🤓 Based on our data, we think this question is relevant for Professor Dixon's class at UCF.

We are asked for the ratio of HCO_{3}^{-} to H_{2}CO_{3} in blood of pH 7.4.

Since the solution is a buffer, we can use the **Henderson-Hasselbalch Equation** to calculate the ratio of [HCO_{3}^{-}]/[H_{2}CO_{3}] to have a pH of 7.4:

$\overline{){\mathbf{pH}}{\mathbf{}}{\mathbf{=}}{\mathbf{}}{\mathbf{pKa}}{\mathbf{}}{\mathbf{+}}{\mathbf{}}{\mathbf{log}}{\mathbf{}}\left(\frac{\mathbf{conjugate}\mathbf{}\mathbf{base}}{\mathbf{weak}\mathbf{}\mathbf{acid}}\right)}$

**In the solution, identify the weak acid and the conjugate base. **

*Recall from the Bronsted-Lowry definition that an acid is a proton (H ^{+}) donor and the base is a proton (H^{+}) acceptor*

What is the ratio of HCO_{3}^{-} to H_{2}CO_{3} in blood of pH 7.4?