Heat and temperature are different concepts. Heat is a form of energy measured in Joules, while temperature is a measure of the average kinetic energy of particles measured in Kelvin or Celsius. Objects can contain various forms of energy including kinetic energy from motion and potential energy from forces between particles. Thermal equilibrium occurs when objects in contact reach the same temperature after heat transfer. Specific heat capacity is the amount of energy required to change an object's temperature and depends on the material. Phase changes from solid to liquid or liquid to gas require latent heat and occur when particles gain enough energy to overcome attractive forces.

1. HEAT AND TEMPERATURE
Thermodynamics

2. 1.0 - Introduction
 Heat is a form of energy.
 It is measured like other forms
of energy in J (Joules).
 The above statement should
have led you to realise that heat
and temperature are two
different things as temperature
is measured in K (Kelvin) or
more commonly in °C (Degree
Celsius or Centigrade) or °F
(Degree Fahrenheit).
 An object may contain various
different types of energy.

3. 1.1 - Introduction Cont.
 Take for example a student
throwing a jar containing bees.
The object has the potential
energy caused by the
gravitational field of the earth
and kinetic energy as it is
moving. However, the individual
bees also have their own kinetic
energy - we will call this the
random kinetic energy of the jar.
 If we said that the jar is now a
metal ball and the bees in the
jar are the atoms of the metal
ball, the average random kinetic
energy caused by the vibration
of the atoms is the temperature
of the metal ball.

4. 1.2 - Introduction Cont.
 In a substance, Kinetic Energy is
present due to the masses and
velocities of its particles being
vibrated, rotated or translated, and
Potential Energy is present due to
the attractive forces between each
of the particles as bonds AND
between separate particles.
 The sum of the kinetic and potential
energies of all the particles is called
the internal or thermal energy of
the substance.
 The term heat is used to describe
the internal energy of a substance.
 The study of the transfers of this
energy is called thermodynamics.

5. 2.0 - Thermal Equilibrium
 Thermal Equilibrium is the state at which
two objects in an isolated environment gain
the same temperature after the process of
heat transfer from the body containing
more heat (TB) to the other (TA).
 Note that in an isolated environment, there
is no heat lost to the surroundings –
therefore, the heat lost by the hotter object
(TB) is equal to the heat gained by the less
hotter object (TA).
 Note that this is theoretical and in practise,
heat is lost through radiation (even if the
experiment is conducted in space) – all
objects that have temperatures above
absolute zero, radiate energy in the form of
electro-magnetic radiation. And when
conducted on earth heat is also lost
through conduction and convection (refer
to Slide 5.0 – Calorimetry).

6. 3.0 - Measuring Temperature
 There are many different
scales used to measure
temperature. Below are the
three scales that are mainly
used in the present day;
 The Fahrenheit Scale
 Developed by German
physicist, Gabriel Fahrenheit
(1686 – 1736)
 In this scale, the freezing
point of a salt solution is 0°F,
the freezing point of pure
water is 32°F, and the boiling
point of pure water is 212°F.
 This scale is mainly used in
the US, UK and Canada.

7. 3.1 - Measuring Temperature Cont.
 The Celsius Scale
 Developed by Andres Celsius
(1701 – 1744)
 In this scale, the freezing
point of pure water is 0°C and
the boiling point of pure
water is 100°C.
 The Kelvin/Absolute Scale
 Developed by Lord Kelvin
(1824 – 1907)
 In this scale, 0 K is the
absolute zero temperature –
this means that at this
temperature, there is
absolutely no particle motion.
 Note; (0 K = -273.15°C).

8. 4.0 - Specific Heat Capacity
 In a room, that has a constant temperature
(say 23°C), all the objects have the same
temperature (23°C – thermal equilibrium).
However, if we humans touched a metallic
object in the room it would feel much more
cold than a non-metallic object in the same
room. This is due to the fact that metals are
good conductors of heat. The heat from our
bodies is conducted faster to the metals
than to the non-metals. And because our
body senses the rate at which heat is
transferred to or away from our body, the
metals feel more cold (remember that the
metals still have the same temperature).
 “Good conductors of heat” refers to
substances with a low heat capacity – i.e.
they require relatively less amounts of
energy to raise its temperature (refer to
next slide for a more precise and detailed
description of specific heat capacity).

