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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.
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