Heat
Kinetic Theory of Matter
• Proposes that heat is a measure of the internal
kinetic energy (microscopic jiggles) of the
molecules and atoms making up the substance.
• The precise relationship between the average
molecular kinetic energy KEave and absolute
temperature T is found to be
KEave = (3/2) kT
Where K, known as Boltzmann constant has the value
k = 1.38 x 10-23 J/K
Heat
• A measure of the total kinetic energy of a
system.
• Internal energy in transit from one body of
matter by virtue of a temperature difference
between them.
• Unit of Measurement: Joule, kcal, Btu
• Kilocalorie
• The amount of heat required to raise the
temperature of 1 kg of water by 1°C.
• Relationship between joule and kilocalorie
1 J = 2.39 x 10-4 kcal
1 kcal = 4185 J
Quantity of Heat
Heat lost or gained = mass x c x Temperature change
Q = m c ∆T
Where: Q is the heat lost or gained in kcal
m is the mass in kg,
c is the specific heat kcal/kg⁰C
∆T is the temperature change in ⁰C
• If you are given c in cal/g ⁰C; this works when heat is in calories,
mass in grams, temperature in ⁰C.
Example
• How much heat must be added to 5 kg of water
to increase its temperature by 20 ⁰C? The
specific heat of water is 1cal/g ⁰C.
• Given: m = 5kg or 5000g
c = 1 cal/g ⁰C
∆T = 20 ⁰C
Solution:
Q = m c ∆T
= 5000g x 1 kcal/g ⁰C x 20⁰C
= 100000kcal
Specific Heat
• The movement of molecules depends on the
type of material used.
• Property of a material that explains why
some foods and liquids remain hotter longer
than others.
Specific Heat and Water
• Water has one of highest specific heats at
1cal/g ⁰C.
• One gram of water is able to absorb and
hold more heat than other liquids.
• When water is cold, it can absorb more heat
as it warms up.
11.2.5 Thermal Equilibrium
• Occurs when one or more substances reach a common
temperature by exchanging heat.
Heat lost = Heat gained
Qlost =Qgained
For example, if you put ice in warm water, the warm
water gets cold, the ice gets warm, and melts. When
the ice completely melts, they will all have the same
temperature. They are in equilibrium.
• A 0.5-kg aluminum block at 75⁰C is dropped into a cup
of water with a mass of 2.0 kg at 40⁰C and comes to
thermal equilibrium. What is the final temperature (Tf)
of the water and aluminum? The specific heat of
aluminum is 0.21 kcal/kg⁰C.
Given:
Aluminum Block
Water
mAl = 0.5kg
mwater = 2.0kg
cAl = 0.21 kcal/kg⁰C
cwater = 1kcal/kg⁰C
TiAl = 75⁰C
Tiwater = 40⁰C
• The initial temperature of the aluminum block is greater
than the initial temperature of water. Since heat flow is form
hot to cold, then the aluminum block will lose heat, while
water will gain heat.
• Aluminum block will lose heat;
∆TAl = (Ti–Tf)
= 75⁰C – Tf
• Water will gain heat;
∆Twater = (Tf - Ti)
= T - 40⁰C
• Apply the equilibrium condition:
Qlost =Qgained
mAl x cAl x ∆TAl = mwater x cwater x
∆Twater
(0.5kg)(0.21 kcal/kg⁰C)(75⁰C – Tf) = (2kg)(1
kcal/kg⁰C)(Tf - 40⁰C)
11.2.6 Transmission of Heat
• Conduction
• Convection
• Radiation
11.2.6.1 Conduction
• Conduction involves the transfer of heat
through direct contact
• Heat conductors conduct heat well,
insulators do not
Example of Conduction
11.2.6.2 Convection
• Takes place in liquids and gases as
molecules move in currents
• Heat rises and cold settles to the bottom
Example of Convection
11.2.6.3 Radiation
• Heat is transferred through space
• Energy from the sun being transferred to the
Earth
Example of Radiation
11.2.7 Change in State
• When the atoms and molecules are no longer
closely associated with one another, all bonds
have been broken and the liquid has become a
gas.
11.2.7.1 Melting and Heat of Fusion
In melting, it takes an amount of heat
Qf = m Hf
to change from a solid to a liquid
Where:
Qf = amount of heat to melt the solid in kcal
Hf = heat of fusion in kcal/kg
m = mass of the solid in kg
11.2.7.2 Evaporization and Heat of
Vaporization
In evaporating, it takes an amount of heat
Qv = m Hv
to change from a liquid to a gas
Where :
Qv = amount of heat to evaporate the liquid in kcal
Hv = heat of vaporization in kcal/kg
m = mass of the liquid in kg
11.2.7.3 Sublimation
• More rarely, the solid can change directly to
a gas.
11.2.7.4 Equation for Change of State
The amount of heat needed to turn mass of ice at
some initial temperature into a gas at a final
temperature can be computed by the following :
Isuyat ko ra  
11.3 Thermodynamics
•
•
What we’ve learned thus far about heat and
thermal energy is summed up in the laws of
thermodynamics. The word thermodynamics
stems from Greek for “movement of heat.”
Is a branch of physics looking at how changes
in energy. Work and the flow of heat influence
each other.
11.3.1.1 The First Law
• Whenever heat flows into or out of a system, the
gain or loss of thermal energy equals the amount of
heat transferred.
Mathematically speaking,
Total Heat = change in internal energy + work done
11.3.1.2 The Second Law
It is impossible for heat to flow from a point
of lower temperature to a point of higher
temperature without the application of
energy from an external source.
11.3.1.3 Entropy
• Is a measure of the amount of wasted heat in the
system.
• It is measure of the amount of disorder in the
system.
Entropy change = heat added / Temperature (in K )