# 1. 1.” (Source: http://www.sciencebuddies.org ) Equation 1: q

1.
INTRODUCTION

1.1  BACKGROUND

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“Have you ever used a hot pack to warm your
hands or a cold pack on an injury? How can something produce heat or cold
without any microwaving or refrigeration involved? The answer is: chemistry.
Chemical reactions that produce heat are called exothermic. The burning of gasoline in a car engine is an
example of an exothermic reaction.

Reactions that are accompanied by the absorption
of heat are called endothermic. As
an example of an endothermic reaction, when the chemical ammonium nitrate is dissolved in
water, the resulting solution is colder than either of the starting materials.
This kind of endothermic process is used in instant cold packs. These cold
packs have a strong outer plastic layer that holds a bag of water and a
chemical, or mixture of chemicals, that result in an endothermic reaction when
dissolved in water. When the cold pack is squeezed, the inner bag of water
breaks and the water mixes with the chemicals. The cold pack starts to cool as
soon as the inner bag is broken, and stays cold for over an hour. Many instant
cold packs contain ammonium nitrate. Ammonium nitrate is a white crystalline
substance. When it is dissolved in water, it splits into positive
ammonium ions and
negative nitrate ions. In the process of dissolving the crystal, the water
molecules “donate” some of their energy. As a result, the water cools
down. How much heat energy is
“lost” when ammonium nitrate dissolves in water? You can measure the
amount of heat that is involved using Equation 1.” (Source: http://www.sciencebuddies.org )

Equation 1:

q
=

c
m (T1 -T2)

q =
energy, measured in joules (J)
c = heat capacity, measured in joules
per gram per degree Celsius, J/(g°C)
m =
mass of solution, measured in grams (g)
T1 =
initial temperature, in degrees Celsius
T2 =
highest/lowest temperature, in degrees Celsius
J = joules (J), unit of energy
°C =
degrees Celsius

“Equation 1 states that “the amount of heat
energy that is lost when water changes from temperature T1 to the lower
temperature, T2, equals the difference in the two temperatures, times the heat
capacity, times the mass of the solution.”

The heat capacity of a substance tells you how
much the temperature will change for a given amount of energy exchanged. For
water at 25°C, the heat capacity is 4.18 J/(g°C).” (Source: http://www.sciencebuddies.org )

Equation 2:

c
(water) =

4.18
J
(g°C)

c =
heat capacity, measured in joules per gram per degrees Celsius, (J/g°C)
J =
joules (J), unit of energy
g =
grams (g) of water
°C =
degrees Celsius

“Equation 2 says that
“the heat capacity of water is 4.18 joules per gram of water per degree
Celsius.” What this means is that if you add 4.18 J of heat energy to 1 g
of water, its temperature will increase by 1.0°C. Substances other than water
have different heat capacities.” (Source: http://www.sciencebuddies.org )

Throughout
my life, I did once use the cold pack and once use the hot pack. It is all
because I had sprained my ankle when I was about to lift a heavy box and it
swollen right after the accident. My family got a cold pack at home since my
brother is a practical Physical Education teacher just somewhere near my house.
Even though the cold pack did not cure the swollen totally, but it reduced the
inflammation faced by me.

There
was once again, my leg’s muscle suddenly got stiff and I could not move even
inch. All I could do was just held the pain while waiting for my brother to get
a hot pack. It was a relieve though for having him as a brother. After some
minutes, the stiffness reduced and eventually lost.

What
I knew about those two packs is they are made from calcium chloride and
ammonium nitrate. But I did learn that there are many other chemicals that
produced exothermic and endothermic reactions. By investigating the effect of
other chemicals on the temperature of water, I can identify any other chemicals
that can be use for these reactions.

1.2 RESEARCH QUESTION

How
does the effect on temperature of solution over time determine the suitability
of a salt to be used in cold pack or hot pack while the temperature changes
were recorded constantly for 4 minutes?

1.3 AIM

To
investigate the relationship between the effects on temperature of solution
over time and the suitability of a salt to be used in cold pack or hot pack.

1.4 HYPOTHESIS

In making hot pack, the salts that
are suitable to be used are the one which can increase the temperature of
solution after being mixed with water in short time, as well as maintaining the
temperature for a long time or having only little decreases from time to time.
The maximum temperature reached by the salt solution should also relevant to be
hold by skin to avoid scald.

As for cold pack, the salts used in
making it can cause decrease in temperature of solution in short time after
dissolve in water. The temperature able to maintain for sometimes and increases
only bit by bit for minutes until it reaches the room initial temperature
again.

