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2002 AP

2002 AP. Calculate the value [H+] in an HOBr solution that has a pH of 4.95. Calculate [H+]. Given the pH is 4.95 Use the exponent 10 raised the –pH This will yield the [H+].

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2002 AP

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  1. 2002 AP

  2. Calculate the value [H+] in an HOBr solution that has a pH of 4.95

  3. Calculate [H+] • Given the pH is 4.95 • Use the exponent 10 raised the –pH • This will yield the [H+]

  4. Write the equilibrium constant expression for the ionization of HOBr in an HOBr solution with a [H+]= 1.8 x 10^(-5)

  5. Equilibrium constant expression • Given the chemical equilibrium reaction and the concentration of [H+] • [H+] = [OBr-]: This is true because as HOBr dissociates the products form in same proportions • Exclude x because it is negligible. • Use the Law of Mass action Products over Reactants raised to their stoichiometric coefficients.

  6. Equilibrium constant expression • Use the equilibrium constant. • The product of the concentration divided by the Ka yields the concentration of HOBr.

  7. Calculate the volume of 0.115M Ba(OH)2 needed to reach equivalence when titrated into 65mL sample of 0.146M of HOBr

  8. Calculate Volume Needed to Reach Equivalent Point • Convert the Volume of HOBr to liters • Multiply the volume in Liters by the Molarity of HOBr in Solution • This will give you the moles of the HOBr in solution • The titrant used is Ba(OH)2 (Strong Base) so the concentration of OH- must be doubled therefore you must multiply the concentration of Ba(OH)2 to get the concentration of OH-

  9. Calculate Volume Needed to Reach Equivalent Point • Take the Moles of HOBr (HOBr = OH-) that was calculated and divide it by the molarity of OH- ([Ba(OH)2] x 2). • This will give you the volume in liters necessary to added to reach eq point

  10. Indicate whether the pH at equivalence point is less than 7, 7, or greater than seven. Explain

  11. What will the pH at the eq point be? • You can go through the calculations to determine the pH specifically at the eq point. • The basic rule of thumb is that if you are titrating a weak acid with a strong base then the pH at the end point will be greater than 7.

  12. Calculate the number of the moles of NaOBr(s) that would have to be added to 125mL of 0.160M HOBr to produce a buffer solution with a [H+] of 5.00 x 10^(-9)

  13. Number of Moles needed to produce certain Concentration • The Henderson-Hasselbauch equation can be used for this scenario (See equation Below) • Given the [H+] take the –log([H+]) and determine the pH for the equation

  14. Number of Moles needed to produce certain Concentration • You have already been given the Ka so plug that in the equation as well • Separate the ratio of concentrations into log(B) – log (A) (Refer to Log Rules)

  15. Number of Moles needed to produce certain Concentration • Solve for the Log (Base) and then eliminate the log function by raising the solution using the base of 10. • This will yield your concentration of Base

  16. Number of Moles needed to produce certain Concentration • Multiply the concentration by the volume of solution in liters • This will yield the number of NaOBr moles

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