Acid reflux is a common ailment that is frequently combated with the ingestion of antacids, compounds that will react with acid and neutralize it. Various chemicals are incorporated into antacids such as \( \ce {Mg(OH)2, NaHCO3, CaCO3} \) and other neutralizing bases. Each vendor will blend its own special recipe and sell the product at a given price, often claiming its product to be “the best” or “the most cost effective.”
Most antacids found on the market today are equally capable of neutralizing acid if taken in enough quantity. Each antacid has the common quality of acid neutralization for which the net ionic equation is shown in Reaction AN.1, an acid-base neutralization reaction. \begin {equation} \ce{\label {net_acid+base} H3O+(aq) + OH-(aq) -> 2 H2O(l)} \end {equation}
Note:
\( \ce {H3O+} \) is often abbreviated as \( \ce {H+} \). However, charged protons (\( \ce {H+} \)) do not exist in aqueous solution and, thus, \( \ce {H3O+} \) is the correct way to present the hydronium ion.
Titration is a method to determine the unknown concentration of a reagent by adding a known volume of a different reagent that reacts chemically with the first and for which the concentration is known.
In this experiment, the method of back titration will be employed. In the back titration method, excess acid is added to the antacid base which is placed in a flask; the antacid is allowed to neutralize as much of the acid as it can, but so much acid will be added that some acid will remain un-neutralized. Once the reaction is complete then a base will be added through titration to quantify how much acid was not neutralized by the antacid. If the original quantity of acid is known and the amount of un-neutralized acid in determined, then the difference between the two values will be the amount of acid that was neutralized by the antacid. The method of back titration is chosen because some of the antacids effervesce and to ensure complete reaction they need time to react. Direct titration with an acid would not afford this reaction time; thus, the end point of the titration would not be clear.
Since the stoichiometry of the acid-base reaction that is to be used is 1:1, at the endpoint of the titration the number of moles of NaOH added through titration will be equal to the number of moles of HCl in the flask (Reaction AN.3). \begin {equation} \ce{\label {HCl+NaOH} 1 HCl(aq) + 1 NaOH (aq) -> 1 H2O(l) + 1 NaCl(aq)} \end {equation}
Using the original molarity of the HCl and the original volume of HCl added to the antacid tablet, the total number of moles of HCl added can be determined. Using the molarity of the base and the volume of base used, the moles of NaOH consumed can be determined. This value is equal to the moles of HCl that was ‘untouched’ by the antacid. Finally, the difference between the total moles of HCl added and the moles of remaining HCl will be the moles of HCl neutralized by the antacid.
Hopefully molarity has been covered in the lecture course but if not then your lab instructor will guide you to use the definition of molarity, given in Equation AN.1: \begin {equation} \label {def_molarity} \text {Molarity} =\frac {\text {moles\ of\ solute}}{\text {liter\ of\ solution}} \end {equation}
Assume one tablet of an antacid has been dissolved in 30 mL of water in an Erlenmeyer flask. (The exact amount of water does not matter.) In this experiment, the molarities of the HCl and the NaOH solutions will be given on the reagent bottles.
Pour about 60-70 mL of the standardized acid into a beaker. Now using a 25.00 mL glass transfer pipet, pipet the acid up into the glass pipet using suction by a pipet bulb such that your finger is used to stop the liquid at the calibration line. Now, exactly 25.00 mL of 0.3636 M HCl is transferred by moving the pipet to the Erlenmeyer flask and allowing the liquid to drain into the flask. Do not blow out the remaining liquid at the tip of the pipet; the pipet is calibrated with that amount staying in the pipet.
Remember that we made sure more HCl was added than the antacid (basic) can possibly neutralize.
Therefore, the solution after the antacid has reacted with the HCl is acidic (has excess \( \ce {H3O+} \)).
With a buret you began at 1.42 mL and the buret drained down to 17.84
mL. The volume of the sodium hydroxide added ise the difference between
the final measurement on the buret and the initial measurement on the
buret:
\begin {equation} \text {Volume added by the buret} = \text {Final Volume reading}-\text {Initial Volume reading of the buret} \end {equation} This means you add 16.42 mL of 0.1017 M NaOH to neutralize the excess
acid.
You need to answer the following questions:
Thus:
Calculate the total moles of acid added to the flask containing the antacid pill: \begin {equation} \frac {0.3636\ \text {mole\ HCl}}{\text {L\ solution}}\times 0.02500L\ \text {solution\ added}= 0.009090\ \text {mole\ HCl\ total} \end {equation}
Note:
When molarity is multiplied by volume (M \(\times \) V) the units of volume cancel so we are left with moles
Once the number of mole of HCl neutralized by the antacid tablet is determined this value can be used to determine two other values, the cost effectiveness and the mass effectiveness.
