Faculty of Natural and Agricultural Sciences
Fakulteit Natuur- en Landbouwetenskappe
Chemistry Department/ Departement Chemie
CHM 172 (2014)
Practical 3
NB. Remember to bring your LABORATORY BOOK (with flow diagram
completed), your LAB COAT & SAFETY GLASSES, a box of matches and marker
pen. Remember correct Dress code!
Preparation before the practical session:
• Lab preparation: Read through this document intensively. Study the experimental procedures
and make a flow diagram for EACH EXPERIMENT in your LAB BOOK. (plan how you would
execute the procedure in the laboratory). Only your book will be allowed in the lab, not this
original document!!
th
• Reading: Indicated sections from Kotz & Treichel (8 Ed).
• Photo series: View the photo series on “The burette” (see ClickUP Course content page, left
hand side column “Titrations”). Answer the questions at the end of this Instruction sheet in your
Lab book.
• Videos: Before coming to the laboratory, view the short YouTube video at
www.youtube.com/watch?v=9DkB82xLvNE
Experiment 1: Preparing the antacid for analysis
Experiment 2: Determine the neutralizing effectiveness per gram of commercial antacid (ENO
Tums ® and Gaviscon ®)
Introduction
Required self-study: Kotz & Treichel and Townsend Section 4.1, 4.2, 4.7
(Important P181-185)
Qantitative chemical analysis
If you know (i) the balanced equation for a reaction and (ii) the exact quantity of one of the reactants,
you can then calculate the quantity of any other substance consumed or produced in the reaction.
This is the essence of quantitative chemical analysis, the determination of a given constituent in a
mixture. To determine the concentration of a particular solute in a solution, chemists often carry out a
titration, a process which involves combining a reagent of known concentration (standard solution),
from a burette in controlled amounts, to a sample of the solution of unknown concentration. The point
at which chemically equivalent quantities are present is known as the stoichiometric or equivalence
point. To determine when the equivalence point of a titration has been reached, a chemical
indicator, which changes colour at, or very near to, the equivalence point is used. The colour change
of an indicator takes place at the end point of the titration.
Use of stoichiometry in calculations
As an example, consider two compounds, H2SO4 and NaOH, both in solution, reacting with one
another according to the following equation:
H2SO4 (aq) + 2NaOH (aq) → Na2SO4 (aq) + H2O ()
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Suppose that the concentration of the acid solution (Ca) is known exactly. Let Ca = 0.200 mol/dm .
This is called the standard solution. The concentration of the base (Cb) is unknown and should be
determined in this example experiment.
1.
A clean burette is filled with the standard solution
2.
A known volume (called an aliquot) of the solution of unknown concentration (the NaOH) is
transferred to a conical flask by means of a clean pipette. Let this volume be Vb = 25.00
3
cm .
1
3.
4.
5.
6.
7.
A small amount of a suitable indicator is added to the solution in the conical flask.
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The standard solution is then carefully added drop-by-drop (with swirling ) from the burette
to the aliquot of the unknown in the conical flask.
This is done until the addition of just one drop suddenly causes the solution in the flask to
change colour, indicating the end-point of the reaction.
The exact volume of the standard solution (added from the burette) is determined by
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recording the initial and final burette readings. Let this volume be Va = 30.00 cm .
The concentration of the unknown solution can then be calculated from the stoichiometric
relationship, as is expressed by the balanced chemical equation of the reaction between A
and B, as follows (three basic steps):
30.00
7.1 Moles of acid added = na = Ca x Va = 0.200 mol x
dm3
3
1000
dm
-3
= 6.00 x 10 mol H2SO4
7.2 According to the equation, 2 moles of base react for every 1 mole of acid added to
the flask (stoichiometric ratio).
-3
∴ Moles of base in the flask = nb = 6.00 x 10 mol H2SO4 x 2 mol NaOH
1 mol H SO
2 4
-2
= 1.20 x 10 mol NaOH
n
7.3 ∴ Concentration of NaOH solution = Cb =
b = 1.20 x 10 - 2 mol
25
Vb
3
dm
1000
-3
= 0.480 mol.dm NaOH
The equivalence point is when there is 1 mole of H2SO4(aq) added from the burette for every 2 moles
of NaOH(aq) in the solution (from the balanced equation).
This equivalence point of any titration is that stage at which stoichiometrically equivalent quantities of
the main reactants have been brought together; i.e. the point at which there is no excess of either
reactant. This point is indicated by an indicator.
If a strong acid is titrated with a strong base, the pH of the solution will be ±7 at the endpoint. An
indicator showing a colour change at this pH, should be used for titration. Therefore bromothymol blue
is the best choice.
The results of a volumetric analysis will only be accurate if the necessary precautions are taken in
the planning and execution of an analysis. Thus, you must use the volumetric apparatus (pipettes,
burettes and volumetric flasks) correctly, and be completely acquainted with the underlying
principles of the procedure. For a burette, as well as the pipette, the volume should be
recorded to two decimal places.
Antacids
+
All antacids have one purpose – to neutralize the excess acid (hydronium ion - H3O ) in the stomach
to relieve acid indigestion. Commercial antacids claim to be the “most effective” for relieving acid
indigestion – and our aim with this experiment is to compare the effectiveness of two commercial
products.
The pH of the gastric juice in the stomach ranges from 1.0 to 2.0. The acid in the stomach (mainly
hydrochloric acid) is necessary for the digestion of the food we eat. It is continually secreted while
eating, therefor overeating can result in an excess of stomach acid and a pH lower than normal. This
condition causes irritation of the lining of the stomach (“heartburn”). An antacid (with base properties)
reacts with the hydronium ion to relieve the symptoms.
