CONTENT

TITLE PAGE

ABSTRACT 2

INTRODUCTION: 1.1 INTRODUCTION 1.2 OBJECTIVE

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METHODOLOGY 6

RESULTS AND DISCUSSION 7

CONCLUSION 8

REFERENCES 8

APPENDICES 9

Abstract

1

After doing this experiment, we are able to determined the specific heat capacity of a

metal, copper. We also had determined the quantity and the direction of heat flow for the

dissolution of salt. Based on our group result, we got 0.0759 Jg-1°C-1 as the specific heat of

copper while the direction of heat flow of the dissolution salt is exothermic. The methods

that we use are weighing, heating, and recording. However, there are some problems

occurred. Our result is not 100% accurate but nearly. The inaccurate result is due to some

errors while doing the experiment. For example, we do not shut the fan off. Thus, some of

the copper powders are dispersing into the air. Moreover, the air pressure also effects the

decreasing temperature of the copper.

As conclusion, we successfully manage to determine the specific heat of a copper and the heat flow for the dissolution of the salt.

1.0 INTRODUCTION

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1.1 Introduction

A calorimeter is an apparatus for measuring the quantity of heat energy released or

absorbed during a process. Since there are many processes that can be studied over a wide

range of temperature and pressure, a large variety of calorimeters have been developed.

q = (specific heat) ( m )( Dt) = S.H. ( m)( Dt) 

Non-isothermal calorimeters measure the temperature change that occurs during the

process. An aneroid-type non-isothermal calorimeter is normally constructed of a material

having a high thermal conductivity, such as copper, so that there is rapid temperature

equilibration. It is isolated from its surroundings by a high vacuum to reduce heat leaks. This

type of calorimeter can be used for determining the heat capacity of materials when

measurements involve low temperatures. Aneroid-type non-isothermal calorimeters have

also been developed for measuring the energy of combustion for small samples of rare

materials.

With most non-isothermal calorimeters, it is necessary to relate the temperature rise

to the quantity of energy released in the process. This is done by determining the

calorimeter constant, which is the amount of energy required to increase the temperature of

the calorimeter itself by 1°. This value can be determined by electrical calibration or by

measurement on a well-defined test system. For example, in bomb calorimeter the

calorimeter constant is often determined from the temperature rise which occurs when a

known mass of a very pure standard sample of benzoic acid is burned.

Isothermal calorimeters make measurements at constant temperature. The simplest

example is a calorimeter containing an outer annular space filled with a liquid in equilibrium

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with a crystalline solid at its melting point, arranged so that any volume change will displace

mercury along a capillary tube. The Bunsen ice calorimeter operates at 0°C (32°F) with a

mixture of ice and water. Changes as a result of the process being studied cause the ice to

melt or the water to freeze, and the consequent volume change is determined by

measurement of the movement of the mercury meniscus in the capillary tube. While these

calorimeters can yield accurate results, they are limited to operation at the equilibrium

temperature of the two-phase system. Other types of isothermal calorimeters use the

addition of electrical energy to achieve exact balance of the heat absorption that occurs

during an endothermic process.

All calorimeters consist of the calorimeter proper and a jacket or a bath, which is

used to control the temperature of the calorimeter and the rate of heat leak to the

environment. For temperatures not too far removed from room temperature, the jacket or

bath contains liquid at a controlled temperature. For measurements at extreme

temperatures, the jacket usually consists of a metal block containing a heater to control the

temperature. With non-isothermal calorimeters, where the jacket is kept at a constant

temperature, there will be some heat leak to the jacket when the temperature of the

calorimeter changes. It is necessary to correct the temperature change observed to the

value it would have been if there were no leak. This is achieved by measuring the

temperature of the calorimeter for a time period both before and after the process and

applying Newton's law of cooling. This correction can be avoided by using the technique of

adiabatic calorimeter, where the temperature of the jacket is kept equal to the temperature

of the calorimeter as a change occurs. This technique requires more elaborate temperature

control, and its primary use is for accurate heat capacity measurements at low

temperatures.

In calorimetric experiments it is necessary to measure temperature differences

accurately; in some cases the temperature itself must be accurately known. Modern

calorimeters use resistance thermometers to measure both temperatures and temperature

differences, while thermocouples or thermistors are used to measure smaller temperature

differences.

