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Homework Help: Using polyprotic acid as a titrant

  1. Aug 7, 2010 #1
    1. The problem statement, all variables and given/known data

    I'm currently doing an experiment that requires me to use a weak polyprotic acid (triprotic) to neutralize a strong base. I'll be using the acid as the titrant. I need to find out the molarity of the base. Can I still use the M1V1=M2V2 equation in this case?

    2. The attempt at a solution
    Since it's a triprotic acid , should I divide the number of moles of the base that I have calculated in the by 3?
  2. jcsd
  3. Aug 7, 2010 #2


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    Staff: Mentor

    You are on the right track, but it is more complicated and depends on the acid and the end point pH.

  4. Aug 8, 2010 #3
    1.Thanks a lot for your response . I've been trying to read as much as I can on the subject of polyprotic acids, but I am still confused.

    2.Information on the polyprotic acid I will be using
    The acid I am using now is citric acid ( H3C6H5O7) which has pKa values of 3.13, 4.76, and 5.41 . The end point of citric acid is pH 9.37.

    The ionization of citric acid is such:
    H3(C6H5O7) --->C6H5O73- + 3H+

    Actually the substance I'm going to titrate it with has a pH of 5.4. It is composed of some alkaloids and other materials. I'm trying to titrate the alkaloids with the citric acid because my final aim is to get neutralize the alkaloids.

    3.A little of my working

    I was thinking of using the pH indicator Bromocresol purple which has a transition pH of 5.2–6.8 or methyl red with a transition pH of 4.5–5.2.

    How would the hydrogen atoms react to the OH- in the analyte. Since this time the acid is the titrant, would the acid use up all first hydrogen ion before moving to the next one? or with each drop of acid into the analyte all 3 hydrogen ions would be used up?

    I am thinking that the first hydrogen is used up first until it reaches a certain pH before the next hydrogen is dissociated. Is that true? If so, what are the equations that I need to know?
    Last edited: Aug 8, 2010
  5. Aug 8, 2010 #4


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    Staff: Mentor

    You mean pH of the end point when citric acid is titrated with a strong base.

    I don't get what you are saying. Solution of alkaloids should be - as the name implies - a little bit alkaline, 5.4 is on the wrong side of 7. And if you have a solution with pH 5.4 you can't neutralize it with acid, you need a base for that.

  6. Aug 8, 2010 #5
    Sorry I do mean the end point of citric acid when it is titrated with strong base. Thanks.

    The analyte is an extract of a vegetable. It contains alkaloids which I hope the citric acid would interact with. Initally when we tested the pH of the vegetable extract we were quite surprised that the pH was less than pH7 since it was supposed to contain quite a bit of alkaloids. But then we found out that in the extract there is also ascorbic acid and several other compounds.

    My lecturer requires me to do addition titration of the citric acid to alter the taste of the vegetable extract. Can the same concepts of an acid-base titration still be used?
  7. Aug 9, 2010 #6


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    Staff: Mentor

    Basic concepts are universal, but how they should be applied is not clear to me, as I am not sure I know what you are going to do.

    Are you sure citric acid is going to be your titrant? What is the procedure for addition titration?

  8. Aug 14, 2010 #7
    The pH is acidic because within the extract there are other compounds as well. I'll try to understand the experiment a little better and try to figure it out. Thanks for all your help.
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