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17.E: Exercises

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    17.1: Common-Ion Effect in Acid-Base Equilibria


    1. The solubility product Ksp for bismuth sulfide \(\ce{Bi2S3}\) is \(1.6 \times 10^{-72}\) at 25 °C. What is the molar solubility of bismuth sulfide in a solution that is 0.0010 M in sodium sulfide \(\ce{Na2S}\)?
    2. John poured 1.0 mL of 0.10 M \(\ce{NaCl}\), 1.0 mL of 0.10 M \(\ce{KOH}\), and 1.0 mL 0.20 \(\ce{HCl}\) solutions together and then he made the total volume to be 100.0 mL. What is the \(\ce{[Cl- ]}\) in the final solution?
    3. The Ksp for \(\ce{AgCl}\) is \(1.0 \times 10^{-10}\). From which of the following solutions would silver chloride precipitate?
      A: A solution 0.10 M in \(\ce{Ag+}\) and 1.00 M in \(\ce{Cl-}\)
      B: A solution \(1.0\times 10^{-5}\; M\) in \(\ce{Ag+}\) and 0.20 M in \(\ce{Cl-}\)
      C: A solution \(1.0 \times 10^{-7}\; M\) in \(\ce{Ag+}\) and 1.0E-7 M in \(\ce{Cl-}\)
      1. A only
      2. B only
      3. C only
      4. A and B
      5. A, B, and C
      Substance Ksp
      magnesium hydroxide \(1.2\times 10^{-11}\)
      magnesium carbonate \(1.6\times 10^{-5}\)
      magnesium fluoride \(6.4\times 10^{-9}\)
    4. Addition of which of the following substances will cause the precipitation of a salt from one liter of a \(1\times 10^{-4}\; M\) \(\ce{Mg^2+}\) solution?
      1. \(1\times 10^{-4}\) mole NaOH
      2. \(1\times 10^{-1}\) mole nitric acid
      3. \(1\times 10^{-5}\) mole potassium acetate
      4. \(1\times 10^{-4}\) mole ammonium nitrate
      5. \(1\times 10^{-2}\) mole sodium fluoride
    5. The Ksp for strontium chromate is \(3.6 \times 10^{-5}\) and Ksp for barium chromate is \(1.2\times 10^{-10}\). What concentration of potassium chromate will precipitate the maximum amount of either the barium or the strontium chromate from an equimolar 0.10 M solution of barium and strontium ions without precipitating the other?
    6. Iron(II) hydroxide is only sparingly soluble in water at 25 °C; its Ksp is equal to \(7.9\times 10^{-16}\). Calculate the solubility of iron(II) hydroxide in a solution of pH 6.0.


    1. Answer 1.08e-13
      x = molar solubility; (2x)2 (0.0010 + 3x)3 = Ksp; x = ?
    2. Answer 0.0030 M
      Evaluate \(\ce{[Na+]}\), \(\ce{[K+]}\), \(\ce{[H+]}\) and \(\ce{[OH- ]}\) for fun!
    3. Answer d.
      Calculate the \(\ce{[Ag+] [Cl- ]}\) for A, B, and C, and compare their values with Ksp.
    4. Answer e.
      A lot of calculations to figure out, but it's a common ion problem.
    5. Answer 3.6e-4
      This problem requires some thinking.
    6. Answer 7.9 M
      \(\ce{[OH- ]}\) = 10(-14+6) = 1e-8. x * (1e-8 + x)2 = Ksp; x = ?

    17.3: Acid-Base Indicators


    There are numerous natural indicators present in plants. The dye in red cabbage, the purple color of grapes, even the color of some flowers are some examples. What is the cause for some fruits to change color when they ripen?


    \(\ce{[H+]}\) of the juice changes. The changes in pH or \(\ce{[H+]}\) cause the dye to change color if their conjugate acid-base pairs have different colors. There may be other reasons too. Do colors indicate how good or bad they taste?


    Choose the true statement:

    1. All weak acids are indicators.
    2. All weak bases are indicators.
    3. Weak acids and bases are indicators.
    4. All indicators are weak acids.
    5. An acid-base conjugate pair has different colors.
    6. Any indicator changes color when the pH of its solution is 7.


    d. Color change is a requirement for indicators.


    Do all indicators change color at pH 7 (y/n)?


    No! Phenolphthalein changes color at pH ~9. Bromothymol blue has a pKn value of 7.1. At pH 7, its color changes from yellow to blue. Some indicators change color at pH other than 7.

    17.E: Exercises is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts.

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