Protein pH titration (Henderson-Hasselbalch)

In summary, the question involves calculating the final pH of a histidine solution after 2.5 moles of NaOH are added, using the Henderson-Hasselbalch equation. The professor's approach of using 0.5 mol increments of alkali to neutralize protons may not be accurate. The histidine can be treated as a triprotic acid, similar to phosphoric acid in a buffer problem. However, the wording of the question may be misleading.
  • #1
physicsstudent14
13
0

Homework Statement



A solution of histidine has all acid-base groups protonated. The pKa of the carboxy group is 1.8, the amino group is 9.2, and the side chain is 6.0. For each mole of amino acid, 2.5 moles of NaOH are added. What is the pH of the solution?

Homework Equations



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

The Attempt at a Solution



I am trying to answer this question for a practice exam. I know that the base will deprotonate the protein, starting from the carboxy group followed by the side chain. I assume that the NaOH is very basic, so it will definitely cause a pH of greater than 6.0. How do I proceed from there?

My attempt: pH = unknown
pKa = 1.8, 6.0, 9.2
[A-]/[HA] = 2.5

pH = 9.2 + log(2.5) = 9.59, but this is wrong.

Could someone please help me with this problem. Thank you very much.
 
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  • #2
Honestly - no idea how to solve the question. The most obvious approach will be to calculate initial pH and continue from there, just calculating stoichiometry and NaOH excess. However...

A solution of histidine has all acid-base groups protonated.

This is not a quantitative statement. All groups are protonated - no matter wht pH is, it is just a matter of where the equilibrium lies. At pH 10 there is one caroboxylic group protonated for every 108 molecules, at pH 0.8 there is one NOT protonated carboxylic group for evey nine protonated.

Could be they want us to assume molecule like H3Histidine3+ that reacts with excess base. That's wrong, as such protonated molecule will exist only in the presence of huge excess of strong acid, which can't be ignored when calculating final concentration of the base.

So - either I have not yet waken up, or the question has no solution as worded.
 
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  • #3
According to my professor, adding alkali neutralizes protons so that...
Adding 0.5 mol of alkali will give a pH of the carboxyl group pKa = 1.8
Adding 1.0 mol of alkali will give a pH between 1.8 and 6.0
Adding 1.5 mol of alkali will give a pH of the histidine group side chain pKa = 6.0
Adding 2.0 mol of alkali will give a pH between 6.0 and 9.2
Adding 2.5 mol of alkali will give a pH of the amino group pKa = 9.2
Adding 3 mol of alkali will fully deprotonate the histidine.

I'm confused as to how he go the 0.5 mol increments of alkali.
 
  • #4
Take a look at the Henderson-Hasselbalch equation - at pH=pKa acid is exactly 50% neutralized ([HA]=[A-]).

Trick is - while he is right about neutralization stoichiometry, he still has no idea what the final pH is.
 
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  • #5
Treat the hisitidine exactly as you would treat a phosphoric acid buffer problem (H3PO4). Histidine can be though of as a triprotic acid in this example.

Have you done any problems using phosphoric acid in a like manner? For example, what would the pH be for a phosphoric acid solution that has had 2.5 equivalents of NaOH added? The pKa's are different but not the approach.

Borek, I know you know how to do this problem!
 
  • #6
chemisttree said:
Have you done any problems using phosphoric acid in a like manner?

Borek, I know you know how to do this problem!

What I don't like about the question, is the wording.

Would you say that phosphoric acid in its solution is fully protonated? 0.1M solution has pH around 1.6 - that means around 25% is in the form of H2PO4-. Perhaps that's nitpicking, but I still don't like it.
 
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1. What is protein pH titration?

Protein pH titration, also known as Henderson-Hasselbalch titration, is a technique used to determine the isoelectric point (pI) of a protein. It involves measuring the pH at which a protein has a net charge of zero, indicating equal numbers of positive and negative charges.

2. How is protein pH titration performed?

To perform protein pH titration, a solution of the protein is titrated with an acid or base, while the pH is measured at each step. The titrant is added in small increments and the pH is monitored until a plateau is reached, indicating the pI of the protein.

3. What is the Henderson-Hasselbalch equation?

The Henderson-Hasselbalch equation is used to calculate the pH of a solution containing a weak acid and its conjugate base. It is expressed as pH = pKa + log ([A-]/[HA]), where pKa is the acid dissociation constant, [A-] is the concentration of the conjugate base, and [HA] is the concentration of the weak acid.

4. What factors can affect protein pH titration?

The accuracy and reproducibility of protein pH titration can be affected by factors such as temperature, ionic strength, and the presence of other molecules in the solution. It is important to control these variables in order to obtain reliable results.

5. What is the significance of protein pH titration in biochemistry?

Protein pH titration is an important technique in biochemistry as it allows for the determination of the pI of a protein, which is a crucial parameter in protein characterization. The pI can affect the stability, solubility, and function of a protein, making it a key factor in understanding its properties and behavior.

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