In the very first laboratory course of the first semester, there are five
experiments based on the titrimetric analysis besides the elemental analysis of
organic compounds and experiments based on chromatography. The
introductory unit of the laboratory course has been designed to familiarise you
with the apparatus which will be used in performing the titrimetric experiments
and some of the basic concepts on which the experiments are based.
You might have studied in your previous classes that in titrimetry we estimate a
substance in solution by titrating it against the standard solution of an
appropriate substance. The whole process involves the use of specific
apparatus that you need to be thoroughly familiar with. Therefore, first of all we
introduce you to the apparatus commonly used in titrimetric analysis, and
explain its correct use. We also tell you how to make a standard solution and
express its concentration. The weighing of chemicals being a part and parcel of
these experiments, the correct use of analytical balance has been dealt in
detail. Finally, we introduce you to the common safety measures one should
observe in a Chemistry laboratory. BCHCL-132 Chemistry Lab I
8
Expected Learning Outcomes
After studying this unit, you should be able to:
measure and deliver sample volumes by selecting and using appropriate
apparatus for titrimetric measurement;
determine the mass of a sample by correctly using analytical balance;
perform basic laboratory skills, including pouring reagents and transferring
solids, preparing solutions of known concentrations;
list and explain different types of titrations; and
explain the safety measures as applied to a chemistry laboratory.
1.2 APPARATUS COMMONLY USED
Titrimetric analysis involves reliable and accurate measurement of volumes of
solutions. Three pieces of apparatus, namely, a pipette, a burette and a
volumetric flask are indispensable for this purpose. Their use is described
here. Before doing the experiment you should go through the instructions
given below carefully and work accordingly.
1.2.1 How to Use a Pipette
Pipette is used to measure and transfer known volume of a liquid from one
container to the other.
A pipette is shown in Fig. 1.1 (a). As you can see, it is a long tube with a bulb
in the middle. On the narrow upper part of the pipette a horizontal line is
marked. This line indicates the level to which the pipette has to be filled to
deliver the liquid equal to the volume indicated on the bulb when used in the
way described below. Pipettes can be of different capacities like 1, 2, 5, 10,
20, 25, 50 cm3
, etc. You will use pipettes mostly of 10 and 20 cm3
for your
experiments.
Fig. 1.1: (a) Pipette. (b) Handling of a pipette. (c) Correct way to drain
out the solution.
Pipettes which can
measure volumes of less
than 1 cm3
are also
available with special
accessories.
CAUTION!
Do not suck corrosive
liquids like strong acids
and alkalies by mouth.
You can use a rubber
teat for this purpose. Unit 1 Titrimetry: An Introduction
9
Before using a pipette, it has to be thoroughly washed with a good quality
detergent followed by plenty of water and finally with distilled water. This
removes all the grease. It is then rinsed with the solution which has to be
measured. For rinsing, the solution is taken in a clean and dry beaker. The
pipette is dipped deep into the solution and the solution is sucked into the
pipette to fill it up to about half its volume. It is then taken out and the solution
is made to wet it completely from inside by moving the solution up and down
and also around its axis. The solution is drained out and the whole process is
repeated. The pipette is then filled with the solution until the level is about
2 cm above the mark. The top of the pipette is then quickly closed by slightly
moist (not wet) index finger; see Fig. 1.1 (b). The pressure of the finger is
slowly released so as to allow the solution to run out until the lower meniscus
just touches the mark. The solution from the pipette is transferred into the
container in which titration has to be done. The solution is allowed to run out
on its own. The last drop of the solution which does not seem to drain out by
itself is taken out gently by touching the tip of the pipette with the walls of the
container for about 3-4 seconds; see Fig. 1.1 (c). Do not blow out the last
drop. The pipette is calibrated to include the liquid column trapped at the tip.
Further, blowing makes it dirty and CO2 in the breath may react with the
solution being pipetted. The volume of the liquid thus transferred through the
pipette is equal to the volume marked on the pipette.
Another type of pipette is designed to deliver definite but different volumes of a
liquid. It is called a graduated pipette, Fig. 1.2. It has got marking
corresponding to different volumes. It is also used in a similar fashion, with the
only difference that the liquid is not completely drained out; instead the volume
required is transferred.
SAQ 1
Why should you not blow the last drop out of the pipette?
1.2.2 How to Use a Burette
A burette is designed to transfer definite but variable volumes of a liquid into
another container.
A burette is a long glass tube, commonly of 50.0 cm3
capacity in 0.1 cm3
unit
graduation marks, Fig. 1.3. It has a stop cock at the lower end to control the
amount of solution drained. The burette also has to be washed, first with a
detergent followed by plenty of water and finally with distilled water. It is then
rinsed with the solution to be measured. For rinsing it is filled a little less than
half with the solution and by repeatedly rotating and tilting the burette, the
solution is made to wet it completely from inside. This solution is discarded.