Heat capacities of materials and heats of combustion are processes that are routinely

measured with calorimeters. Calorimeters are also used to measure the heat involved in

phase changes, for example, the change from a liquid to a solid (fusion) or from a liquid to a

gas (vaporization). Calorimetry has also been applied to the measurement of heats of

hydrogenation of unsaturated organic compounds, the heat of dissolution of a solid in a

liquid, or the heat change on mixing two liquids.

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1.2 Objective1) To determine the specific heat of a metal

2) To determine the quantity and direction of heat flow for the dissolution of a salt

2.0 METHODOLOGY

Materials

Calorimeter, thermometer, beaker 400 ml, test tube, bunsen burner, wire gauze, graduated

cylinder, three finger clamp, copper, natrium tiosulphate

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Procedure:

A) Specific Heat of a Metal

1. 10 g of copper is weighed and transferred it into a dry test tube. The test tube is

placed in a 400 ml beaker. Then, the beaker is filled with water until it is well

above the level of the metal sample in the test tube. The water is boiled using a

heater and maintain this temperature for at least 5 minutes so that the metal

reaches thermal equilibrium with the water. The temperature is recorded.

2. The mass of calorimeter is weighed. By using a graduated cylinder, 20.0 ml of

water is added into the calorimeter. The combined mass of the calorimeter and

water are determined. The thermometer is secured with a clamp and position it

below the water surface. 5 minutes is allowed for the system to reach thermal

equilibrium. The temperature is recorded over 60 second intervals.

3. The test tube from the boiling water is removed and only the copper is quickly

transferred into the calorimeter. The lid is replaced and the content is stirred

gently. The water temperature is recorded as a function of time (about 30 second

intervals) for 3 minutes.

4. The temperature (y axis) versus time (x axis) graph is plotted. ΔT is determined

from our curve and then do the calculations indicated on the report sheet.

A) Entalphy (Heat) of Solution for the Dissolution of a Salt

1. The mass of the dry calorimeter is weighed. Using graduated cylinder, 20.0 ml of

distilled water is added into the calorimeter and temperature is recorded. The

temperature is recorded for 60 seconds with 15 seconds interval. The combined

mass of the calorimeter and water are reweighed. After that, the thermometer is

positioned below the water surface.

2. 5.0 g of the natrium thiosulphate, Na2S2O3 salt is weighed.

3. The salt is added to the calorimeter, the lid is replaced and stirred gently until the

salt dissolves. The temperature and time are read and recorded at 15 seconds

intervals for 3 minutes.

4. The temperature versus time curve is constructed and T is determined.

3.0 RESULTS AND DISCUSSIONS

(A) Specific Heat of Metal Temperature of copper = 81.5 °C Mass of calorimeter = 6.8g Mass of calorimeter + water = 25.5g Mass of water = 18.70g

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Time (s) Temperature of water in calorimeter (°C)60 29.0120 29.0180 29.0240 29.5300 29.5

Time (s) Temperature of water + cooper (°C)30 30.060 30.090 30.0120 30.5150 30.5180 30.5

∆Twater = Tfinal(water) – Tintial(water)

= (29.5 – 29.0) °C

= 0.5 °C

∆Tcopper = Tfinal(copper) – Tintial(copper)

= (30.0 – 81.5) °C

= - 51.5 °C

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Heat lost by copper = - (Heat gained by water)

Specific heat (copper) x mass (copper) x ∆Tcopper = - [Specific heat (water) x mass (water) x

∆Twater ]

q x 10 x (- 51.5) = - (4.184 x 18.7 x 0.5)

q = 0.0759 Jg-1°C-1

The specific heat for copper is 0.0759 Jg-1°C-1.

Using the q = mcΔT equation, the heat of copper determined. The result we get is far from

the real specific heat capacity of copper. This is due to some errors during handling the

experiment. Which are, we open the lid of calorimeter too long and this might effects the

temperature of calorimeter. Besides, when we measured the temperature of thermometer,

the bulb of the thermometer is touched the surface of coffee cup. This situation has affects

the reading of the thermometer taken. Parallax error also occur during the thermometer

reading is taken. We should have ensured that our eyes are perpendicular with the

thermometer scale. Apart from that, there are some errors when the weight of copper is

taken.

