a) Knowing that the pKa of carbonic acid is 6.37, what is the pH of a buffer that has equal concentrations of carbonic acid and bicarbonate? b) Pre...

a) To calculate the pH of a buffer solution with equal concentrations of a weak acid and its conjugate base, you can use the Henderson-Hasselbalch equation:

pH=pKa+log⁡([A−][HA])

pH=pKa+log([HA]

[A−

]

​)

In this case, carbonic acid (H2CO3) is the weak acid, and bicarbonate (HCO3-) is its conjugate base. Given that the pKa of carbonic acid is 6.37, and assuming equal concentrations, the equation simplifies to:

pH=6.37+log⁡(1)=6.37

pH=6.37+log(1)=6.37

Therefore, the pH of the buffer is 6.37.

b) The Henderson-Hasselbalch equation can be rearranged to solve for the ratio of the concentrations:

[H2CO3][HCO3-]=10pH−pKa

[HCO3-]

[H2CO3]

​=10pH−pKa

For a blood pH of 7.4 and a pKa of 6.37:

[H2CO3][HCO3-]=107.4−6.37

[HCO3-]

[H2CO3]

​=107.4−6.37

Calculate this to find the necessary proportion.

c) The carbonate buffer is effective against both acids and bases. This is because it consists of a weak acid (H2CO3

H2CO3) and its conjugate base (HCO3-

HCO3-). The buffer's ability to maintain a stable pH is based on the equilibrium between the weak acid and its conjugate base, allowing it to neutralize both added acids and bases.