Experiment 3
Evaluating Commercial Antacids
Lab Owl Announcement

    Upon completion of this lab go log onto OWL.  A Lab Owl section should now appear in your courses and your first assignment, Lab Owl: Exp 3, should appear in this section.  You have one week to complete this assignment.  Two more assignments will appear here as the semester progresses.  Remember, these Lab Owls are worth 25% of your laboratory grade.

Introduction

Reactions between Metal Carbonates and Acid

    Heart burn and acid indigestion are primarily due to excess acid in the stomach. There are many preparations sold over the counter to relieve these conditions, many of which act by neutralizing the excess acid in the stomach, and some may be more effective than others! In most of these products, other ingredients are added, flavoring agents and substances to soothe the walls and lining of the stomach, binders to make the tablet hold together and so forth. The active ingredients, however, are there to neutralize the excess acid in the stomach. 

    Among the most common active ingredients are: Mg(OH)2 (the active ingredient in milk of magnesia), NaHCO3 (sodium bicarbonate) and CaCO3 (calcium carbonate) and in some Al(OH)3 (aluminum hydroxide). In our continued foray into the world of stoichiometry we are going to evaluate some of commercially available antacids by determining the amount of HCl (stomach acid!) that one tablet will neutralize. The reaction between the metal hydroxides and the stomach acid is an acid base reaction very similar to that in the previous experiment:

Mg(OH)2(s) + 2 HCl(aq) = MgCl2(aq) + 2 H2O(l)

     Metal carbonates, which are the most common active component in antacids, neutralize the stomach acid by forming a salt, water and a gas. 

CaCO3(s) + 2 HCl(aq) = CaCl2(aq) + H2O(l) + CO2(g)

Back Titration

    It would seem possible to perform the analysis of the antacids in a similar fashion to that used in the last experiment, i.e. dissolve the antacid in water, add an indicator and titrate with a base whose concentration is known. However, there are prohibitive drawbacks to this: 

    1. The active ingredients are at best, only sparingly soluble in water. 

    2. It requires knowing what the active component is. 

    3. The technique would have to be modified for every different antacid. 

    4. Many of the antacids contain more than one active ingredient 

This is where the process of back titrating comes in. Dissolve the tablet in a relatively large quantity of acid (relative by virtue of the amount of acid that a tablet can neutralize). Let the active component/components neutralize the amount of acid that they are capable of, and then determining the amount of acid that remains using simple acid base chemistry: 

HCl(aq)remaining + NaOH(aq) = NaCl(aq) + H2O(l)

Evaluation of Commercial Antacids

    The tactic as indicated above is to dissolve the antacid in some volume of HCl whose concentration is known. From your knowledge of solutions you can thus determine the number of moles of HCl that was added. Some of this acid will then be neutralized by the antacid and the excess acid titrated with a sodium hydroxide solution whose concentration is known. The number of moles of sodium hydroxide added to neutralize the excess acid can be determined using: 

#mol = Molarity of the NaOH  (from the previous experiment.) x V(L) 

From the balanced chemical equation, this can be converted to moles of acid remaining. Now it is a case of subtraction to determine the moles of acid neutralized. 

In order to do a comparison between the antacids you will need to calculate: 

      1. The number of moles of acid neutralized per gram of tablet. 
           (The amount of acid neutralized divided by the mass of the tablet) 
      2. The number of moles of acid neutralized per $1 cost of the tablet. 
           (The amount of acid neutralized divided by the cost of the tablet in $)

Experimental Procedure:

Determining the ~Volume of HCl to Add

    Since the antacids vary to quite an extent in their effectiveness, your first goal is to determine the amount of HCl that you need to dissolve your tablet in.  The goal being to add sufficient HCl such that the amount remaining after the antacid has reacted requires 20-30mL of your 0.1M NaOH to neutralize it.

  1. Obtain three antacid tablets (same brand).  Note the brand name, number of tablets in a bottle and the cost per bottle.
     
  2. Record the mass of a tablet and place it in a clean dry Erlenmeyer flask.  Add ~ 25mL of distilled water and 15mL of ~1M HCl.  Make sure you note the exact molarity of the HCl.
     
  3. Add ~5 drops of phenolphthalein.  If the solution turns pink immediately add another 5mL of ~1M HCl.  Continue this until you obtain a solution that does not immediately turn pink upon the addition of phenolphthalein.
     
  4. Titrate (roughly) with your NaOH solution.  From the end point you will have to determine whether you need a little more or a little less HCl when you do the analysis on the remaining two tablets.

Evaluation of Commercial Antacids

  1. Record the mass of a tablet and place it in a clean dry Erlenmeyer flask.  Add ~ 25mL of distilled water and the volume  of ~1M HCl determined above.
     
  2. Add ~5 drops of phenolphthalein and titrate to a faint pink end point.
     
  3. Determine:
      1. The moles of acid neutralized by the tablet.
      2. The moles of acid neutralized per gram of tablet.
      3. The moles of acid neutralized per $1 cost of the tablet.
        Use the following data when determining this value:
        Rolaids 150 Tablets $4.49
        Tums 150 Tablets $4.79
        Top Care 150 Tablets $3.79
         
  4. Repeat this procedure with your second tablet.