ANTACIDS
Antacids are over-the-counter formulations that are used to treat indigestion caused by
production of an excess of stomach acid (hydrochloric acid). Manufacturers of antacids make
extravagant claims about the abilities of their products to neutralize stomach acid. In this
laboratory experiment you will determine the relative efficiencies of two antacids to neutralize
stomach acids. You will find the number of grams of acid neutralized per gram of antacid for
both Tums® and Rolaids®.
In order to conduct this experiment, you will need to prepare a standard sodium hydroxide
solution (approximately 0.15M) and a standard hydrochloric acid solution (approximately 0.1M).
The hydrochloric acid solution will be made to simulate stomach acid. The solutions will be
standardized by doing acid-base titrations. The most important principle to be aware of when
doing acid-base titrations is that the number of moles of hydrogen ion from the acid is equal to
the number of moles of hydroxide from the base at the equivalence point (or endpoint). If the
acid or base is in a solution, the number of moles is equal to the volume of the solution times its
molarity. If the acid or base is a solid, the number of moles is equal to its mass divided by the
molar mass.
The sodium hydroxide (NaOH) solution is standardized by titrating it with the primary
standard, potassium hydrogen phthalate (KHP). The neutralization reaction between NaOH and
KHP is given below.
KHP + NaOH → NaKP + H2O
The molarity of the NaOH solution is calculated by first dividing the mass of KHP by its
molar mass. This gives the number of moles of KHP which is equal to the number of moles of
NaOH. The resulting quantity is then divided by the number of liters of NaOH solution titrated
to get its molarity.
The hydrochloric acid solution is standardized by titrating it with the standardized sodium
hydroxide solution. The molarity of the HCl solution is calculated by first multiplying the
molarity of the NaOH solution times the number of liters of this solution used in the titration.
This gives the number of moles of NaOH and is the same as the number of moles of HCl. This
resulting quantity is then divided by the number of liters of HCl solution used, to get the
molarity.
Tums® contains calcium carbonate as an active ingredient. The carbonate reacts with
hydrogen ion from hydrochloric acid by the chemical reaction
CO32- + 2H+ → CO2 + H2O
Notice the two moles of hydrogen ions are neutralized by one mole of carbonate. There are two
active ingredients in Rolaids®. They are calcium carbonate and magnesium hydroxide. In
addition to the reaction of carbonate with hydrogen ion, the hydroxide ion reacts with hydrogen
ion by the chemical reaction
OH- + H+ → H2O
In this case, one mole of hydrogen ions reacts with one mole of hydroxide.
In order to determine the relative efficiency of an antacid, a solution containing a massed
sample of the antacid is titrated past the endpoint with the standardized HCl solution. Then the
mixture is back titrated with the standard NaOH solution to a pink endpoint. The number of
moles of HCl and number of moles of NaOH titrated are calculated by multiplying their
molarities by the number of liters of solution. The number of moles of HCl needed to react with
the antacid equals the number of moles of HCl used minus the number of moles of NaOH back
titrated. The grams of HCl used is equal to the moles of HCl times its molar mass, and this
quantity divided by the grams of antacid is the relative efficiency.
PROCEDURE:
Note: Procedure will be completed over two laboratory periods and will be performed in groups
of two!!
Part 1: Standardization of Hydrochloric Acid and Sodium Hydroxide Solutions (to be
completed in first laboratory period)

As a group:
1. Measure out 50 mL of 1.0 M HCl in your 100 mL beaker. Pour this solution into
a 600 mL beaker (located at the bench tops, one beaker per group). Add enough
distilled water so that the liquid level is equal to 500 mL. Distribute half of this
solution to each student’s 250 mL plastic bottle with a lid. Label each bottle as
HCl.
2. Measure out 50 mL of 1.5 M NaOH in your 100 mL beaker. Pour this solution
into a 600 mL beaker (located at the bench tops, one beaker per group). Add
enough distilled water so that the liquid level is equal to 500 mL. Distribute half
of this solution to each student’s 250 mL plastic bottle with a lid. Label each
bottle as NaOH.
One student will titrate KHP while the other student will titrate HCl.

