Environmental Science 430

Fall 1999

Lab #2: Total Acidity

Purpose: To calculate the total acidity of solutions of simple strong acids, a solution of weak acids, and French Creek.

Materials (number of meters in parentheses):

• 0.01 M NaOH
• Phenolphthalein indicator
• pH meter
• Titration equipment
  • 50 ml burette burette stand and clamp
  • magnetic stirrer and stirring bar
  • 400 ml beaker
• 0.01 M HCl
• 0.01 M Acetic acid (CH3CO2H)
• solution of your choice such as coke or seltzer or tap water
• 0.01 M H2SO4
• A sample from French Creek

Procedure:

1. From the concentrated solution of NaOH prepare a 0.01 M solution of NaOH by dilution.
2. Measure the pH of 15 mL of your test solution (0.01 M HCl, acetic acid, solution of your choice, French Creek). If pH is above 4.0 add 5 ml increments of 0.01 M H2SO4 to reduce pH to 4.0 or less.
3. Measure sample pH. Add "standard" 0.1 M NaOH in incre-ments of 0.5 ml or less, such that a change of less than 0.2 pH units occurs per increment. After each addition, mix thoroughly but gently with a magnetic stirrer. Avoid splash-ing.
4. Record pH when a constant reading is obtained. Continue adding titrant and measure pH until pH 8.3 is reached. If you overshoot 8.3 back titrate by adding more H2SO4 and then trying again.
5. Construct the titration curve by plotting observed pH values versus cumulative milliliters titrant added. A smooth curve showing one or more inflections should be obtained. A ragged or erratic curve may indicate that equilibrium was not reached between successive alkali additions. In the titration of a single acidic species, such as HCL, the most accurate end point is obtained from the inflection point of a titration curve. The inflection point is the pH at which curvature changes from convex to concave or vice versa.
6. Acidity is normally reported as mg of CaCO3/liter. That is acidity is reported as the equivalent of the number of moles of hydrogen ions that could be neutralized by a quantity of CaCO3 expressed in mg. The formula weight of CaCO3 is 100 and every molecule of CaCO3 can neutralize two hydrogen ions.

   In addition acidity is pH dependent and might best be thought of as base neutralizing capacity (BNC). The amount of H+ consumed by titrating with base is necessarily dependent on the final pH you are titrating to. By convention pH 8.3 is a commonly used endpoint for titrating BNC, or acidity, of natural waters and agricultural soils. This, of course, is an operational definition of acidity and it is reported as follows:

   "The acidity to pH = ________mg CaCO3/L."

7. Acidity, as mg CaCO3/L = [(A X B) - (C X 2D)] X 50,000/mL sample where:

• A = mL NaOH titrant used
• B = molarity of NaOH
• C = mL H2SO4 used to back titrate, if necessary because you overshot the endpoint
• D = molarity of H2SO4 and
• 50,000 is the conversion from equivalents of H+ to mg of CaCO3.

Due: One week from the beginning of your lab section.

Reminders: Answers must be typed, and calculations presented, even when not asked for explicitly in the questions, and late papers will not be accepted.

SHOW ALL CALCULATIONS AND GRAPHS

Questions:

1. (10 points) Turn in graphs of your titration curves with the end-points, inflection points and the point you used to measure the pKa of acetic acid. How does your value for the pKa of acetic acid compare to the book value? Show all of your calculations.
2. (10 points) What is going on when you backtitrate? In other words, if you are supposed to be measuring the moles of H+ in solution by tabulating the ml of base added to your solu-tion, why doesn't the addi-tion of new acid to your solu-tion screw up your results? Could you backtitrate using 0.261 M H2SO4?
3. (10 points) How does the acidity (base neutralizing capacity) of the HCl and CH3CO2H compare? Explain the differences in pH and acidity you observed between the two acids.
4. (10 points) What controls the acidity of your test solution? French Creek?

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