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PHYSICAL CHEMISTRY (LAB MANUAL)
EXPERIMENT 1
AIM : Determination of enthalpy of neutralization of hydrochloric
acid with sodium hydroxide.
LEARNING OBJECTIVES
After performing this experiment student will be able to:

explain the cause of change in thermal energy of a system during
any physical or chemical change;

correlate the change in thermal energy with heat gained or lost by
the system;

define the terms calorimetric formula, heat capacity of calorimeter,
water equivalent of calorimeter, enthalpy of neutralization;

distinguish
between
heat
capacity
and
water
equivalent
of
calorimeter;

explain why the enthalpy of neutralization of any strong acid with a
strong base is the same;

identify the apparatus and chemicals required for performing this
experiment;

set up the apparatus for determination of heat capacity of
calorimeter and enthalpy of neutralization;

record the data and draw the required plot to determine the initial
and final temperatures at the ‘time of mixing’;

calculate the desired quantities from the data collected;

identify the common sources of error and suggest how these can be
minimized.
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PHYSICAL CHEMISTRY (LAB MANUAL)
REQUIREMENTS
(a)
Chemicals and Solutions
Conc.HCl, NaOH, oxalic acid, phenolphthalein indicator.
(b)
Apparatus
Weighing bottle, three 250 mL standard flasks, 250 mL beakers, Dewar/
thermos flask, rubber cork (3” dia), glass stirrer, two 0.1oC thermometers,
stop watch / stop clock, burette, 10 mL pipette, 100 mL conical flask,
tripod stand, wire gauze.
THEORY
The enthalpy of neutralization is defined as the change in enthalpy
when one mole of H+(aq) ions from an acid is neutralized by one mole of
OH-(aq) ions from a strong base in their dilute solutions at a constant
temperature and under 1 bar pressure. The neutralization reaction
between a strong acid HA and a strong base BOH occurs as:
HA (aq) + BOH (aq)  BA (aq) + H2O(l)
It is observed that the enthalpy of neutralization is the same in almost all
cases of neutralization reaction between a strong acid and a strong base
and is found to be
– 57.2 kJ mol-1. The reason for this can be
understood as follows. All strong acids (HA), and bases (BOH) are
completely ionized in their aqueous solutions as:
HA (aq)

H+ (aq) + A- (aq)
BOH (aq)