The burette is then mounted on the stand in an upright position and is filled
carefully with the help of a funnel. After taking out the funnel, the meniscus is
adjusted to a definite graduation mark by drawing out some solution through
the stop cock. The bottom of the meniscus should just touch the graduation
mark. While reading the solution level in the burette, your eyes should be on
level with the graduation mark, otherwise there would be error due to parallax,
Fig. 1.4. It is not necessary to adjust the meniscus at the zero mark level, if it
The curved surface of a
liquid in a container is
known as the meniscus.
The meniscus in case of
liquids which stick to the
container, is concave,
e.g. for water and
aqueous solutions, while
it is convex in case of
liquids which do not stick
to the container, e.g. for
mercury.
Fig. 1.2: Graduated
Pipette. BCHCL-132 Chemistry Lab I
10
is too high for the level of your eyes. You can adjust it at, say 10.0 cm3
or any
other convenient level.
Error in burette reading is among the most common sources of error in
titrimetric analysis. To make the meniscus more distinct and to ensure that it
looks the same always, it is convenient to place a screen behind the burette as
shown in Fig. 1.5. This can be made from a small piece of cardboard covered
with white paper with the lower half blackened with ink. The black part is to be
held downward. This is called a parallax card. You can ask your counsellor to
show you how to make a parallax card.
Fig. 1.3: Burette. Fig. 1.4: Eye level for
burette reading.
Fig. 1.5: Reading the burette with
the use of the parallax
card.
After adjusting the meniscus, the level of the solution in the burette is
recorded. This is called the initial reading or initial volume. Then the titration
is performed and at the end of the titration, the level of the solution is
recorded. It is called the final reading or final volume. The difference of the
two readings, (final reading – initial reading), gives the volume of the solution
transferred to the titration flask. The correct way of delivering a liquid from
burette is shown in Fig. 1.6.
Fig. 1.6: Delivery of liquid from a burette.
PRECAUTION!
No standard apparatus
is to be heated above
298 K. Unit 1 Titrimetry: An Introduction
11
1.2.3 How to Use a Volumetric Flask
A volumetric flask is used to prepare a definite volume of a solution of
precisely known concentration.
Volumetric or measuring flask has a flat bottom with a long, narrow neck,
Fig. 1.7. It has a calibration mark on its neck which indicates the level up to
which the flask is to be filled to get a volume equal to the one indicated on the
flask.
Fig. 1.7: Volumetric Flasks.
You will be using volumetric flasks of 100 cm3
and 250 cm3
capacity. The
flask, before use, is cleaned thoroughly, washed with a detergent and plenty of
distilled water and allowed to drain. The weighed compound is transferred into
the flask with the help of a funnel. It is first dissolved in just enough water; the
solution is then made up to the mark by carefully adding more distilled water.
This can be done with a wash bottle or better with a pipette. The flask has to
be stoppered tightly and shaken well before use to get a homogeneous
solution.
1.2.4 How to Use an Analytical Balance
In titrimetric analysis, you will invariably have to prepare a standard solution.
You would be required, for this purpose to weigh a solid accurately by using
an analytical balance. It is very important to learn the use of an analytical
balance because accurate weighing is important for the accuracy of any
titrimetric experiment.
Generally two types of analytical balances are used in chemical laboratories.
These are:
(i) Double pan analytical balance
(ii) Single pan electronic analytical balance
Double pan analytical balance is now outdated but still some of our
laboratories do not have electronic analytical balance and may have only
double pan balance. A commonly used double pan analytical balance is
shown in Fig. 1.8. The various parts of the balance are labeled in the figure.
Before using the balance, you have to first determine the zero point of the
balance. For this purpose, the side doors of the balance are closed and the
arrest knob (1) is slowly and carefully turned counter-clockwise. Avoid jerks
as they may disturb the setting of the balance.
Zero point is the point
on the scale at which the
pointer of the unloaded
balance comes to rest. BCHCL-132 Chemistry Lab I
12
When the arrest knob is turned fully to the left, the pointer (2) starts swinging
around the centre of the scale (3). The first two swings are ignored and
starting with the third swing, the extreme positions of the swing are noted. The
swings to the right are positive and those to the left are negative.
Fig 1.8: Analytical balance.
The readings to the left and right are averaged separately and the mean of
these averages is found, which is the zero point. The following example will
make it clear.
Reading on the left Reading on the Right
1.
2.
3.
4.
– 5.0
– 4.0
– 3.0
– 2.0
+ 5.0
+ 4.0
+ 3.0
– 14.0 + 12
Average 3.5
4
14.0
−=
−
=
4.0
3
12.0
=
Mean Value 0.25
2
43.5
=
+−
=
The zero point is + 0.25, i.e. 0.25 units to the right.