Student A:
1. Weigh out between 0.7 and 0.9 grams of potassium acid phthalate (KHP) into a
weighing boat.
2. Transfer the KHP quantitatively to a 250 mL Erlenmeyer flask by using a squirt
bottle of distilled water to wash all of it out of the weighing boat and into the
flask.
3. Add about 50 mL of distilled water to the flask to dissolve all of the KHP. Then
add three drops of phenolphthalein indicator to the solution.
4. Fill your buret to about the 40 mL mark with distilled water. Let a small portion
run through the tip. Slowly pour the rest out the top of the buret as you rotate it.
Repeat this procedure with the NaOH solution you made in step 2. Fill your buret
slightly above the zero mark with your NaOH solution. Let a small portion run
through the tip to ensure are no bubbles. (Air bubbles can be seen as a gap in the
tip). Record the initial volume of NaOH on your data sheet.
5. Titrate the KHP solution with the NaOH solution by slowly adding base from the
buret and then swirling the flask. Initially the pink color caused by adding base
will swirl away. Continue the titration until the color changes to pink by addition
of one drop of base solution and stays pink (the endpoint). The goal is to achieve
as light of a pink color as possible.
6. Record the volume of the base solution needed to reach the endpoint and calculate
the concentration of the NaOH solution.
7. Do a second trial.

Student B:
1. Fill your buret to about the 40 mL mark with distilled water. Let a small protion
run through the tip. Slowly pour the rest out the top of the buret as you rotate it.
Repeat this procedure with the NaOH solution you made in step 2. Fill your buret
slightly above the zero mark with your NaOH solution. Let a small portion run
through the tip to ensure are no bubbles. (Air bubbles can be seen as a gap in the
tip). Record the initial volume of NaOH on your data sheet.
2. Rinse your pipet first with distilled water and then with your HCl solution.
3. Transfer 50 mL of your HCl solution (made in step 1) into a clean 100 mL beaker.
4. Pipet 20 mL of this HCl solution from the beaker to a 250 mL Erlenmeyer flask
using a volumetric pipet using a volumetric pipet. Add about 50 mL of distilled
water to the flask and then add three drops of phenolphthalein.
5. Titrate the acid with NaOH solution from your buret to a pink endpoint. Initially
the pink color caused by adding base will swirl away. Continue the titration until
the color changes to pink by addition of one drop of base solution and stays pink
(the endpoint). The goal is to achieve as light of a pink color as possible.
6. Record the volume of the base solution needed to reach the endpoint (final NaOH
volume) and calculate the concentration of the HCl solution.
7. Do a second trial.

Each student must do the following prior to leaving:
1. Drain any leftover NaOH solution from your buret back into your plastic bottle.
Save your NaOH and HCl solutions for next week.
2. Rinse your buret and pipet with distilled water prior to returning.
3. Exchange data and observations with your partner. Calculations and the formal
report must be done individually.