B+ (aq) + OH- (aq)
The neutralization reaction can be rewritten as:
H+(aq) + A-(aq) + B+(aq) + OH-(aq)  B+(aq) + A-(aq) + H2O(l)
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PHYSICAL CHEMISTRY (LAB MANUAL)
Since B+(aq) and A-(aq) ions are common on both the sides of this
equation, they can be cancelled to obtain the equation.
H+(aq) +
OH-(aq)  H2O(l)
nHo = – 57.2 kJ mol-1
Since this reaction is common to all neutralization reactions between
strong acids and strong bases, their enthalpy of neutralization is also the
same.
This experiment is performed in the following two steps:
1(a)
Determination of heat capacity of calorimeter.
1(b) Determination of enthalpy of neutralization of HCl
and NaOH .
1(a). Determination of heat capacity of calorimeter.
The amount of heat absorbed or evolved during a thermochemical process
(chemical or physical) is carried out by measuring the change in
temperature. This method is called calorimetry and the apparatus used
is called calorimeter. The experimental set up is shown in Figure 1.
Fig 1: Calorimeteric set up
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PHYSICAL CHEMISTRY (LAB MANUAL)
Calorimetric Calculations
The amount of heat absorbed or evolved during a thermochemical process
(chemical or physical change) is carried out by measuring the change in
temperature in the calorimeter. Calculations are based upon the first law
of thermodynamics according to which the total energy of a system and
its surroundings in conserved. In an adiabatic calorimeter the heat
exchange with the surroundings is zero. Hence, the heat lost or gained
inside the calorimeter must balance each other
(i)
Heat lost or gained by water
It is calculated with the help of calorimetric formula
qwater = m x s x t
where
m=
mass of water taken
in the calorimetric vessel. Since density of
water is very nearly 1g mL-1 at room temperature, the volume of
water (or solution) in milliliter is numerically equal to its mass in
grams. The slight inaccuracy in omitting the presence of solutes is
negligibly small.
s=
Specific heat capacity of water. It is the amount of heat required to
raise the temperature of unit mass (1g) of water by 1oC. In joule
(SI unit) the specific heat capacity of water is 4.18 J oC-1 g-1 and in
calorie it is 1 cal oC-1 g-1.
t =
Change in temperature
=
(ii)
tfinal – tinitial
Heat lost or gained by calorimeter
The above formula cannot be used for the calorimeter used here (glass
beaker along with the stirrer and the thermometer). It is because the
glass is a thermal insulator and heat is not uniformly gained or lost by the
entire vessel.
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PHYSICAL CHEMISTRY (LAB MANUAL)
It is almost only that part of the beaker which is actually in contact with
the reacting system which absorbs or loses heat. Mass of this portion
cannot be determined. For this reason instead of using mass and specific
heat capacity, use is made of heat capacity
of calorimeter, C in the
calorimetric formula i.e.
qcalorimeter
=
C x t
Where
qcalorimeter is the heat lost or gained by calorimeter, C is its heat capacity,
which is the amount of heat required to raise the temperature of the
calorimeter by 1oC. t is the change in temperature.
In place of the term ‘heat capacity of calorimeter, the term ‘waterequivalent of calorimeter’ is also used. It is defined as the mass of
water in grams which requires the same amount of heat to raise its
temperature by 1oC as the calorimeter. The heat gained or lost may now
be calculated by the relation;
qcalorimeter
=
mwater
x
swater x
t
where mwater is the water equivalent of calorimeter. By comparing with the
earlier relation in terms of heat capacity C, it can be concluded that
C
=
mwater
x
swater
If calculations are made in calorie, the specific heat capacity of water is 1
cal oC-1 g-1 and
C
=
mwater
Thus, the heat capacity of water is numerically equal to the water
equivalent, if the heat absorbed or lost is measured in calories.
Yet another term used is ‘calorimeter constant’ which is defined as the
amount of heat required (in calories) to raise the temperature of
calorimeter by 1oC. Numerically it is equal to the water equivalent of
calorimeter and also to the heat capacity of calorimeter in calories.
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PHYSICAL CHEMISTRY (LAB MANUAL)
It may be noted here that the term ‘calorimeter’ here includes not only
the glass beaker but also the ancillary apparatus like thermometer and
stirrer.
Therefore, in all calorimetric measurements the same ‘set’ of apparatus
should
be
used
which
includes
the
glass
beaker,
glass
stirrer,
thermometer and even the Dewar flask/ thermos flask since it also
absorbs some amount of heat.
Another important point to be noted is that the given experimental value
of heat capacity (or water equivalent or calorimeter constant) is valid for
the same total volume of solution (or that of cold water + hot water) for
which it was determined. If in another experiment, the total volume of
solution to be used is different, heat capacity must be determined afresh
using this new total volume of hot water and cold water.
The heat capacity of calorimeter is determined by supplying a known
amount of heat to the calorimeter by mixing known quantities of two
samples of water kept at different temperatures. From the known
quantities of water and the initial and final temperatures, the heat
capacity of the calorimeter can be calculated. The following procedure is
followed to carry out this process:
PROCEDURE
1. Take a clean and dry 250 mL glass beaker and place it in a thermos
flask on a 3” diameter rubber cork which is kept at the bottom of
the flask.
2. Measure 100 mL distilled water into it with the help of a burette or
a 50 mL pipette.
3. Cover of the thermos flask with a lid with two holes in it through
which a glass stirrer and a 0.1oC thermometer are fixed.