Such small discrepancies between the zero point and the middle of the scale
may be ignored as they are insignificant. However, if the deviation is large,
Ideally the zero point and
the mid or zero of the
scale should be the
same.
7 Unit 1 Titrimetry: An Introduction
13
e.g., greater than 1.5 units, the balance must be adjusted by means of the
screws (4), for which you may request your counsellor.
After adjusting the zero point of the balance (if necessary), we come to actual
weighing. For this purpose, we use a glass or a plastic weighing bottle,
Fig. 1.9. First of all, the weighing bottle is weighed on a rough balance to find
its approximate mass to the nearest gram. Then, the left side door of the
analytical balance is opened and the weighing bottle is kept on the left side
pan (5) and the door is closed. Similarly, through right side door, weights equal
to the approximate mass of the weighing bottle are transferred to the right side
pan from a weight box; Fig. 1.10.
Fig. 1.10: Weight box and weights.
The arrest knob is once again turned to the left and the movement of the
pointer is seen. If it moves more to the left, then the weights transferred are in
excess of the mass of the bottle. In that case some weights have to be
removed. On the other hand, if the pointer moves to the right, then the added
weights are not sufficient and we need to add more weights. Arrest the
movement of the beam by turning the arrest knob fully towards the right and
open the right side door to add or remove some weight(s), as the case may
be. Recheck the movement of the pointer by turning the arrest knob. Continue
this process till the addition of 1 gram weight makes the right hand pan heavier
while its removal makes it lighter, e.g., if the weight is say 15.5 g, then 15 g
weight would be lighter and 16 g weight would be heavier. After this, the
fractional weights marked in mg, have to be added in the order of decreasing
weight till the two sides are balanced. Do not use fractional weights of less
than 10 mg, you should use a rider in such cases. A rider, Fig. 1.10, is a thin
metallic wire suitably bent to be seated on the beam of the balance. It is
normally put on the right hand side of the beam (6) with the help of the rider
carrier (7). By varying the position of the rider on the beam (8), the rest point is
found, i.e., the two pans are balanced.
The beam scale has got markings from 0-10 on either side. It is calibrated in
such a way that each main division is numerically equal to mass in milligram,
when the rider is put on it. Each main division is further divided into 5
subdivisions and each subdivision is equivalent to 0.2 mg. thus the accuracy
of such an analytical balance can be only up to 0.2 mg. The mass of an object
can be calculated using the following formula:
You must close both the doors of the balance before raising the pans with the arrest knob.
Always use forceps to
transfer the weights.
Refrain from using
your hands.
mass adjustments
below 10 mg/.01 g. BCHCL-132 Chemistry Lab I
14
Mass of the object = (Weights added in grams) + (Fractional weights added ×
0.001) g
+ (Main division of the rider position × 0.001) g
+ (Subdivision of the rider position × 0.0002) g
Let us illustrate the use of this formula. Suppose that while weighing an object,
the weights added to the right side pan are 15 g, 200 mg and 2 × 20 mg. Let
the rider position be 3 on the subdivisions after 2 main divisions.
Then the mass of the object
= 15.00 g + (240 × 0.001) g + (2 × 0.001) g + (3 × 0.0002) g = 15.2426 g
You have, so far, seen how to weigh an object accurately. If we want to weigh
substance in the weighing bottle, we make use of the method of weighing by
difference. For this, the weighing bottle is first approximately weighed. The
substance to be weighed is put into the bottle (a little more than required) and
weighed accurately (m1 g). The substance is transferred into a volumetric flask
and the bottle is again weighed accurately (m2 g). The difference of the two
masses, i.e., (m1 − m2) gives the exact amount of the compound transferred
(m g).
1.2.5 Single-Pan Electronic Analytical Balance
In case of electronic balance also known as digital balance, the mass of the
object being weighed will be digitally displayed. In electronic balances
(Fig.1.11), substance can be added to or removed from the balance without
any problem.
Handling of Electronic Balance
Balances of all types should never be subjected to harsh treatment of any
description. Always make adjustments smoothly and carefully. You will need to
ensure that the balance reads zero before you start to weigh. However, if you
are using a balance with a tare facility, it is not necessary to zero the balance.
Fig.1.11: Single-pan electronic analytical balance.
Mass of the substance
(m g) = Mass of the
bottle with substance
(m1 g) − Mass of the
bottle after transferring
the substance (m2 g)
m = m1 – m2 g Unit 1 Titrimetry: An Introduction
15
The tare facility allows you to cancel out the mass of the container so that you
can weigh the required mass of material without having to take into account
the mass of the container. Even if you cannot tare out the complete mass of
the container, the tare facility can still be used to adjust the mass indicated by
the scale to a convenient whole number.
