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

The ratio,
`([HCO_(3)^(-)])/([H_(2)CO_(3)])`, is 10.7

Step-by-step 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(([HCO_(3)^(-)])/([H_(2)CO_(3)]))`

Here pH=7.40,
`pK_(a)(H_(2)CO_(3))`=6.37

So,
`7.40=6.37+log(([HCO_(3)^(-)])/([H_(2)CO_(3)]))`

or,
`([HCO_(3)^(-)])/([H_(2)CO_(3)])=10.7`

So the ratio,
`([HCO_(3)^(-)])/([H_(2)CO_(3)])`, is 10.7