Calculations to be included in the formal report:
1. Concentration of HCl and NaOH for each trial
2. Average concentration of HCl and NaOH
Part 2: Analysis of the Antacids (to be completed in second laboratory period)
1.
This lab is to be done with the same partner you worked with last week. One student
will titrate the Tums® tablet while the other is titrating the Rolaids® tablet.
2.
Crush a Tums® or Rolaids® antacid tablet in a mortar using a pestle.
3.
Weigh out about 0.2 gram of the antacid into a weigh boat and transfer quantitatively
to a 250 mL Erlenmeyer flask. Record the mass on your data sheet to the nearest
thousandth of a gram. Label this as Trial 1.
4.
Weigh, transfer, and record the weight of a second sample and label the flask as Trial
2.
5.
Fill your buret to about the 40 mL mark with distilled water. Let a small protion run
through the tip. Slowly pour the rest out the top of the buret as you rotate it. Repeat
this procedure with the NaOH solution you made in step 2. Fill your buret slightly
above the 25 mL mark with your NaOH solution. Let a small portion run through the
tip to ensure are no bubbles. (Air bubbles can be seen as a gap in the tip). Record the
initial volume of NaOH on your data sheet.
6.
Rinse your volumetric pipet first with distilled water and then with your HCl solution.
7.
Transfer 100 mL of your standardized HCl solution (made last week) into a clean 150
mL beaker.
8.
Pipet 25 mL of your standardized HCl solution from the beaker to a 250 mL
Erlenmeyer flask using a volumetric pipet. Swirl to dissolve as much of antacid as
possible. Part of the sample will not dissolve.
9.
Heat the resulting mixture in each of your Erlenmeyer flasks to near boiling on a hot
plate in order to free all of the active antacid ingredient from the inert portion of the
sample and drive off the carbon dioxide gas. When you begin to see bubbles forming
from the boiling of the mixture, remove it from the hot plate.
10.
Add three drops of phenolphthalein to the solution in the flask labeled trail 1.
11.
Slowly back titrate the standardized NaOH solution from the base buret to a pink
endpoint. Record the volume of the NaOH solution added to the nearest 0.01 mL.
12.
Repeat this procedure for the Erlenmeyer flask labeled trial 2.
13.
Any extra NaOH or HCl solution should be poured down the drain.
14.
Rinse your buret and pipet with distilled water prior to returning.
15.
Exchange data and observations with your partner prior to leaving lab. Calculations
and the formal report must be done individually.
16.
Calculations to be included in the formal report:
a. Moles of HCl added, moles of excess HCl, moles of HCl neutralized by the
antacid, mass of HCl neutralized by antacid
b. Relative efficiency 
mass of HCl neutralize d by antacid
mass of antacid
ANTACIDS Data sheet
Your Name_________________________ Date______________
Section___________
Your Partners Name____________________
Part 1: Standardization of Hydrochloric Acid and Sodium Hydroxide Solutions
Molar Mass of KHP = 204.23 g/mol
A. Standardization of NaOH
Trial 1
Trial 2
1. Mass of KHP (g)
____________
____________
2. Initial Volume of NaOH (mL)
____________
____________
3. Final Volume of NaOH (mL)
____________
____________
4. Volume of NaOH delivered (mL)
____________
____________
B. Standardization of HCl
Trial 1
5. Volume of HCl (mL)
Trial 2
20.00 mL
20.00 mL
6. Initial Volume of NaOH (mL)
____________
____________
7. Final Volume of NaOH (mL)
____________
____________
8. Volume of NaOH delivered (mL)
____________
____________
Trial 1
Trial 2
____________
____________
25.00 mL
25.00 mL
3. Initial Volume of NaOH (mL)
____________
____________
4. Final Volume of NaOH (mL)
____________
____________
5. Volume of NaOH delivered (mL)
____________
____________
Part 2: Analysis of the Antacids
Your Antacid________________
1. Mass of antacid (g)
2. Volume of HCl (mL)
Your Partner’s Antacid _________________________
Trial 1
6. Mass of antacid (g)
Trial 2
____________
____________
25.00 mL
25.00 mL
8. Initial Volume of NaOH (mL)
____________
____________
9. Final Volume of NaOH (mL)
____________
____________
10. Volume of NaOH delivered (mL)
____________
____________
7. Volume of HCl (mL)
Formal Report: Answer the following questions in your formal report. If the answer requires a
calculation, show your work.
1. Part I:
a. Average concentrations of NaOH and HCl
b. What is the relationship between the number of moles of hydrogen ions and the
number of moles of hydroxide ions present at the endpoint of a titration?
2. Part II:
a. Moles of HCl added
b. Moles of excess HCl = moles of NaOH titrated
c. Moles of HCl neutralized by antacid
d. Mass of HCl neutralized by antacid
e. Relative efficiency 
mass of HCl neutralize d by antacid
mass of antacid
f. Which antacid is more efficient at neutralizing stomach acid?
g. Why do we add excess acid and then back titrate with NaOH to the endpoint in
the titration of each antacid instead of titrating directly to the endpoint with acid?