Back Titration: Analysis of a Commercial Antacid



An excess of stomach acid (primarily HCl) causes heartburn and acid indigestion. Commercial antacids consist of a number of ingredients, such as binders and flavorings, but the active ingredient is simply a basic salt such as Mg(OH)2 (milk of magnesia), NaHCO3 (sodium bicarbonate), CaCO3 (calcium carbonate) or Al(OH)3 (aluminum hydroxide), to neutralize the acid. In this experiment we will analyze a typical antacid containing CaCO3 which reacts with an acid to form a salt, water and a gas:

CaCO3 + 2HCl à CaCl2 + H2O + CO2

A conventional acid/base titration is very difficult in this case as the active ingredient is only sparingly soluble in water. In order to overcome this limitation we will completely dissolve the tablets in excess acid and back titrate the remaining acid with NaOH to determine the amount of CaCO3 present. Of course, in order to obtain an accurate value for the concentration of acid titrated, we need to know the exact concentration and volume of NaOH used. The exact value can be measured with a buret but the exact concentration is more problematic. NaOH is hygroscopic (absorbs water) so that preparing a solution of accurately known concentration by weighing out solid NaOH is difficult. The prepared solution also tends to react with carbon dioxide in the air over time. We, therefore, have to perform a standardization of the NaOH before beginning the back titration. This is done by titrating the NaOH solution with a pure, stable acid (known as a primary acid). Potassium hydrogen phthalate KHP (MW = 204.22 g/mol) is often used for this purpose. An added advantage of this method is that it can be used to determine the neutralizing power of any antacid even if the active ingredient is not known.




  1. Obtain a 500mL volumetric flask and prepare approximately 0.6M solution of sodium hydroxide.
  2. Dissolve between 3 and 4 g of KHP in approximately 50mL of distilled water in a 250mL Erlenmeyer flask.
  3. Add 3 or 4 drops of phenolphthalein and titrate with your NaOH to a faint pink endpoint.
  4. Repeated the standardization procedure until you have 2 values for NaOH concentration that differ by less than 1%.

Analysis of Antacid Tablets

  1. Obtain an antacid tablet.
  2. Wash thoroughly a 250mL Erlenmeyer flasks with distilled water.
  3. Into the flask add the tablet and about 20mL of distilled water followed by 5.00mL of 6M HCl.
  4. Heat gently on a hotplate until all the effervescence has ceased. Boil for 1-2 minutes more. Some of the inactive tablet material may not dissolve, however, this should not interfere with the titration.
  1. Add 3 or 4 drops of phenolphthalein and titrate the cool solution with your standardized NaOH.
  2. Repeat with 3 more tablets.



Calculate the number of moles of NaOH used for each titration of a tablet, this is equivalent to the moles of excess HCl. Subtract this value from the initial number of moles of HCl added to the tablet to give the number of moles of acid neutralized. Calculate the number of moles of CaCO3 that this corresponds to and, therefore, the mass of CaCO3 in each tablet. Repeat for each trial and calculate the mean and standard deviation of your results.



  1. How did your mean value for the mass of CaCO3 per tablet compare to the value of 500mg given on the label?
  2. The label states that 500mg of CaCO3 correspond to 200mg of elemental. Show a simple calculation to confirm this statement.
  3. During the heating process some of the liquid will evaporate. Should we account for this change in our calculations? Explain.
  4. Some commercial antacids are colored. How might this be a problem in our analysis?
  5. Some antacids contain a mixture of active ingredients. If our antacid contained an equal number of moles of CaCO3 and Al(OH)3 what mass of CaCO3 must be present to neutralize the same amount of HCl as 500mg of CaCO3 alone? What can you say about the relative neutralizing power of these two ingredients?