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Acids, Bases, and Buffers 1: Monoprotic and Polyprotic Acids |
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Prelab Problems
1. Bottles of vinegar are labelled, " 5% acetic acid, by mass." Calculate the mass of acetic acid in the 10.0 mL of vinegar in the above example, assuming that the solution has the same density as water at room temperature. Convert the label concentration in % mass units to molarity and compare to the result of the titration example on the Chemistry and Background page.
2. Write the balanced chemical equation for the reaction of acetic acid and sodium hydroxide. If 50.0 mL of 0.131 M CH3COOH is titrated with 0.203 M NaOH, how many mL of base must be added to reach the equivalence point? What will be the pH and mL values of the half-equivalence point?
3. Describe how to make 100 mL of buffer solution from 10.0 mL of 1.5 M acetic acid and an equal number of moles of sodium acetate. Sodium acetate will be supplied as solid sodium acetate trihydrate (MW = 136.080), so calculate the number of grams needed. What will be the pH of your buffer? How would you make 100 mL of a buffer of pH 5, starting from 10.0 mL of 1.5 M acetic acid and a supply of solid sodium acetate trihydrate? Equation 6 for acetic acid and the pKa values in Table 1 will be useful here. You can also use the molecular mass calculator in the periodic table applet, elsewhere on this website.
4. 50 mL of the buffer solution prepared in Problem 3 is titrated with 0.12 M NaOH and requires 46.18 mL of solution to change pH from 4.75 to 5.75. Calculate the solution's basic buffer capacity in the units defined in the introduction. Would you expect the acidic buffer capacity to be greater or less than the basic buffer capacity for this solution? Why?
5. Consider the titration of 100.0 mL of 0.0200 M H3PO4 with 0.121 M NaOH. Calculate the milliliters of base that must be added to reach the first, second, and third equivalence points.
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