9. 4.1 - Specific Heat Capacity Cont.
 Specific Heat Capacity is the measure of
how much energy is required to raise the
temperature of 1 kg of a substance by 1
K or 1°C (note – a change of 1 K is exactly
the same as a change of 1°C).
 Different substances have different
specific heat capacities. On the left side
is a table containing the specific heat
capacities of some commonly seen
substances
 On the bottom left is the formula for
specific heat capacity. In this formula, Q
is the heat energy required (J), m is the
mass (kg), c is the specific heat capacity,
ΔT is the change in temperature
(measured in either °C or K – refer to the
first dot point in this slide).

10. 5.0 - Calorimetry
 When two substances are
placed together in a closed
system, thermal equilibrium
occurs.
 In practice, there is always some
heat lost to the surroundings.
There are two main ways in
which such heat loss could be
minimised;
 Carrying out the experiment
quickly.
 Use calorimeters, which have
good insulation to limit the loss
of heat to the surroundings. This
process is called Calorimetry.
Return to Slide 2.0 – Thermal Equilibrium

11. 6.0 - Change of State
 The amount of energy required to
melt 1 kg of an object is called the
specific latent heat of fusion.
 The amount of energy required to
vaporise 1 kg of an object is called the
specific latent heat of vaporisation.
 On the left is the formula for the
energy required to change the state of
a substance. In this formula; Q is the
heat energy required (J), m is the mass
(kg), and L (specific latent heat) of the
object becomes;
 Lf for the specific latent heat of fusion (or)
 Lv for the specific latent heat of
vaporisation.

12. 7.0 - Changing the melting and boiling
Points
 Most substances have fixed
melting and boiling points as
long as they are in pure form.
 To change the melting and
boiling points of various
substances, there are two
main methods which could
be used;
 Adding Impurities to the
substance (and/or)
 Changing the pressure of the
substance and/or its
environment.

13. 8.0 - Evaporation
 Liquids turn into gas without boiling.
 <Image of clothes This process is called evaporation and
occurs all the time.
drying>  For a substance to change state,
energy is required. But note that not
all the individual particles of a
substance have exactly the same
energy (also note – temperature is the
measure of the average random
kinetic energy of the particles of a
substance).
 This is the reason for evaporation –
individual particles with relatively
higher energy are able to reach the
surface of the substance and escape
(e.g. when you leave a bowl of water
at room temperature, it will
eventually evaporate to nothing).

14. 9.0 - Laws of Thermodynamics
 There are two laws of
thermodynamics;
 The first law of
thermodynamics states that
the total increase in the
thermal energy of an isolated
system is equal to the sum of
the heat added to it and the
work done on it. Note that
this is just an extension of
the law of conservation of
energy.
 The second law of
thermodynamics relates heat
transfer to differences in
temperature.

15. 9.1 - Laws of Thermodynamics Cont.
 The laws of thermodynamics also
 <Image of Sand Castle helped develop a new term in
– Entropy> physics, called entropy.
 Entropy is the measure of the
disorder of a system – the more
disorder, the more entropy. It states
that in nature, all ordered systems
head towards becoming disordered.
 An example of this (from
thermodynamics) would be when two
objects, say water and ice, are placed
in contact with each other and allowed
to reach thermal equilibrium. After
equilibrium is reached, the ordered
molecules of the ice become less
ordered; and therefore it now has a
lower entropy.

16. FINALLY! WE’RE DONE!

17. REMEMBER!
This presentation is only designed to help you learn
easier not thorough. So, refer to you textbook for
detailed information on this chapter! And practice
the questions in your textbook if any!