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March 29, 2017

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Solve an equilibrium problem (using an ICE table) to calculate the pH of each solution.

a solution that is 0.266 M in CH3NH2 and 0.134 M in CH3NH3Br.

Kb = 4.4e-4

  • chemistry 1046 - ,

    It's easier to solve this one with the Henderson-Hasselbalch equation BUT if you must use the ICE chart, here it is.

    ......CH3NH2 + HOH ==> CH3NH3^+ + OH^-
    I.....0.266.............0..........0
    C........-x.............x..........x
    E.....0.266-x............x.........x

    .......CH3NH3Br ==> CH3NH3^+ + Br^-
    I.......0.134.........0..........0
    C.......-0.134.......0.134....0.134
    E..........0.........0.134.....0.134

    Kb = 4.4E-4 = (CH3NH3^+)(OH^-)/(CH3NH2)
    Substitute as follows:
    (CH3NH3^+) = 0.134 +x = 0.134 from salt and x from base.
    (OH^-) = x
    (CH3NH2) = 0.266-x from salt
    Solve for x and convert to pH.

    If using the HH equation first convert Kb to pKb = -log 4.4E-4 = 3.36, then
    pH = pKa + log(base)/acid)
    pH = 10.64 + log(0.134/0.266) = ?
    Same answer either way.

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