(b) What is the ratio of HCO 3 - to H 2 CO 3 in an exhausted marathon runner whose blood pH is 7.1?

Hey everyone. So here we have an ecstatic asset which is a weak acid. And it's going to associate. And to acetate ion and hydro any um ions and solution. The association constant Acetic acid is 1.78. I'm sent to negative five and it asks if we require a buffer of ph 5.25, what would be the ratio of the conjugate base to the weak asset that we would use? Okay Is 1.78 Times 10 to the - in the ph It's 5.25. And since we're dealing with but first we need to use our Henderson Hasselbach equation and this is going to be ph equals P. K. A. Plus the log of the concentration of the conjugate base divided by the concentration of the weak acid. So we're giving the K. And we need to find the P. K. A. So we're gonna use P. K. A. Equals negative law times K. Eight. So the P. K. A. He goes negative law Times 1.78 Times to the -5. And for P. K. A We're gonna get 4.75. So we're going to get 5.25 Equals 4. plus the law. And we can just call this X. So it's attract both sides by 4. And we're going to get 0.5 equals log of X. So to solve for X We're going to do 10. That is 0.5 power dot exe Which is going to be the concentration of the conjugate base divided by the concentration of the weak acid. It's gonna be 3.16. Thanks for watching my video and I hope it was helpful.

Ch.17 - Additional Aspects of Aqueous Equilibria

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Brown 14th Edition

Ch.17 - Additional Aspects of Aqueous Equilibria

Problem 29b

Chapter 17, Problem 29b

(b) What is the ratio of HCO₃^- to H₂CO₃ in an exhausted marathon runner whose blood pH is 7.1?

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Identify the relevant chemical species involved, which are bicarbonate ion (HCO3-) and carbonic acid (H2CO3).

Use the Henderson-Hasselbalch equation for the bicarbonate buffer system in blood: pH = pKa + log([HCO3-]/[H2CO3]).

Find the pKa value of carbonic acid, which is typically around 6.1 at body temperature.

Substitute the given pH value (7.1) and the pKa value into the Henderson-Hasselbalch equation.

Solve the equation for the ratio [HCO3-]/[H2CO3] by isolating this term and using basic algebraic operations.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Buffer Systems

Buffer systems in the body, particularly the bicarbonate buffer system, help maintain pH levels in the blood. This system involves a dynamic equilibrium between carbonic acid (H2CO3) and bicarbonate ions (HCO3-), which can absorb excess hydrogen ions (H+) or release them to stabilize pH. Understanding this balance is crucial for analyzing changes in blood pH, especially in physiological conditions like exhaustion.

Acid-Base Balance

Acid-base balance refers to the mechanisms the body uses to maintain a stable pH in the blood and other fluids. The normal blood pH range is approximately 7.35 to 7.45, and deviations can indicate metabolic or respiratory issues. In the case of an exhausted marathon runner with a pH of 7.1, this indicates acidosis, which can result from increased lactic acid production and depletion of bicarbonate.

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation relates the pH of a solution to the concentration ratio of an acid and its conjugate base. For the bicarbonate buffer system, it can be expressed as pH = pKa + log([HCO3-]/[H2CO3]). This equation is essential for calculating the ratio of bicarbonate to carbonic acid in the blood, allowing for a quantitative understanding of how the body compensates for changes in pH.

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Henderson-Hasselbalch Equation

Related Practice

Textbook Question

A buffer contains 0.10 mol of acetic acid and 0.13 mol of sodium acetate in 1.00 L. (c) What is the pH of the buffer after the addition of 0.02 mol of HNO3?

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Open Question

A buffer contains 0.15 mol of propionic acid (C2H5COOH) and 0.10 mol of sodium propionate (C2H5COONa) in 1.20 L. (a) What is the pH of this buffer? (b) What is the pH of the buffer after the addition of 0.01 mol of NaOH? (c) What is the pH of the buffer after the addition of 0.01 mol of HI?

Textbook Question

(a) What is the ratio of HCO3^- to H₂CO₃ in blood of pH 7.4?

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Textbook Question

You have to prepare a pH = 3.50 buffer, and you have the following 0.10 M solutions available: HCOOH, CH3COOH, H3PO4, HCOONa, CH3COONa, and NaH2PO4. Which solutions would you use?

You have to prepare a pH = 5.00 buffer, and you have the following 0.10 M solutions available: HCOOH, HCOONa, CH3COOH, CH3COONa, HCN, and NaCN. Which solutions would you use?

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Textbook Question

You have to prepare a pH = 5.00 buffer, and you have the following 0.10 M solutions available: HCOOH, HCOONa, CH3COOH, CH3COONa, HCN, and NaCN. How many milliliters of each solution would you use to make approximately 1 L of the buffer?

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(b) What is the ratio of HCO3- to H2CO3 in an exhausted marathon runner whose blood pH is 7.1?

Verified Solution

Key Concepts

Buffer Systems

Acid-Base Balance

Henderson-Hasselbalch Equation

(b) What is the ratio of HCO₃^- to H₂CO₃ in an exhausted marathon runner whose blood pH is 7.1?