Upset Tummy? MOM to the Rescue!


Mix milk of magnesia (MOM) with universal indicator and observe the dramatic rainbow of colors as the antacid dissolves in the simulated stomach acid! This is a great demonstration to teach concepts of acids and bases, solubility, Ksp and “antacid testing” consumer chemistry.


  • Acid–base neutralization
  • Solubility and Ksp
  • Antacids


Hydrochloric acid, HCl, 3 M, approximately 20 mL*
Milk of magnesia, 20 mL*
Universal indicator solution, 1%, 4–5 mL*
Water, distilled or deionized, 800 mL
Beaker, 1-L (or other large beaker)
Graduated cylinder, 25-mL or 50-mL
Ice, crushed (or ice cubes)
Magnetic stir bar
Magnetic stir plate (or stirring rod)
Pipets, Beral-type, 2*
*Materials included in kit.

Safety Precautions

Milk of magnesia is intended for laboratory use only; it has been stored with other non–food-grade laboratory chemicals and is not meant for human consumption. Hydrochloric acid solution is toxic by ingestion and inhalation and is corrosive to skin and eyes. Universal indicator solution is an alcohol-based flammable solution. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please review current Safety Data Sheets for additional safety, handling and disposal information.


Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. The final solution may be neutralized according to Flinn Suggested Disposal Method #24b. Excess milk of magnesia may be placed in the trash according to Flinn Suggested Disposal Method #26a.


  1. Measure 20 mL of milk of magnesia using a graduated cylinder and pour it into a 1-L beaker.
  2. Place the 1-L beaker on a magnetic stir plate. Add a magnetic stir bar to the beaker.
  3. Add water and crushed ice (or ice cubes) to give a total volume of approximately 800 mL. Turn on the stir plate so as to create a vortex in the mixture.
  4. Add about 4–5 mL (about 2 pipets full) of universal indicator solution. Watch as the white suspension of milk of magnesia turns to a deep purple color. The color indicates that the solution is basic.
  5. Add 2–3 mL (1 pipet full) of 3 M HCl. The mixture quickly turns red and then goes through the entire universal indicator color range back to purple.
  6. Repeat this process, adding HCl one pipet full at a time, waiting after each addition until the mixture turns back to blue–purple.
  7. The process can be repeated a number of times before all of the Mg(OH)2 has dissolved and has reacted with the HCl. As more acid is added, the color changes will occur more rapidly and eventually the suspension will be completely dissolved. This will be evidenced by a clear, red solution.

Student Worksheet PDF


Teacher Tips

  • This kit contains enough chemicals to perform the demonstration seven times: 150 mL of milk of magnesia, 250 mL of 3-M hydrochloric acid solution, 50 mL of universal indicator solution and 15 Beral-type pipets.
  • If a 1-L beaker is not available, use a 600-mL or 400-mL beaker. Adjust chemical amounts accordingly. (Note: The actual milk of magnesia concentration does not have to be exact in order for the demonstration to work.)
  • If a magnetic stir plate is not available, the mixture can be stirred with a stirring rod.
  • The acid used in the demonstration is 3 M hydrochloric acid (HCl). Actual stomach acid ranges in pH from 0.1 to 1.0 M HCl. However, 3 M HCl is used in this demonstration in order to limit the total acid volume and allow the reaction to go to completion with a reasonable volume of acid. If desired, dilute the 3 M acid to 1 M and perform the experiment as written. The volume of acid needed will be three times greater.
  • The reaction is performed on ice in order to slow down the color changes so that all colors in the universal indicator color range can be viewed. The reaction may be performed without the use of ice.
  • Consider performing this demonstration at different temperatures—5 °C, 25 °C and 60 °C—to compare the effect of temperature on the rate of reaction.
  • An excellent follow-up to this antacid demonstration is Flinn’s Antacid Testing Lab Kit—How Powerful Is Your Antacid? (Catalog No. AP1741).

Correlation to Next Generation Science Standards (NGSS)

Science & Engineering Practices

Analyzing and interpreting data
Constructing explanations and designing solutions

Disciplinary Core Ideas

MS-PS1.A: Structure and Properties of Matter
MS-PS1.B: Chemical Reactions
HS-PS1.B: Chemical Reactions

Crosscutting Concepts


Performance Expectations

MS-PS1-2. Analyze and interpret data on the properties of substances before and after the substances interact to determine if a chemical reaction has occurred.
HS-PS1-2. Construct and revise an explanation for the outcome of a simple chemical reaction based on the outermost electron states of atoms, trends in the periodic table, and knowledge of the patterns of chemical properties.

Answers to Questions

  1. Describe what happened in this demonstration.

Universal indicator was added to a mixture of milk of magnesia and ice water, turning the solution purple, indicating a pH of around 10. Hydrochloric acid was added one pipet-full at a time. Each time the solution flashed red before going through a series of color changes, from red to orange to yellow to green to blue to purple again. This process became more and more rapid, until finally the milky suspension turned into a clear, red solution.

