# 3: Analysis of a Commercial Antacid (Experiment)


## Introduction

In this experiment, you will use the standard acid and standard base that you prepared earlier to analyze the neutralization capacity of a commercial antacid.

Safety Precautions:

Be especially careful when using the acid or base solutions as they can cause severe burns. All waste solutions may be disposed of by rinsing them down the drain.

## Procedure

1. Obtain an unknown antacid sample from the front counter. Using a mortar and pestle, grind the sample to a fine powder. To efficiently do this experiment you will first do a quick titration to get a rough idea of the neutralization capacity of your unknown. You will then do accurate titrations.

Question $$\PageIndex{1}$$

Write a balanced chemical equation for this reaction.

1. Remove the beaker from the heat source and cool to room temperature under running tap water, being careful not to contaminate the contents of the beaker. Just prior to titrating your sample, add about 5 drops of phenolphthalein indicator to the solution and titrate to the visual end point with your standard 0.1 M NaOH solution. Wash down the sides of the beaker with deionized water using your water bottle to be sure all the reactants are in the solution. The indicator will change colors from colorless to pink. This is the same as the first endpoint in the carbonate titration. For reasons of accuracy it is important that the volume of NaOH required be between 25 mL and 45 mL. If your volume is not within this range, you will have to adjust the mass of antacid sample for the next analyses. Using the information gained from this rough titration you are now ready to make an accurate determination.
3. Cool the flasks to room temperature under running tap water, being careful not to contaminate the contents. Titrate your sample, to the visual end point with your standard 0.1 M NaOH solution. The repeat titrations will go much faster once the endpoint is found.

## Calculations

1. Report the effective mass percent of $$\ce{OH^{–}}$$ or $$\ce{CO_{3}^{2–}}$$ for the antacid you measured. For remote students, just calculate it for $$\ce{CO_{3}^{2–}}$$. The effective mass is the mass of $$\ce{OH^{–}}$$ or $$\ce{CO_{3}^{2–}}$$ that would be in the unknown if that were the only base present. The antacid samples actually contain a variety of bases.
2. Report the average percent by mass of $$\ce{OH^{–}}$$ or $$\ce{CO_{3}^{2–}}$$ in your unknown along with a standard deviation, a 95% confidence limit, and a relative deviation. This can easily be done on a computer. or remote students, just calculate it for $$\ce{CO_{3}^{2–}}$$.

Question $$\PageIndex{2}$$

Do you have more confidence in your analysis of the $$\ce{KHP}$$ or of the antacid? Why?

Question $$\PageIndex{3}$$

Stomach acid is about 0.12 M $$\ce{HCl}$$. Does your unknown consume 47 times its own weight of 0.12 M $$\ce{HCl}$$ solution? Show by calculation.

Clean-up. After the experiment is completed, drain any remaining solution from the burette. Rinse each burette with deionized water. Then, fill the burette with deionized water and carefully place it in the burette rack.

Post-lab exercise. The data from your laboratory section will be collected. From this information, determine the average effective mass percent $$\ce{OH^{–}}$$ or $$\ce{CO_{3}^{2–}}$$ for each type of tablet along with the standard deviation, 95% confidence limit, and a relative deviation. Based on this data, what product would you recommend to the consumer? Explain.

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