For example, suppose you are using a balance that can only tare out 10 g, but
your container weighs 22.45 g and you require to weigh 8.70 g of material. In
such a case instead of
(1) weighing the container,
(2) adding its mass to the required mass of material, and
(3) adding material to the container until the balance reads the calculated
mass,
you can use the tare facility to adjust the displayed mass of the container to
read 20.00 g and then add material to the container until the balance reads
28.70 g. This avoids errors in weighing that arise due to miscalculations! Of
course, when you have finished weighing, you should remove any tare and
adjust the balance to read zero.
An important point to note here is that you should always use a container for
weighing any material. Even a piece of filter paper is not very satisfactory as
solid chemicals can easily spill onto the balance pan. As already mentioned
small glass or plastic weighing bottles are available and these are most useful.
For very accurate weighing, e.g. in the preparation of a standard solution it
may be necessary to use a weighing bottle.
Having learnt about the general apparatus to be used in the experiments for
the first laboratory course, let us now understand the various terms and
concepts used in these experiments. Before this, try to answer the following
SAQ.
SAQ 2
What is the mass of a substance if the following weights are needed to weigh
it?
g mg position of rider
5 200 8.2
2 100
1 50
1.3 EXPRESSION OF CONCENTRATION
In a qualitative sense, the term concentration deals with the "crowdedness" of
the particles of solute in a solution. A solution having more number of solute
particles per unit volume is said to be more concentrated. In quantitative
analysis, one very often comes across this term. Before we give an expression
for this, it would be worthwhile to recapitulate a few relevant fundamental
concepts here.
Solute is the dissolved
substance in a solution.
Solvent is the liquid in
which the solute is
dissolved. Solution is
the homogeneous
mixture of a solute and a
solvent.
Never weigh anything
(apart from the
container) directly on the
balance pan. BCHCL-132 Chemistry Lab I
16
Mole, denoted as mol, is the amount of a substance that contains as many
elementary entities as are there in 0.012 kg of C12 isotope of carbon. The mole
may be of atoms, ions, molecules, electrons or any other entity. The number of
elementary entities in a mole of any substance is fixed and is given by a
constant called the Avogadro's number, NA which equals 6.022 × 1023
.
Relative Molecular Mass (Molecular Weight) denoted as Mr
, is the mass of
one molecule in atomic mass unit (a.m.u.) relative to 1/12th of the mass of the
pure C12 isotope (12.000 a.m.u.). For most titrimetric analyses, purpose of this
is the same as the old atomic mass and molecular mass. We find it by
multiplying the atomic mass of each element in the molecule by its subscript in
the formula and then addin
Cuts: The most common accidents in the chemistry laboratory are cuts
from broken glassware. If you have a cut, wash the wound well with cold
water immediately. If bleeding is severe, apply pressure directly on to the
wound to stop the bleeding. Then an antiseptic cream can be applied to
the wound with a proper dressing.
ii) Burns: Burns generally caused by hot equipment can be treated as the
cuts are treated, that is, wash the burnt part with cold water for sometime
and then apply Burnol to it.
iii) Fire: A small fire in a beaker, caused by the vapours of an inflammable
liquid, can be extinguished by covering it with a watch glass.
If the clothes catch fire one should lie on the floor and fire can be
smothered by wrapping a blanket around the body.
iv) Poisoning: If one happens to swallow a poisonous chemical, plenty of
water should be given if the person is conscious. For a corrosive poison,
calcium hydroxide solution (lime water) should be given as soon as
possible. An antidote should be given only in the case of non-corrosive
poisons. BCHCL-132 Chemistry Lab I
28
v) Explosion: Sometimes a faulty technique during the experiment can lead
to an explosion. 'You should work with highly oxidizing or explosive
chemicals only under strict supervision'.
Table 1.3 gives the remedies for a few common chemical reagents used in the
laboratory.
Table 1.3: Remedies for a few chemical reagents
Chemical Neutralising wash
Acid like HNO3, H2SO4, HCl Initial action should be a thorough washing with cold
water. Then NaHCO3 or 2M ammonium carbonate
(leaves no residue on clothes), apply Vaseline or a
soothing cream.
Alkalies, e.g., NaOH, KOH
etc.
1M acetic acid, then apply Vaseline or a soothing
cream.
Bromine 2M Ammonia, keep the affected part dipped in
NaHSO3 till bromine is washed off, then apply
Vaseline.
Phenol Ethanol and then hospital treatment
Sodium Ethanol on a cotton wood pad
1.8 ANSWERS
Self-Assessment Questions
1. The pipette is calibrated to include the liquid column trapped at the tip.
Further, blowing it makes it dirty and CO2 in the breath may react with the
solution being pipetted.
