Specific Heat Capacity Measurement and Calibration (Lab Report)

Introduction
The objective of this document is to introduce the learning outcomes for ‘Specific Heat Capacity
Measurement and Calibration’ laboratory exercise and to provide guidance for the required
laboratory work and report.
The heat capacity is a physical quantity that can be measured as the ratio of the heat
added/removed from an object to an object’s temperature change caused by the energy change.
In industry, the heat capacity is one of the key components in design and optimisation of every
process and technology for materials physical or chemical treatment. The SI unit of heat capacity
is joule per Kelvin, [J K-1
]. Heat capacity is an extensive property of materials as it is a function of
the size of a sample. In practical calculation, for majority of experimental and theoretical
purposes, the use of the intensive properties is more convenient. Similarly, the SI unit of energy is
the joule, J [(kg m2
)/s2
]. In medical science, energy unit is calories. A calorie is defined as the
amount of energy required to increase the temperature of one gram of water by 1°C.
This gives 1 calorie equals to 4.184 joules.
When expressing the heat capacity as an intensive property, the physical quantity is divided by the
amount of substance, mass, or volume, which makes the quantity independent of the size of the
sample. The molar heat capacity is the heat capacity per unit amount (mole) of a pure substance
and the specific heat capacity (Cp), known as specific heat, is the heat capacity per unit mass of a
matter. The SI unit of specific heat capacity is [kJ/(kg K)]. Cp is generally a temperature dependant
function. However within a good approximation it can be taken as a constant over a moderate
temperature range for a pure single phase materials. The Cp has been measured and tabulated for
the wide range of liquids and solids. These tables are widely available online and in the specialised
book in the University Library [Domalski, E.S. (1984), Furukawa, G.T. (1968), Jiang, Q. W. (2011)].
Calorimetry can be characterised as the analytical technique aims to measure the quantities of
heat released or absorbed by a system during, as example, a chemical reaction.
The amount of heat that flows in or out of a system depends on:
 the quantity of matter in the system,
 the nature of that matter,
 a system temperature change as it absorbs or releases heat.
Calorimetry is performed with a device called calorimeter that provides good heat insulation of
internal volume from its surroundings and a possibility to measure an internal volume
temperature changes in order to determine specific heats of a material located in the internal
volume.
Specific heat determination (calorimetry) approach was developed with account of the
thermodynamic laws.
The Zeroth Law of Thermodynamics states that, “If two samples at different temperatures
(indicated as THot and TCool) are placed in direct physical contact, heat will be lost by the hotter
sample THot and gained by the cooler one TCool. This heat exchange will continue until the moment
of both samples achieved the same final temperature, TFinal”.
On the other hand, the First Law of Thermodynamics states that, “During heat exchange heat is
neither created nor destroyed.” Thus calorimetry investigation employs the fact that heat lost by
one part of the system equals to the heat obtained by other part of the system if the system is
thermo insulated.
Calorimetry experiment can be either performed using constant pressure or a constant volume
method. The constant volume approach is traditionally used for measuring the heat of
combustion and is done with a bomb-type (constant-volume) calorimeter. In the Laboratory, you
will use the constant pressure (isobaric) calorimetry to determine Cp of different samples

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