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Chapter 24: Problem 102

The amino acid glycine can act as a weak acid: If the $1^{\text {st }} \mathrm{p} K_{a}\( for the protonated amino group of glycine is \)9.8,$ what is the ratio of the neutral to anionic form of glycine in blood at $\mathrm{pH} 7.4 ?$

Short Answer

Expert verified
The ratio of the neutral form of glycine to the anionic form of glycine in blood at pH 7.4 is approximately 0.00398, using the Henderson-Hasselbalch equation.

Step by step solution

01

Understand the Henderson-Hasselbalch equation

The Henderson-Hasselbalch equation is given by: \[ \mathrm{pH} = \mathrm{p}K_a + \log{\frac{[\text{Conjugate Base}]}{[\text{Acid}]}} \] In our case: - Acid: Protonated amino group of glycine (neutral form) - Conjugate Base: Anionic form of glycine We are given the pH and pKa value, and we are to find the ratio, which is: \[ \text{Ratio} = \frac{[\text{Anionic form of glycine}]}{[\text{Neutral form of glycine}]} \]
02

Use the Henderson-Hasselbalch equation to find the ratio

Plug in the given values into the Henderson-Hasselbalch equation: \[ 7.4 = 9.8 + \log{\frac{[\text{Anionic form of glycine}]}{[\text{Neutral form of glycine}]}} \]
03

Solve for the ratio

We'll perform the following steps to solve for the ratio: 1. Subtract 9.8 from both sides of the equation: \[ -2.4 = \log{\frac{[\text{Anionic form of glycine}]}{[\text{Neutral form of glycine}]}} \] 2. Use the properties of logarithms to remove the log: \[ 10^{-2.4} = \frac{[\text{Anionic form of glycine}]}{[\text{Neutral form of glycine}]} \] 3. Calculate the value of \(10^{-2.4}\): \[ 0.00398 \approx \frac{[\text{Anionic form of glycine}]}{[\text{Neutral form of glycine}]} \] The ratio of the neutral form of glycine to the anionic form of glycine in blood at pH 7.4 is approximately 0.00398.

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Most popular questions from this chapter

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