The percentage deviation of actual specific heat capacity of copper with the result we

get is as below:

(0.385-0.0759) x 100% = 80.29%

0.385

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(B) Enthalpy of Solution for the Dissolution of a Salt Mass of dry calorimeter = 6.7g Mass of calorimeter + water = 21.6g

Time (s) Temperature of water in calorimeter (°C)15 2830 2845 2860 28

Time (s) Temperature of water + dissolve salt (°C)15 2330 2345 2360 2475 2490 24105 24120 24135 24150 24165 24180 24

Based on the graph,∆T = gradient of graph

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=

= 15

Heat lost by water = - (Heat absorbed for the dissolution of Na2S2O3)

Heat lost by water, QH2O = Mass (water) x Specific heat (water) x ∆T

= 14.9 x 4.184 x (15)

= 935.12 J

Heat absorbed for the dissolution of Na2S2O3 = 935.12 J

For experiment B, the result that we get is 935.12J. While handling this experiment B, there

are also some errors that occur. One of it is, the Na2S2O3 did not dissolve in water thoroughly

even we had stir it. Besides, the other errors that occur during this experiment B is same as

what we had in experiment A which IS the lid of the calorimeter is opened too long. This has

affected the reading of thermometer.

4.0 CONCLUSION

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At the end of this experiment, the objectives are achieved. The specific heat flow of copper is 0.0759 Jg-1°C-1 and the quantity of the dissolution is 935.12 J. From the graph we sketch, the reaction show the released of heat, exothermic.

EXERCISES:

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1. What is the difference between specific heat and heat capacity? What are the units

for these two quantities? Which is the intensive property and which is the extensive

property?

Specific heat capacity is the amount of heat needed to raise the temperature of 1 gram

of any substance by 1°C at constant pressure. And the units for specific heat capacity

is Jg-1°C-1 or calg-1°C-1. Specific heat capacity is an intensive property.

While, heat capacity is the amount of heat required to change a sample of substance

by 1°C. The units for heat capacity is Jmol-1 °C-1 and it is an extensive property.

2. A 20.94 g sample of an unknown metal is heated to 99.4oC in a hot water bath until

thermal equilibrium is reached. The metal is quickly transferred to 100 ml of water

at 22.0 oC contained in a Styrofoam cup. The thermal equilibrium temperature of

the metal plus water mixture is 24.6 oC. What is the specific heat of the metal?

qsystem = 0

q = mcΔT - CcalΔT

= (120.94)(4.184)(2.2)

= 1113.23

qsystem = 1113.23 + [(120.94)(א)(-74.8)]

[(74.8-)(א)(120.94)] + 1113.23 = 0

א 9046.31+ = 1113.23

=0.123 J/g°C

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3. Magnesium metal reacts with hydrochloride acid according to the following

equation.

Mg (s) + 2HCl (aq) Mg Cl2 (aq) + H2 (g)

When 0.425 g of magnesium was added to 150.0ml of 1.00 M HCl in a coffee-cup

calorimeter the temperature of the solution increased from 24.5 oC to 35.3 oC.

Given that the heat capacity of the calorimeter is 125 J/oC and that the density of

the HCl solution is 1.00 g/ml, calculate the heat released per mole of magnesium.

q = mcΔT

= (0.425)(125)(10.8)

= 573.75

let,

0.0175 mole ≡ 573.75

1 mole ≡ ?

So, 1 mole = 32785.71

REFERENCES

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William L.Masterton, Cecile Nespral Hurley. Chemistry principles and Reactions.5th edition.

Belmont,USA: Thomson

John W.Moore, Conrad L.Stanitski and Peter C.Jurs. Chemistry the Molecular Science. 2nd

edition. Belmont.USA: Thomson

Melanie M. Cooper. Cooperative Chemistry Laboratory Manual. 1st edition. Clemson

University: McGraw-Hill INTERNATIONAL EDITIONS

Ignacio Tinoco Jr, Kenneth Sauer, James C. Wong and Joseph D. Puglisi. Physical Chemistry

(Principles and Applications in Biological Sciences). 4th editon. Upper Saddle River, New

Jersey: Prentice Hall

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