2. (5 + 2 + 1) g + (200 + 100 + 50) × 0.001 g + 8 × 0.001 g + 2 × 0.0002 g.
= 8 g + 0.350 g + 0.008 g + 0.0004 g
= 8.3584 g
3. From Eq. 1.1, M =
M V
m
m
1000 mol dm-3
Where Mm = 40 g mol-1
m = 4.000 g
V = 500 cm3
Therefore,
3
dm
40 500
1000 4 000 −
×
×
=
. M = 0.200 mol dm-3
Thus, molar concentration = 0.200 M Unit 1 Titrimetry: An Introduction
29
4. Again consider Eq. 1.1,
3
m
dm 1000 −
=
M V
m M
Where Mm = 169.87 g mol−1
V = 1 dm3
= 1000 cm3
M = 0.1 M
On substituting these values in the above equation, we have
1000
10 × 169 87 × 1000
=
. . m
= 16.987 g
Thus, mass of AgNO3 required for 0.1M solution = 16.987 g.
5. (a) i) NaOH is hygroscopic,
ii) It is not available in pure form as it combines with CO2 from the air
and some part of it is converted into sodium carbonate.
(b) Benzoic acid fits most of the criteria, but its solubility in water is low,
although in non-aqueous solvents such as ethanoic acid (acetic acid)
or ethanol it is not so.
6. i) -
ii) ×
iii) ×
iv) BCHCL-132 Chemistry Lab I
30
EXPERIMENT1
DETERMINATION OF SODIUM
CARBONATE AND SODIUM
HYDROGEN CARBONATE
PRESENT IN A MIXTURE
Structure
1.1 Introduction
Expected Learning Outcomes
1.2 Principle
1.3 Requirements
1.4 Procedure
1.5 Observations
1.6 Calculations
1.7 Results
1.8 Answers
1.1 INTRODUCTION
You might be familiar with the basic principle of acid-base titrations. In this
experiment we are expanding acid-base titration methods further for the analysis of
a mixture of sodium carbonate (Na2CO3) and sodium hydrogen carbonate
(NaHCO3) or sodium bicarbonate. This method of titration will help you in
understanding the basic principle of some important industrial analyses such as that
of soda ash (anhydrous sodium carbonate), washing soda (hydrated sodium
carbonate, Na2CO3.10H2O), baking soda (sodium hydrogen carbonate), mixture of
sodium carbonate−sodium hydroxide, commercial caustic soda (NaOH), etc. All
these commercial products have a tendency to absorb moisture and carbon
dioxide from atmosphere. This phenomenon of absorption of moisture and carbon
dioxide by these chemicals is called weathering. After such weathering soda ash
and washing soda contain appreciable moisture and sodium hydrogen carbonate
and caustic soda has appreciable amount of moisture and sodium carbonate.
Therefore, it becomes necessary to analyse the purity of these chemicals before
using them for any fine use. The procedures, such as, conductometry,
potentiometry or acid-base indicator methods can be used to analyse the above
substances. Here, we will discuss the acid-base indicator method only for the
analysis of a mixture of sodium carbonate and sodium hydrogen carbonate. Experiment 1 Determination of Sodium Carbonate and sodium Hydrogen Carbonate Present in a Mixture
31
Procedure used for the titration of a mixture of sodium carbonate and sodium
hydrogen carbonate is basically the same as that of the acid-base titration which
you have performed in your earlier classes, except that there are two analytes in
our sample.
Expected Learning Outcomes
After performing the experiment given, you should be able to:
state and explain the principle of acid-base titration with special reference to
the titration of sodium carbonate and sodium hydrogen carbonate
mixture;
prepare a standard solution of sodium carbonate;
standardise the given solution of hydrochloric acid and use it in estimating the
strength of basic solutions; and
determine the strength of sodium carbonate and sodium hydrogen
carbonate in a given solution.
1.2 PRINCIPLE
During an acid-base titration, the pH changes in a characteristic way. The pH
changes during titration can be understood by plotting a pH curve. A pH curve is
formed, if the pH of the solution being titrated is plotted against the volume of
solution added. The titration of sodium carbonate with a strong acid such as HCl
produces the titration curve shown in Fig. 1.1.This titration curve has two
equivalence points. You may like to ask, why does sodium carbonate solution
behave this way? To answer this question we should study the behaviour of
sodium carbonate in aqueous solution.