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PHYSICAL CHEMISTRY (LAB MANUAL)
4. Take a 250 mL glass beaker, rinse it with distilled water and then
measure 100 mL distilled water into it.
5. Keep this beaker on a tripod stand over a wire gauze. Fix a 0.1oC
thermometer in it and note down the temperatures of the water.
6. Heat the water till its temperature rises by 8o-10oC. Now remove
the beaker from the tripod stand and keep it on the laboratory
table.
7. Start a stop clock and record the temperatures of the cold water in
the thermos flask and the hot water kept out side with constant
stirrring, alternately every 15 seconds. Continue recording the
temperature till 15 seconds before the time of mixing (i.e. upto 4
min 45 s, if the time of mixing has been fixed as 5 min).
8. Remove the thermometer from the beaker containing hot water.
Hold this beaker in your hand near the mouth of thermos flask.
9. At about 3-4 seconds before the time of mixing (5 min), remove
the lid of the thermos flask partly and start pouring the hot water
gently into the beaker kept in the thermos flask. When completely
poured, quickly replace the lid and start stirring the water in the
thermos flask gently but continuously.
10.After 30 seconds from the time of mixing (i.e. from 5 min 30 s)
start recording the temperature of mixed water at regular intervals
of 30 seconds. Take 8-10 readings and record them in the table.
11.Plot a graph of temperature versus time (fig. 2). Mark the time of
mixing on the graph by drawing a vertical line.
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PHYSICAL CHEMISTRY (LAB MANUAL)
Fig 2. : Temperature versus time graph for determination of heat
capacity of calorimeter
12.Join the points for cold water, hot water and mixture by three
straight lines and extend them to the vertical line representing the
time of mixing as shown in
figure 2.
13.From the points of intersections of the three lines, note down the
temperatures of (i) cold water, tc, (ii) hot water, th and (iii) mixture,
tm.
14.Calculate the heat capacity of the calorimeter as given later.
15.Repeat the experiment one more time to obtain second value of the
heat capacity
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PHYSICAL CHEMISTRY (LAB MANUAL)
VI SOURCES OF ERRORS AND PRECAUTIONS
(A)
Handling of Chemicals and Apparatus
1. Dewar/ thermos flask is very fragile and should be handled carefully
avoiding jerks etc.
2. Always keep a rubber cork at the bottom of the flask and keep the
glass beaker on it.
3. Thermometer should be clamped carefully such that its bulb does
not touch the bottom or side of the beaker.
4. While stirring, the stirrer should not strike the beaker (bottom or
side) or the bulb of the thermometer.
5. Before lifting the beaker containing hot water touch it briefly to
ascertain that it can be lifted comfortably with bare hands. Else use
a piece of cloth or paper to hold it.
(B)
Temperature
6. While recording the temperature, bring your eye at the level of
mercury in order to avoid parallax error.
7. Since two thermometers are being used in this experiment, their
readings should match. It can be checked by recording the
temperature of the same sample of water by both (Preferably use
two separate samples of water with at least 10oC difference in their
temperatures and check the thermometers in them).
In case, there is a difference in their readings, the thermometer
fixed in the thermos flask may be taken as standard and a label
pasted on it. Correction factor should be applied to the readings of
the other thermometer. Only the corrected readings should be
plotted on the graph paper.
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PHYSICAL CHEMISTRY (LAB MANUAL)
8. The difference between temperature of cold and hot water should
not be more than 8o-10oC.
9. While recording the temperature, the water must be stirred
constantly but gently.
(C)
Errors in measurements and calculations
10.The beaker is which water is to be heated (from which a liquid has
to be poured into the other beaker) should be rinsed with distilled
water before taking a known volume of water in it. This will take
care of water that remains sticking to the sides after pouring it into
the other beaker.
11.The time of mixing is the time when one half of hot water has been
poured. A rough estimate of the time taken for pouring can be
made before start of the experiment by pouring 100 mL water from
one beaker into the other gently. During the experiment, the
pouring should be started one half of this time before the predecided time of mixing (5 min).
12.After pouring the hot water, stirring should be done gently but
continuously.
OBSERVATIONS
I.
Masses of cold and hot water
Mass of cold water taken = 100 g
Mass of hot water taken = 100 g
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PHYSICAL CHEMISTRY (LAB MANUAL)
II. Temperature - Time Data
Temperature /
Time
/ min
Time
Cold
water
Hot water
Recorded
Corrected
/ min
Temperature
of Mixture
/ oC
CALCULATIONS
Density of water  1 g mL-1
Mass of cold water = 100 g
Temperature of cold water at the time of mixing (from graph) = tc =
Mass of hot water = 100 g
Temperature of hot water at the time of mixing (from graph) = th =
Temperature of mixture at the time of mixing (from graph) = tm =
Specific heat capacity of water = 4.18 J 0C-1 g-1
Let the heat capacity of calorimeter be= C J
0
C-1
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PHYSICAL CHEMISTRY (LAB MANUAL)
Heat lost by hot water = m x s x t
= (100g) x (4.18 J oC-1g-1) x (th – tm) J
Heat gained by cold water = m x s x t
= (100 g) x (4.18 J 0C-1g-1) x (tm - tc) J
Heat gained by calorimeter= C x t
= C x (tm – tc)
Heat gained by cold water
J
and calorimeter = Heat lost by hot water
{100 x 4.18 x (tm - tc)} + {C x (tm – tc)} = 100 x 4.18 x (th – tm)
(100 x 4.18 + C) (tm - tc) = 100 x 4.18 x (th - tm)
100 x 4.18 ( t h  t m ) 
o 1
C
  {100 x 4.18} J C
(t m  t c )