  1. Write the balanced chemical equation for each of the following reactions.

a. Neutralization reaction between magnesium hydroxide and hydrochloric acid

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

b. Dissociation of magnesium hydroxide

Mg(OH)2(s) → Mg2+(aq) + 2OH(aq)

c. Reaction between hydrogen ions from the acid and hydroxide ions from the base

H+(aq) + OH(aq) → H2O(l)

  1. Using Le Chatelier’s Principle, explain why adding hydrochloric acid causes more magnesium hydroxide to dissolve in solution.

The reaction in Equation c uses up hydroxide ions from Equation b. To reestablish equilibrium Equation b therefore shifts to the right, causing more magnesium hydroxide to dissolve and react with the acid. As more acid is added, more magnesium hydroxide dissolves until eventually there is none left.

  1. Explain why the solution is red and clear at the end of the demonstration.

All the solid magnesium hydroxide, which was responsible for the milky appearance of the solution, has dissolved, resulting in the solution’s clarity. The red color is due to the universal indicator, which is red in the presence of an acid, in this case the excess hydrochloric acid.


The active ingredient in milk of magnesia is magnesium hydroxide, Mg(OH)2. Magnesium hydroxide forms a suspension in water since it has a very low solubility—0.0009 g/100 mL in cold water and 0.004 g/100 mL in hot water.

Initially in the demonstration, the solution is basic due to the small amount of Mg(OH)2 that goes into solution. The universal indicator gives the entire solution a violet color, indicating a pH of about 10. (See Universal Indicator Color Chart in Table 1.) Upon adding hydrochloric acid (the simulated “stomach acid”), the mixture quickly turns red because the acid disperses throughout the beaker, first neutralizing the small amount of dissolved Mg(OH)2, and then turning the solution acidic from the excess acid that is present.


The excess acid causes more Mg(OH)2 from the suspension to gradually dissolve. As more of the Mg(OH)2 goes into solution, the acid is neutralized and eventually the solution becomes basic again from the excess Mg(OH)2 that is present. The addition of universal indicator allows this process to be observed. During the process, the color of the mixture goes through the entire universal indicator color range—from red to orange to yellow to green to blue and finally back to violet. By adding more “stomach acid,” the process can be repeated several times before all of the Mg(OH)2 is dissolved and eventually neutralized.

Magnesium hydroxide is classified as a weak base (in most textbooks) due to its very limited solubility in water. This limited solubility makes it an ideal compound to use in commercial antacids since it slowly dissolves as it neutralizes stomach acid rather than dissolving all at once. The neutralization reaction is the reaction between Mg(OH)2 (a weak base) and HCl (a strong acid). The overall equation for the reaction is shown in Equation 1.

The reacting species for the strong acid, HCl, is the hydrogen ion, H+. In contrast, since Mg(OH)2 is a weak base, the principal reacting species is the undissociated Mg(OH)2 molecule. The acid–base reaction involves the Mg(OH)2 molecule and the H+ ion as reactants. The products are a H2O molecule and a Mg2+ ion in solution. Because the chloride ion, Cl, from HCl is a spectator ion, it is not included in the net ionic equation.

While Mg(OH)2 is practically insoluble, a certain amount of Mg(OH)2 dissociates into ions when put in water. The extent of dissociation of Mg(OH)2 is indicated by its solubility product constant, Ksp. The Ksp at 25 °C for Mg(OH)2 is 6 x 10–12, indicating that only a small amount of Mg(OH)2 dissociates into ions and the reaction equilibrium lies far to the left, according to Equation 2.

In the demonstration, the initial milk of magnesia suspension in water contains very few Mg2+ and OH ions before the acid is added. As HCl is added to the beaker containing milk of magnesia, the H+ ions from the HCl react with the OH ions (those that are actually in solution from the Mg(OH)2) according to Equation 3.

The reaction between H+ (stomach acid) and OH (antacid) to form water uses up some of the OH– ions and drives Equation 2 to the right, causing more Mg(OH)2 to dissolve and dissociate into ions. As OH– ions are removed from solution by the H+ ions, more and more Mg(OH)2 is forced to dissociate to replace those ions, according to Le Chatelier’s Principle. As more acid is added, the Mg(OH)2 continues to dissociate until all of it is dissolved and Equation 2 lies all the way to the right. The final solution in the milk of magnesia demonstration will thus be clear and acidic (red in color from the universal indicator), indicating that the Mg(OH)2 is fully dissolved. At this point, the “antacid power” or “acid-neutralizing ability” of the milk of magnesia is depleted.


Special thanks to Bette Bridges, Bridewater-Raynham High School, Bridgewater, MA; Annis Hapkiewicz, Okemos High School, Okemos, MI; and Penney Sconzo, Westminster School, Atlanta, GA, for separately bringing this demonstration to our attention.

Summerlin, L. R.; Borgford, C. L.; Ealy, J. B. Chemical Demonstrations: A Sourcebook for Teachers, Vol. 2; American Chemical Society: Washington, DC. 1988; p 173.

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