Sodium carbonate is a salt of a weak acid and a strong base; when such salts are
dissolved in water, they behave as bases due to the basicity of the conjugate base
CO2−
3
. The equilibrium, which is often called hydrolysis, is given by the reaction:
CO 2−
3 + H2O HCO −
3 + OH–
...(1.1)
(carbonateion) (hydrogen carbonate ion)
The hydrogen carbonate ion is furtherhydrolysed to carbonic acid:
HCO −
3
+ H2O H2CO3 + OH–
...(1.2)
(carbonic acid)
The OH–
ions so produced in solution are responsible for the basic character of sodium
carbonate.
When sodium carbonate is titrated with a strong acid, such as hydrochloric acid, the
reaction is completed in two steps. First, the carbonate ions are converted to the
hydrogen carbonate ions, and then to carbonic acid. This is due to the fact that a
strong acid displaces a weak acid from the conjugate base of the latter.
CO 2−
3 + H+
→ HCO −
3
...(1.3)
(Na2CO3) (HCl) BCHCL-132 Chemistry Lab I
32
HCO −
3 + H+
→ H2CO3 ... (1.4)
(HCl)
1 Equivalence point st
2 Equivalence point nd
7
14
pH
Volume of strong acid added in cm 3
Fig. 1.1: The titration of sodium carbonate with hydrochloric acid.
As neutralisation takes place in two steps (it is indicated by reactions in Eqs. 1.3 and
1.4), we observe two regions of sharp pH change in the titration curve (Fig. 1.1) and
thus two equivalence points at pH 8.31 and pH 3.69. As shown in Fig. 1.1, at the first
equivalence point (in the pH range 9-7) CO 2−
3
is neutralised to HCO −
3 and at the
second equivalence point (in the pH range 5-3) HCO −
3
is neutralised to H2CO3.
Combining both the above equations we can write complete neutralisation reaction of
sodium carbonate as
CO 2−
3
+ 2H+
→ H2CO3 C(1.5)
Finally, the carbonic acid produced as a result of these titrations can
decompose into carbon dioxide (CO2)
H2CO3 (aq) → CO2 (g) + H2O
In this experiment, we will utilise this behaviour of sodium carbonate in the estimation
of a mixture of sodium carbonate and sodium hydrogen carbonate. Let us now
understand the behavior of a mixture of sodium carbonate and sodium
hydrogen carbonate during titration with hydrochloric acid.
The titration curve for a sodium carbonate and hydrogen carbonate mixture is shown
in Fig. 1.2. As you can see, it has two equivalence points. The first equivalence point
indicates half neutralisation of the carbonate in the given sample, i.e., its conversion
to hydrogen carbonate (cf Eq. 1.3). The second equivalence point indicates
neutralisation of the hydrogen carbonate in the initial sample mixture and the
hydrogen carbonate just generated from the half neutralisation of carbonate
(cf Eq. 1.4). In this experiment we will be using acid base indicators to detect
both the equivalence points. Experiment 1 Determination of Sodium Carbonate and sodium Hydrogen Carbonate Present in a Mixture
33
1 Equivalence point st
2 Equivalence point nd
7
14
pH
Volume of H l added in cm C
3
Na CO +
NaHCO
2 3
3
CO3
2- HCO3
-
HCO3
H -
+
H
+
H CO 2 3
Fig.1.2:Titration curve for sodium carbonate and sodium hydrogen carbonate titrated
with hydrochloric acid.
As you know, acid-base indicators are organic dyes which change colour as
pH changes. This is because the indicator has two forms, one is acid form in
lower pH medium and other is base form in higher pH medium and having two
different structures. For example, phenolphthalein has two forms, one is
benzenoid form (I) in low pH medium (pH < 8) and thus, it is colourless. Its
second form has quinonoid structure (II) in high pH medium (pH > 10) which
has a pink colour. Similarly, methyl orange exists in quinonoid form (III) in
lower pH medium (pH < 3.2) and benzenoid form (IV) in higher pH medium
(pH > 4.4). The colour of its benzenoid form is yellow while that of quinoniod
form is red.
O
O
OH
OH
C
O
-
O
O
O
-
Quinonoid form in basic
solution (pink)
(II)
Benzenoid form in acidic
solution (colourless)
(I)
OH-
H
+
Phenolphthalein
N N -
O3S N
CH3
CH3
NH N -
O3S N
+
CH3
CH3
Quinonoid form in acidic solution (red)
(IV)
Benzenoid form in basic solution (yellow)
(III)
H
+ OH-
Methyl orange
Volume of HCl added in cm3
H2CO3
HCO3 BCHCL-132 Chemistry Lab I
34
Equivalence point so obtained using indicators is called an end point. It is not
necessary that the end point is coincident with the equivalence point, because
of the delay in getting the indicator to show the colour change, and other
factors. Ideally end point and equivalence point should be as close as
possible.
In general acid-base indicators show colour change in a pH range of ±1 pH
unit. The colour change and the pH range of some common indicators are
tabulated below to guide you for selecting appropriate indicator for any acid
base titration.