An example of an over-the-counter antacid is Milk of magnesia, an aqueous solution of magnesium
hydroxide (Mg(OH)2).
1
To cause the liquid in the flask to move around quickly in a twisting, circular fashion.
2
The neutralising reaction in the stomach is the following:
2+
Mg(OH)2(s) + 2H3O(aq) → Mg (aq) + 4H2O(ℓ)……………….(1)
The following are examples of antacids frequently used.
Antacid name
Rolaids
Tums
Maalox
Milk of Magnesia
Alka-Seltzer II
Gaviscon
Amphojel
Main component
CaCO3
CaCO3
Mg(OH)2 and Al(OH)3
Mg(OH)2
NaHCO3 and KHCO3
Al(OH)3
Al(OH)3
Excessive use of antacids can lead to a potentially dangerous condition, namely an additional
excretion of stomach acid due to the fact that the antacid keeps the pH of the stomach fluid above 2.
This increase in acid production, causes a chronic problem.
To decrease the stomach fluid from becoming too basic from an antacid, buffers are added as part of
the
formulation of some antacids. Antacids containing CaCO3 and/or NaHCO3 buffer the pH of the
2stomach by establishing a HCO3 / CO3 buffer system:
CO3 (aq) +H3O (aq) → HCO3 (aq) + H2O(ℓ)………………….(2)
+
HCO3 (aq) +H3O (aq) → CO2(g) + 2 H2O(ℓ)….……………….(3)
2-
-
+
Note that CO2 is generated, and is the cause of the burp that often is expelled.
In this experiment, a strong acid-strong base titration is used to determine the neutralizing power of
two different antacids. To obtain the quantitative data for the analysis, which requires a well-defined
endpoint in the titration, the buffer action is eliminated. This is done by first adding an excess of
standardized hydrochloric acid, HCℓ (concentration known), to the antacid solution, neutralising the
basic properties of the antacid. The solution is then heated to remove developing carbon dioxide. At
this point, ALL moles of base in the antacid (whether or not a buffer is present), have reacted with the
standardized HCℓ solution. The unreacted (excess) HCℓ is then titrated with a standardized sodium
hydroxide solution, NaOH (concentration known). The more NaOH required, the less effective the
antacid was.The analytical technique used here is called back titration.
Take notice: the number of moles of base in the antacid of the commercial sample plus the number
of moles of NaOH used in the titration is equal to the number of moles of HCℓ added to the original
antacid sample.
To summarize:
moles (HCℓ) = moles (base in antacid) + moles (NaOH)
The moles of base in the antacid per gram of antacid provide the data required for a comparison of
the effectiveness of the two commercial antacids.
Experiment 1: Preparing the antacid for analysis
Aim: To prepare the antacid for analysis by reacting it with an excess of HCℓ and removing all
developing CO2 gas.
Apparatus and equipment
Bunsen burner
Glass beaker (250 mL)
Glass funnel
Burette
Chemicals
NaOH(aq) (0.100 M)
HCℓ (aq) (0.114 M)
Bromophenol blue indicator
Sample 1: ENO tums ®
Sample 2: Gaviscon ®
Safety
Wear eye protection and a laboratory overcoat.
4x Erlenmeyer flasks (250 mL)
Wooden tongs
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Boiling stone
Hydrochloric acid is a strong acid. Avoid contact
with skin and do not spill.
Procedure:
Do this procedure in duplicate for ENO tums:
1. Determine the mass of an Ehrlenmeyer flask and record it.
2. Add no more than 0.20 g of the pulverized commercial antacid to the flask and measure the
combined mass. Record.
3. From the burette, carefully release 25 mL of a standardized 0.114 M HCℓ (stomach acid
equivalent) into the Erlenmeyer flask and swirl. Record the initial and final volumes on the
burette and calculate the exact volume released.
4. Add a boiling stone and warm the solution slowly to a very gentle boil on a Bunsen burner.
Maintain the heat for 1 min to remove dissolved CO2. Swirl gently every few seconds (use the
wooden tongs).
5. Add 2 drops of bromophenol blue indicator. If the solution is blue, add an additional 10.0 mL
of hydrochloric acid into the solution and boil again. Record the total volume of HCℓ that is
added to the antacid.
6. Repeat the procedure in duplicate with Gaviscon.
7. Leave the prepared antacid to cool.
Experiment 2: Determine the neutralising effectiveness per gram of a commercial antacid (ENO
Tums and Gaviscon)
Aim: To master the technique of a back-titration
Procedure
1. Rinse a burette three times with water to clean it.
2. Rinse it with 5 mL of standardised NaOH (0.100 M) solution.
3. Fill the burette with NaOH solution; be sure no air bubbles are in the burette tip. Read and
record the initial volume to two decimal places.
4. Titrate the sample with the NaOH solution until a faint blue endpoint. Watch closely, the
endpoint may appear after only a few millilitres of titrant, depending on the concentration of
the antacid in the sample.
5. When a single drop of NaOH solution changes the sample solution from yellow to a faint
green-blue, stop. Wait for 10-15 seconds, if the colour change is permanent, read and record
the final volume.
6. Flush the burette with tap water, dispense through the burette tip. Repeat this 3 times.
Leave the burette upside down to dry.
Complete the report sheet and submit your results at the end of the practical session.
Please clean the laboratory bench where you were working – throw all pieces of cleenex
paper in the dustbins provided.
Thoroughly clean all the apparatus used.
Return all equipment used to the locker.
Questions (to answer in your Lab book):
1.
2.
3.
4.
5.
Why do we do a “rough” titration first?
How do we calculate the “volume titrated”?
To how many decimals should the volume be recorded?
Why do we have to remove the funnel from the burette top before we start the titration?
How do we know when the “end point” is reached?
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