2.      METHODOLOGY

2.1  SAFETY PRECAUTION

1.      Wear indirectly vented chemical-splash goggles and
chemical-resistant gloves and apron while in the laboratory.

2.      Handle all glassware with care.

3.      Wash your hands with soap and water before leaving
the laboratory.

2.2  VARIABLES

VARIABLE

HOW TO CONTROL

a)      Independent:
The
type of salts

Use
different type of salts which are ammonium nitrate, ammonium chloride,
calcium chloride, sodium chloride, magnesium sulphate and copper(II) sulphate

b)      Dependent:
The
temperature of solution over time

Measure
the temperature of the solution for every 15 seconds.

c)      Controlled:
I)       The
weight of salt

II)    The
volume of water

III) The
time taken

Use
same weight of salt which is 5 g for each trial.
Mix
salt with 25mL of distilled water for each trial.
Record
the temperature changes for 4 minutes for each trial.

2.2.1
APPARATUS
AND MATERIALS

APPARATUS

MATERIALS

Distilled water

Copper (II) sulphate

Spatula

Magnesium sulphate

25mL measuring cylinder ± 0.25

Dehydrated calcium
chloride

Digital stopwatch, ± 0.01s

Sodium chloride

Digital weighing scale, ± 0.01g

Ammonium chloride

Thermometer, ± 0.25º

Ammonium nitrate

Calorimeter with lid

Small bowl

2.2.2
PROCEDURE

1.
Measure 20mL of distilled water using measuring cylinder.

2.
Pour the measured distilled water into the calorimeter.

3.
Measure the initial temperature of the water using thermometer and
record it.

4.
Put small bowl on the digital weighing scale.

5.
Zero the scale and slowly put down 5g of copper
(II) sulphate using the spatula.

6.
Pour the salt into the water-filled calorimeter and immediately
close it as well as start the stopwatch.

7.
Stir the solution gently throughout the process.

8.
For each 15 seconds in 4 minutes period, measure the solution’s
temperature using the thermometer and record the value.

9.      Dispose of the copper(II)
sulphate solution down the sink.

10. Repeat steps 1 to 8 four
more times to get five trials.

11. Calculate the average temperature
over times and the average amount of energy involved to be included in the
record.

12.  Repeat the steps 1 to 10
by using different salt which are ammonium nitrate,
ammonium chloride, dehydrated calcium chloride, sodium chloride and magnesium sulphate.

3.
RESULTS AND ANALYSIS

3.1  QUANTITATIVE

Type of salt

Average
temperature of solution over time (ºC)

Average
amount of energy involve (J)

0

15

30

45

60

75

90

105

120

135

150

165

180

195

210

225

240

Copper (II) sulphate, CuSO4

30.8

36.7

37.1

36.8

36.2

35.8

35.3

34.9

34.8

34.4

34.2

34.0

33.9

33.8

33.4

33.4

33.4

-526.68
± 21.07

Magnesium sulfate, MgSO4

29.8

38.5

39.6

39.5

39.1

39.1

38.9

38.7

38.4

38.0

37.5

37.4

36.9

36.7

36.4

36.2

35.9

-819.28
± 32.77

Dehydrated calcium chloride, CaCl­2

29.8

35.6

35.1

34.4

33.9

33.4

33.0

32.8

32.4

32.3

32.2

32.0

31.8

31.8

31.8

31.7

31.5

-484.88
±  19.40

­Sodium chloride, NaCl

28.6

27.8

27.8

27.7

27.7

27.7

27.7

27.7

27.9

27.9

28.1

28.1

28.1

28.1

28.1

28.1

28.2

75.24
± 3.01

Ammonium chloride, NH4Cl

27.9

19.1

19.1

19.2

19.8

20.6

21.0

21.4

22.1

22.5

22.9

23.1

23.3

23.6

23.7

24.0

24.2

735.68
± 29.43

Ammonium nitrate, NH4NO3

28.8

21.9

19.6

20.3

21.4

22.2

22.8

23.4

23.7

24.2

24.6

24.9

25.1

25.2

25.4

25.7

25.8

769.12
± 30.76

3.2  QUALITATIVE

Type of salt

Observation

Copper (II) sulphate, CuSO4

Solution turns from colourless to
blue
Blue precipitate formed

Magnesium sulfate, MgSO4

No
changes

Dehydrated calcium chloride, CaCl­2

White
precipitate formed

­Sodium chloride, NaCl

White
precipitate formed

Ammonium chloride, NH4Cl

No
changes

Ammonium nitrate, NH4NO3

White
precipitate formed

3.3  PROCESSING
AND ANALYSIS

3.3.1
Example of calculation

i.
Average temperature of solution
over time (ºC), copper(II) sulphate at 0s

=

=

= 30.8

ii.      Amount of
energy involve (J), copper(II) sulphate

= c m (T1 -T2)

