How Does Adding Nitric Acid Affect the pH of a Carbonate Buffer Solution?

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Discussion Overview

The discussion revolves around the effect of adding nitric acid to a carbonate buffer solution, specifically focusing on the resulting pH changes. Participants explore the application of the Henderson-Hasselbalch equation and the stoichiometry of the acid-base reaction within the context of a homework problem.

Discussion Character

  • Homework-related
  • Mathematical reasoning
  • Technical explanation
  • Conceptual clarification

Main Points Raised

  • One participant calculates the initial pH of the carbonate buffer solution as 10.50 using the Henderson-Hasselbalch equation.
  • Another participant questions how the addition of nitric acid affects the concentrations of carbonate (CO32-) and bicarbonate (HCO3-).
  • It is proposed that the addition of HNO3 increases the concentration of H3O+, shifting the equilibrium to produce more HCO3- and decrease CO32-.
  • A participant suggests that the protonation of carbonate goes to completion, implying that no carbonate ions remain in the solution.
  • There is a discussion about determining the limiting reagent in the reaction.
  • One participant attempts to calculate the new concentrations of HCO3- and CO32- after adding nitric acid, arriving at a pH of 10.35, which is later challenged by another participant.
  • Another participant points out that the total concentration of carbonate and bicarbonate must remain constant, leading to a correction in the calculated concentrations.
  • There is a clarification regarding the effect of volume change on the Henderson-Hasselbalch equation, emphasizing that the ratio of concentrations remains constant despite volume changes.

Areas of Agreement / Disagreement

Participants express differing views on the final concentrations of bicarbonate and carbonate after the addition of nitric acid, with some calculations being corrected and refined throughout the discussion. The discussion remains unresolved regarding the exact pH after the acid addition, as participants are still working through the implications of their calculations.

Contextual Notes

Participants highlight the importance of stoichiometry in acid-base reactions and the assumptions involved in applying the Henderson-Hasselbalch equation. There are unresolved questions about the impact of volume changes on concentration ratios.

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Homework Statement



What is the pH of a carbonate buffer solution prepared by mixing 1.500 mol Na2CO3 and 1.000 mol NaHCO3 and adding water to make a 1.000 L solution? (pKa of (HCO3)- is 10.32)

A 200.00 mL sample of 1.000 M nitric acid was added to the buffered solution above. What is the resulting pH?

Homework Equations



Henderson-Hasselbalch equation (pH = pKa + log ([A-]/[HA]))
pH = -log [H+]

The Attempt at a Solution



I used the equation for the first part, and I calculated a pH of 10.50 (pH = 10.32 + log (1.500/1.000)). Is this correct? I'm not sure how to include the second part though... I'm guessing I'm supposed to calculate the resulting pH when nitric acid is added to 1.000 L of water and then somehow use that pH for the second part... If that makes any sense. Any help is appreciated!
 
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10.50 looks OK, although you added, not multiplied (I guess that's just a typo).

What reaction takes place in the solution when you add strong acid? How does it change amounts of CO32- and HCO3-?
 
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Oops, that is a typo =P Apologies!

So if I add HNO3, it would dissociate, creating more H3O+. And the addition of hydronium ions would shift the equilibrium to increase the concentration of HCO3-, which decreases the concentration of CO32-. Is that right?
 
You are on the right track. Assume carbonate protonation went to completion.
 
Does that mean there wouldn't be any carbonate ions left in the solution?
 
Check what is a limiting reagent.
 
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So this is what I ended up doing: 0.200 L of 1.000 M nitric acid was added, which means that 0.200 moles of nitric acid was added. Since it's a strong acid, it fully dissociates, meaning that there are 0.200 additional moles of hydrogen ions in the solution. Since there are more protons in the solution, the equilibrium shifts to create more HCO3-, meaning that CO32- are used. Using the ICE method, I figured that the concentration of HCO3- is 1.000 M and that of CO32- is 1.083 M. Then I used the H-H equation and got this: pH = 10.32 + log (1.083 M / 1.000 M) = 10.35. Does that look right?
 
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No. As told - use just the stoichiometry of the reaction.

Simple check - you started with total concentration of carbonate and bicarbonate being 2.5M, now you are down to 2.083 - so there must be something wrong, carbon can't just disappear from the solution (unless it bubbles out as CO2, but you are not protonating bicarbonate, so this is not the case).
 
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I'm not sure I'm understanding correctly - would the correct concentrations be 1.2 M and 1.3 M?
 
  • #10
Yes.

Assuming you assigned correct concentration to correct form.
 
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  • #11
Oh, I see! Thank you =)
 
  • #12
Just one more question: so the volume doesn't change even after adding 0.200 L of nitric acid? Wouldn't it be different?
 
  • #13
Changes - but it doesn't matter. You see, in the Henderson-Hasselbalch equation under log there is a ratio of two concentrations. If you add water, both concentrations change in the same way, so their ratio stays constant. In fact, it can be easily shown that volume cancels out, and it is just ratio of number of moles of acid and conjugate base that counts.

Beware: there are some traps here. Please read the page about HH equation that I linked to to find out what is the problem.
 
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