Calculate C using data from 2 sets as C1 and C2.
Average heat capacity =
C1  C 2 0 -1
J C
2
Result: The average heat capacity of calorimeter for 200 mL of water
(100 mL cold water +100 mL hot water) was found to be …. J 0C-1.
DISCUSSION
Discuss the possible reasons for difference in the values of heat capacity
obtained from two sets.
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PHYSICAL CHEMISTRY (LAB MANUAL)
EXTENSION
1. A major source of error in calorimetric measurements is the loss of
heat by radiation. How can the procedure be modified to minimize
it, while using the same apparatus?
2. Suggest alternatives to Dewar/Thermos flask.
TROUBLE SHOOTING
Time of Mixing
In case the mixing of hot and cold water was not started at the
stipulated time as given in the procedure, make a note of the time when
about one-half of mixing had been done. Take this as the time of mixing
and draw a vertical line on the graph accordingly.
1(b). Determination
of
the
Enthalpy
of
Neutralization
of
Hydrochloric Acid with Sodium Hydroxide.
It can be done by the following procedure:
(1) Take the same calorimeter setup (250 mL beaker, thermos flask,
thermometer
and
stirrer
whose
heat capacity
was
determined
in
experiment No.1(a).
(2) Measure 100 mL 1 M HCl solution in the 250 mL beaker used as
calorimeter and keep it inside the thermos flask.
(3) Fix a glass stirrer and 0.10C thermometer in it with the help of the lid
of the thermos flask.
(4) Rinse a 250 mL beaker with 1 M NaOH solution and then measure 100
mL of this solution in it. Fix the second 0.10C thermometer in it.
(5) Start the stop clock and record the temperatures of HCl and NaOH
solutions alternately every 15 seconds till the last reading is taken 15
seconds before the time of mixing.
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PHYSICAL CHEMISTRY (LAB MANUAL)
(6) Remove the thermometer from NaOH solution, open the lid of thermos
flask partly and pour it into the beaker containing HCl solution such that
one half of mixing would occur by the time of mixing.
(7) Replace the lid as soon as the pouring is complete and start stirring
the mixture solution in the thermos flask gently but continuously.
(8) Starting from 30 seconds after the time of mixing, record the
temperature after every 30 seconds for next 4-5 minutes.
(9) Plot graphs of temperatures of HCl, NaOH and their mixture solutions
versus time. Draw a vertical line at the time of mixing and determine the
temperatures of the acid (ta), base (tb) and mixture tm at the time of
mixing from the graph (Fig. 3)
(10) Repeat the experiment (steps 10-17) once more time to obtain
second value of enthalpy of neutralization using the same solutions
Fig 3. Temperature versus time graph for determination of
enthalpy of neutralization of HCl and NaOH
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PHYSICAL CHEMISTRY (LAB MANUAL)
Sources of Error and Precautions
1. Solutions of HCl
and NaOH should be taken with the help of
burette or pipetted carefully. Caution should be taken while using
the pipette
that
you don’t suck them into your mouth. If it
happens, do not swallow the solution, spit it out in the sink and
wash out mouth repeatedly with water and consult a doctor as soon
as possible.
2. If necessary, apply correction factor to the readings of thermometer
used for recording the temperature of the solution kept outside the
thermos flask.
OBSERVATIONS AND CALCULATIONS
In the neutralization reaction the heat is liberated which raises the
temperature of the calorimeter and the resulting solution
Total heat liberated, q = - (heat gained by calorimeter + heat gained by
water)
Heat capacity of calorimeter = C J oC-1 (as determined in part I)
Mass of water in the mixture of acid and base solutions = 200 g
( Mass of 200 mL water  200 g)
Specific heat capacity of water = 4.18 J oC-1 g-1
Rise in temperature = tfinal – tinitial
Where tfinal = tm, the temperature of mixture at the time of mixing
tifinal =

ta  tb
, the average initial temperature at the time of mixing.
2
t = ( t m 
ta  tb
)
2
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PHYSICAL CHEMISTRY (LAB MANUAL)
Heat gained by calorimeter = C x t
Heat gained by water = m x s x t
= 200 g x 4.18J oC-1 g-1 x t
Total heat gained by water and calorimeter
= (200 x 4.18 + C) (t)
Heat evolved during neutralization = q = - (200 x 4.18 + C) t
Enthalpy of neutralization is the heat evolved when 1 mole of H+ (aq) ions
from the acid is neutralized by 1 mole
of OH- (aq) ions from the base.
This can be obtained by calculating the heat evolved, q, when one litre
volume of 1 M HCl is neutralized by 1 litre volume of 1 M NaOH.
nH  
q 1000
q 1000

 10 q
Volume of 1 N acid or base used
100
or nHo = 10 q
Calculate nHo for both the sets and calculate the average value.
RESULT
The enthalpy of neutralization of HCl and NaOH was found to be
… kJ mol-1
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