= 4.18 × 20 × (30.8 – 37.1)

=
– 526.68

Where q =
energy, measured in joules (J), c = heat
capacity, measured in joules per gram per degree Celsius, J/(g°C), m =
mass of solution in grams (g), T1 = initial temperature in degrees Celsius, T2 =
highest/lowest temperature in degrees Celsius, J = joules (J) unit of energy and °C = degrees Celsius

iii.    Uncertainty, average amount of energy involved in copper(II)
sulphate

Percentage
of uncertainty

a.
Digital weighing scale =  × 100% = 0.2%

b.
Measuring cylinder = × 100% = 1.25%

c.
Digital stopwatch = 16(  × 100%) = 1.07%

d.
Thermometer, initial temp =  × 100% = 0.81%

e.
Thermometer, initial temp =  × 100% = 0.67%

Total percentage uncertainty = 0.2 + 1.25 + 1.07 + 0.81 +0.67 = 4%

Absolute uncertainty =  × 526.68J = ± 21.07 J

?, the
average amount of energy involved is (- 526.68 ± 21.07) J

3.3.2
Graph

Graph 1: Average temperature of
solution over time (exothermic)

Graph 2: Average temperature of
solution over time (endothermic)

3.4  DISCUSSION

For the
Graph 1, all three solutions from different types of salt had a rapid increase
for the first 15 seconds. Starting at 30 seconds period, the temperature of
solution of dehydrated calcium chloride decreased until the third minute before
had a plateau for 45 seconds and continued to decrease. As for solutions of
copper (II) sulphate and magnesium sulphate, they still had a little increment
of temperature after 15 seconds and reached its peak when the time is 30
seconds. Copper (II) sulphate solution’s temperature started to decline
throughout the experiment, except for the last 45 seconds, where they already
reached constant temperature. Different with magnesium sulphate solution which
the temperature of the solution decreased for 30 seconds after that and become
constant the next 30 seconds. From 90 seconds period, the temperature of the
solution kept on decreasing until the forth minute. All three types of salt
have the same trend of changes but among them, the temperature changes of magnesium
sulphate solution was the highest, followed by copper (II) sulphate solution and
the lowest was dehydrated calcium chloride solution.

While from the Graph 2, it can be
seen that the three salt solutions all reflect the trend occurs in Graph 1.
They decreased in temperature first before gradually increased again to reach
the room temperature. Sodium chloride solution was clearly shown to have very
small changes of temperature throughout the 4 minutes period. Ammonium nitrate
solution was continuously decreased in temperature for the first 30 seconds and
after that increased steadily till the experiment ended. For the temperature of
ammonium chloride solution, it declined sharply and been constant for 15
seconds. Then, the solution temperature started to increase gradually until the
forth minute.

The changes of temperature were
caused by the loss and gain of energy of the particles of the solution when
they reacted to each other to carry out the reactions. Every substance has
their own heat capacity, and every atom bonding forces are different. Because
of that, the amount of energy involved in a reaction and maximum temperature are
different as well.

4.      EVALUATION

4.1  LIMITATIONS
AND WAYS TO OVERCOME

LIMITATION

WAYS TO OVERCOME

1.      The heat may escape from the
calorimeter just before the lid is placed.

Close
the calorimeter immediately right after the salt is poured into it to
minimize the amount of heat escaped.

2.      Some salts need more water to
allow them to dissolve completely.

Carry out the experiment using much more distilled water
to ensure that all salts dissolved.

5.
CONCLUSION

To
identify if a salt is suitable to be used in making hot pack or cold pack, we
need to consider the increment or decrement of the temperature of solution
which reflect the amount of energy involved to complete the reaction.

As for exothermic reaction, all
three types of salt are suitable to be used in making hot pack since they show
increment in temperature when get dissolved in water. They also decrease slowly
by time which is good for hot pack in maintaining the temperature for a longer
time. But, for magnesium sulphate, the temperature it can reach almost 40°C and
the patient may not be able to stand such temperature. To overcome the problem,
a layer of insulator should be placed inside the hot pack to lower down the
temperature.

In making cold pack, the most
suitable salts to be used are ammonium nitrate and ammonium chloride because
both of them showed the same trend of temperature changes. They released energy
rapidly to decrease the temperature but take some time to back to room
temperature. The minimum temperatures are also suitable to be hold by our body.
Sodium chloride showed only very little changes in its temperature after being
mixed with water so it is not really suitable to be used in cold pack.
Therefore, the hypothesis was supported.