Table 1.1: Colour changes and pH ranges of acid-base indicators
Using Fig. 1.2 and Table 1.3, you can select phenolphthalein and methyl orange as
most suitable indicators for the detection of the first and the second end points,
respectively for the titration of mixture of sodium carbonate and sodium hydrogen
carbonate. Once, these two end points are detected, the volume of HCl used to titrate
sodium carbonate and sodium hydrogen carbonate in the mixture can be determined.
This can be further illustrated by considering Fig. 1.3.
Va
initial Vb
(1st end point) Vc
(2nd end point)
CO2 HCO3
HCO3
H2CO3
(V1
= Vb
- Va
)
(V2 = Vc
- Va
)
(Total hydrogen carbonate)
(V3
=V2
-2V1
)
2V1
− − −
Fig. 1.3: Volumes of HCl used during titration of a mixture of sodium carbonate
and sodium hydrogen carbonate.
Name Acid Colour pH Range of
Colour Change
Base Colour
Methyl violet Yellow 0.0 - 1.6 Blue
Thymol blue Red 1.2 - 2.8 Yellow
Methyl orange Red 3.2 - 4.4 Yellow
Bromocresol
green
Yellow 3.8 - 5.4 Blue
Methyl red Red 4.8 - 6.0 Yellow
Litmus Red 5.0 - 8.0 Blue
Bromothymol
blue
Yellow 6.0 - 7.6 Blue
Thymol blue Yellow 8.0 - 9.6 Blue
Phenolphthalein Colourless 8.2 - 10.0 Pink
The pH range is termed
as the colour-change
interval of the pH
indicator. The position of
the colour-change
interval in the pH scale
varies widely with
different indicators. For
most acid-base titrations,
we can, therefore, select
an indicator which
exhibits a distinct colour
change at a pH close to
the equivalence point.
For example, pH
indicator phenolphthalein
shows colour change in
pH range 8-10 and
methyl orange in the pH
range 3.1-4.4. Therefore,
for detection of end
points in the pH range of
8-10, phenolphthalein
will be the suitable
indicator. Similarly for
detection of the end point
in pH range of 3-4,
methyl orange will be the
suitable indicator. Experiment 1 Determination of Sodium Carbonate and sodium Hydrogen Carbonate Present in a Mixture
35
In this diagram Va, Vb and Vc refer to burette readings−initial, at the 1st end point
with phenolphthalein and at the 2nd end point with methyl orange, respectively.
These values of Va, Vb and Vc are also used in calculating the volumes of HCl
required for neutralizing Na2CO3 and NaHCO3 present in mixture. Thus, for the first
end point we need Vb –Va = V1 cm3
and for the second end point Vc–Va= V2 cm3
of
hydrochloric acid. Thus, V1 is the volume of HCl needed to titrate half of the
sodium carbonate initially present in the mixture (CO 2−
3 + H+
→ HCO −
3
) and 2V1
will be the volume of HCl needed to neutralise whole sodium carbonate in the
mixture (CO 2−
3 + 2H+
→ H2CO3). V2 is the volume of HCl required to neutralize
both sodium carbonate and sodium hydrogen carbonate in the given solution. On
subtraction of 2V1 from this volume, V2, we can find out the volume of HCl, V3,
needed to neutralise initially present sodium hydrogen carbonate in the
mixture,V3 = (V2 – 2V1).
The corresponding chemical reactions may be summarised as:
CO3
H
+
HCO3
(Na2CO3
) (HCl)
HCO3
H
+
H2CO3
(HCl)
end point with phenolphthalein,
Volume of HCl = V1
= Vb − Va
end point with methyl orange,
Volume of HCl = V2
= Vc − Va
(NaHCO3
+ HCO3 of
sodium
carbonate)
2− −
−
−
+
+
Using volume V1, V2, V3 and molarity equations we can calculate amount of sodium
carbonate and sodium hydrogen carbonate in the mixture.
Using reaction of Eq. 1.3, we can write molarity equation for the half
neutralisation of sodium carbonate, here sodium carbonate reacts with
hydrochloric acid in 1:1 molar ratio to gives hydrogen carbonate. Thus the
molarity equation for this step will be:
HCl HCl
Na2CO3 Na2CO3
M V
M V
=
1
1
i.e. Na2CO3 Na2CO3 M V = MHClVHCl C(1.6)
Where MNa CO32
the molarity of sodium is carbonate solution and Na2CO3
V is the
volume of sodium carbonate used in titration. MHCl is the molarity of
hydrochloric acid and VHCl is the volume of the hydrochloric acid used in the
titration of sodium carbonate up to hydrogen carbonate stage.