The cost effectiveness is the amount of acid neutralized by the tablet per price paid (Equation AN.3). This factor can be used to see if the antacid is really the bargain it is advertised to be. \begin {equation} \label {cost_eff} \text {Cost\ effectiveness} =\frac {\text {moles\ of\ HCl\ neutralized\ by\ tablet}}{\text {cost\ per\ tablet}} \end {equation} The mass effectiveness is the ratio of the moles of HCl neutralized by the tablet to the mass of the antacid used (Equation AN.4). This value can be used to determine if the consumer will have to ingest 1 tablet or 5 tablets to neutralize a given amount of acid. \begin {equation} \label {mass_eff} \text {Mass\ effectiveness} =\frac {\text {moles\ of\ HCl\ neutralized\ by\ tablet}}{\text {mass\ of\ tablet(s)}} \end {equation}
Set up a buret using a buret clamp and an iron ring stand as shown in Figure AN.1. Learn about burets by viewing the Labflow video, Using and Reading a Buret.
1.
Obtain one antacid tablet (1000 mg \( \ce {CaCO3} \)) and record the mass
to the precision of the balance.
Note:
If the tablet contains only 500 mg \( \ce {CaCO3} \), obtain two and divide the total number of tablets in the bottle by two.
2.
Record the price of the bottle of antacid and the number of tablets listed on the
bottle (see Note above) and calculate the price per tablet, paying attention
to correct units for the result and the correct number of significant
figures.
3.
Place tablet in a 250 mL Erlenmeyer flask with 30 mL of DI water (by adding
water the number of moles of acid or base is not changed).
4.
Using a stir rod (or a magnetic stir bar and stir plate), gently crush the tablet
so that it dissolves in the liquid or at least is suspended in the liquid
uniformly.
5.
Pour an amount of about 70-85 mL of the standardized HCl solution from the
reagent bottle into a small beaker.
Note:
This volume is NOT the amount of HCl added to the Erlenmeyer flask.
Record the molarity from the bottle and label the beaker appropriately.
6.
Using a volumetric pipet and a pipet bulb, use suction to suck up a bit more
than 25.00 mL of the HCl acid into the volumetric pipet. Place your finger over
the top end of the pipet and wiggle your finger until the volume goes down and
reaches the calibration mark which signifies that 25.00 mL would be transferred
from this mark. Once the volume level is at the calibration mark transfer the
25.00 mL of HCl to the Erlenmeyer flask by placing the tip in the flask and
releasing your finger and allow it to drain. Do not blow out the remainder as the
transfer pipet is calibrated to deliver 25.00 mL such that a few drops may
remain.
7.
Allow the antacid to react with frequent stirring for 10 minutes before
proceeding.
8.
While the acid is reacting with the antacid, rinse a clean buret twice with small
portions of the NaOH solution.
9.
Close the stopcock and fill the buret with the NaOH solution. The buret does
NOT need to be filled to exactly the 0.00 mL mark. Ensure that the tip is full of
liquid by opening and closing the stopcock quickly to fill the tip. Drain extra
into your waste beaker, labeled appropriately.
10.
Hold a white card behind the meniscus and read the initial volume of NaOH
solution in the buret to the nearest 0.01 mL (Figure AN.2).
11.
Add 8-10 drops of the bromophenol blue indicator to the flask. If the indicator is
green or blue consult with your instructor. If the indicator is yellow, then
proceed to the next step.
12.
Open the stopcock and begin the titration of the antacid solution with
constant swirling of the receiving flask. If you have not performed a
titration before, your lab instructor or lab assistant will demonstrate for
you.
13.
When you are nearing the equivalence point, you will begin to see flashes of
bluish-purple color that persists longer and longer. When you start seeing these
flashes of color, slow down delivery of the NaOH from the buret. Otherwise, you
may “overshoot” the endpoint, giving you a false result for the amount of
NaOH required to use up all the HCl. At the end of the titration, you
should be adding the NaOH a drop at a time (that is not a squirt at a
time!).
14.
At the equivalence point (represented by the endpoint), the bluish-purple color
should remain. Note the final volume reading on the buret and record the value
to the nearest 0.01 mL.
15.
Repeat this entire process for an antacid tablet of a different brand (Trial 2).
Complete the report doing the necessary calculations.
16.
Discard your waste in the appropriate waste container and clean up your
workstation.
Name: _________________________________________________
Section: ________________ Date: ______________________
| Description | Trial 1 | Trial 2 |
|---|---|---|
| Name of antacid used | ||
| Mass of tablet(s) with 1000 mg \(\ce{CaCO3}\) | ||
| Price/bottle of antacid | ||
| Number of tablets/bottle | ||
| Price/tablet of antacid |
Calculation: Show calculations for Trial 1 here:
Calculation: Show calculations for Trial 1 here:
Molarity of NaOH used
Calculation: Show titration calculations for Trial 1 here:
