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Very Difficult Chemistry Problem

  1. Apr 13, 2009 #1
    1. The problem statement, all variables and given/known data
    1. a) Calculate pH of solution that results form mixing 16.5 mL of 0.182 M HCN (aq) with 29.2 mL of 0.105 M NaCN. Ka of HCN is 4.9*10^-10

    b) and how could we determine what 'indicator' would be the 'best' to use for a titration between 0.10 M CH3NH2 with 0.10 M HBr?


    2. Relevant equations
    pH = Pka + log([A-]/[HA])
    ka=kw/kb

    3. The attempt at a solution
    a) pH = -log(4.9*10^-10) + log(.105/.182), which turns out wrong, the answer is actually 9.32(why???).

    b) found the kb of CH3NH2 to be 4.4*10^-4, thus ka is found from the given equation. anyway it gives something in the range of 10 or so, the answer is actually 4-6, why??? These questions have no likes in the textbook that my professor issued and the material is confusing. Please help.
     
  2. jcsd
  3. Apr 14, 2009 #2

    Borek

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    These are not concentrations in the final solution. You have not mixed 1:1 volumes (in which case at least ratio of the concentrations would be preserved).

    Hard to say not seeing what you did. However, you have a weak acid that - once dissolved - gives alkalic solution? Gimme a break :wink:

    Equivalence point calculation.

    Selecting indicators for acid-base titration.
     
  4. Apr 14, 2009 #3

    symbolipoint

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    itsmybedtime wrote:
    Borek responded:
    Borek, we are not sure what he meant by "10 or so"; itsmybedtime is confused and maybe is not clear on what some calculated values (possibly miscalculated) mean.

    itsmybedtime next wrote:
    If that is really true, then you really should check some alternative textbooks, either another General Chemistry book, or one or two analytical textbooks with the topic of Neutralization Titrations. Are you certain that your professor did not discuss the topic details of weak acid-base equilibria more thoroughly during class lectures? Some proressors do this to fill some inadequacies of the chosen course textbook.

    Some advice about studying weak acids & bases: Understanding their equilibria often requires studying the topic more than once; learning can be much/somewhat better the second time, even a third time. Also, develop some skill with quadratic equations and logarithms. They are essential PRACTICAL arithmetic for weak acids and bases.
     
  5. Apr 14, 2009 #4

    Borek

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    I can be wrong, but my understanding was that itsmybedtime calculated pH to be 10. As we deal with an acid solution, this result is wrong at first sight - one should expect pH below 7. This is just a reflex, you don't need to check details of the calculations to know that they have to be wrong in this case. That's what I was aiming at.
     
  6. Apr 14, 2009 #5

    symbolipoint

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    itsmybedtime asked about:
    One possible source of trouble is that CH3NH2 is a BASE, NOT an acid. Did you make a mistake in calculating Kb ? Do you know that most of your Kb calculations for this will involve [OH-] and that you might be justified in neglecting hydronium ions?

    In any case, if you are titrating a weak base using a strong acid titrant, and you know Kb for the base, can you predict about at what pH will be the equivalence point? This will give you some idea which indicator to choose.
     
  7. Apr 14, 2009 #6

    Borek

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    That's the correct approach - calculate Ka of conjugate acid from known Kb, use this Ka to calculate pH.

    At the equivalence point you have just a solution of CH3NH3+ - which is a weak, conjugate acid of CH3NH2. This acid determines pH of the solution.
     
  8. Apr 14, 2009 #7
    thanks for the insights everyone! I will be working on problems like this all night tonight. :smile:
     
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