As indicated by the final reaction (Eq. 1.5), complete sodium carbonate reacts with
hydrochloric acid in 1:2 molar ratios. Hence, molarity equation can be written as
HCl HCl
Na2CO3 Na2CO3
M V
M V
=
2
1 BCHCL-132 Chemistry Lab I
36
i.e. 2 Na2CO3 Na2CO3 M V = MHClVHCl C(1.7)
Where MNa CO32
is the molarity of sodium carbonate solution and Na2CO3
V is the
volume of sodium carbonate used in titration. MHCl is the molarity of
hydrochloric acid and VHCl is the volume of the hydrochloric acid used in the
complete titration of sodium carbonate (i.e. volume of HCl used up to methyl
orange end point for standard sodium carbonate solution)). Eq. 1.7 will be
used for the calculation of the strength of hydrochloric acid in the titration of
standardisation of hydrochloric acid.
Chemical reaction of the neutralisation of sodium hydrogen carbonate with
HCl can be written as:
HCO −
3 + H+
→ H2CO3 ...(1.8)
Here sodium hydrogen carbonate reacts with hydrochloric acid in 1:1 molar ratios.
Hence, molarity equation can be written as
M NaHCO3
V NaHCO3
= MHClVHcl ...(1.9)
Where MNaHCO3
is the molarity of sodium hydrogen carbonate solution and NaHCO3
V
is the volume of sodium hydrogen carbonate used in titration.MHCl is the
molarity of hydrochloric acid and VHCl is the volume of the hydrochloric acid
used in the titration.
Hydrochloric acid used for titration of the mixture solution is not a primary standard.
Therefore, before using hydrochloric acid for the titration, it should be standardised
with a suitable primary standard, preferably sodium carbonate. The reaction
between sodium carbonate and hydrochloric acid is shown by Eq .1.5 and molarity
equation used for calculating the molarity of hydrochloric acid is shown by Eq. 1.7.
End point of the titration is detected with methyl orange indicator.
Before proceeding further, answer the following SAQs.
SAQ 1
Suggest whether aqueous solutions of the following substances are acidic, basic or
neutral.
a) NaCN; b) NaCl; c)CH3COONa; d) NaHCO3; e) K2CO3
SAQ 2
Predict the number of pH breaks or sharp pH change(s) which will be observed for
the following titrations:
a) CH3COOH –NaOH
b) NaHCO3 –HCl
c) K2CO3 – HCI Experiment 1 Determination of Sodium Carbonate and sodium Hydrogen Carbonate Present in a Mixture
37
SAQ 3
On the basis of Fig. 1.4 given below, suggest suitable indicators for the titration of
ethanedioic acid or oxalic acid (H2C2O4) against strong base.
1 Equivalence point st
2 Equivalence point nd
7
14
PH
Volume of strong base added in cm 3
H C O 2 2 4
HC O2 4
-
C O2 4
2-
Fig. 1.4: Titration of ethanedioic acid (oxalic acid) with strong base like NaOH.
1.3 REQUIREMENTS
You will need the following apparatus and chemicals for this experiment.
Apparatus Chemicals
Beaker (250 cm3
) 2 Hydrochloric acid
Burette (50 cm3
) 1 Methyl orange
Burette stand with clamp 1 Phenolphthalein
Conical flask (250 cm3
) 1 Sample: Mixture of sodium carbonate and
sodium hydrogen carbonate or baking
soda
Funnel 1 Sodium carbonate (AR grade)
Pipette (20 cm3
) 1
Volumetric flasks (250 cm3
) 2
Weighing bottle 1
Solutions provided
Sample solution: Prepare a sample solution by dissolving a mixture of
sodium carbonate and sodium hydrogen carbonate (8.5 g Na2CO3+5.4 g
NaHCO3) or commercial baking soda (15 g) in 2 dm3 distilled water. BCHCL-132 Chemistry Lab I
38
Phenolphthalein indicator solution: It is prepared by dissolving 0.1 g of
the reagent in 80 cm3
of ethanol and adding adequate distilled water to
make it 100 cm3
. If a precipitate is formed, it is filtered.
Methyl orange indicator solution: It is prepared by dissolving 0.1 g of free
acid/sodium salt of the indicator in 80 cm3
of distilled water and adds 20 cm of
ethanol to make it 100 cm3
.
Hydrochloric acid solution (~ 0.1 M): This solution is prepared by taking 8.5
cm3
conc. HCI (37%) in a 1 dm3
volumetric flask and diluting the acid up to the
mark with distilled water.
1.4 PROCEDURE
First collect 0.1 M hydrochloric acid in a 250 cm3
beaker. Since hydrochloric
acid is a secondary standard, you have to standardise it by titrating it against a
primary standard, Na2CO3 in this case.
1) Standardisation of hydrochloric acid:
i