Question

uman blood is kept at a typical pH of 7.40 mainly by the carbonic acid–carbonate ion buffer system. The corresponding chemical equation describing the buffer would be: H2CO3(aq) H2O(l) ⇌ HCO3-(aq) H3O (aq) The appropriate Ka value for carbonic acid is 4.3×10-7, so its pKa value is 6.37. Calculate the [base]/[acid] ratio in human blood. (Answer to 3 significant figures, no units)

Answer 1

Answer:

The ratio, $\frac{[HCO_{3}^{-}]}{[H_{2}CO_{3}]}$, is 10.7

Explanation:

$H_{2}CO_{3}$ is an weak acid and $HCO_{3}^{-}$ is it's conjugate base.

So, according to Henderson-Hasselbalch equation, pH of this buffer system can be represented as-

$pH=pK_{a}(H_{2}CO_{3})+log(\frac{[HCO_{3}^{-}]}{[H_{2}CO_{3}]})$

Here pH=7.40, $pK_{a}(H_{2}CO_{3})$=6.37

So, $7.40=6.37+log(\frac{[HCO_{3}^{-}]}{[H_{2}CO_{3}]})$

or, $\frac{[HCO_{3}^{-}]}{[H_{2}CO_{3}]}=10.7$

So the ratio, $\frac{[HCO_{3}^{-}]}{[H_{2}CO_{3}]}$, is 10.7