lead to a decrease in pH and affect cell function.

For the second part of the problem, you correctly calculated the [CO2 (aq)] in blood at pH 7.4. This concentration is important for maintaining the equilibrium between CO2 (aq) and H2CO3 (aq).

Overall, it is crucial to understand buffer systems and how they function in maintaining the pH of blood. Keep practicing and solving problems to strengthen your understanding.