13. States Of Matter, Heat capacity and Latent Heat Flashcards
Describe the difference between
solids, liquids and gases.
Substances exist in the solid, liquid or gaseous state.
SOLID
In the solid state a substance has a fixed shape, and the atoms or molecules of which it is composed are arranged in a regular pattern or lattice.
In a molecular substance there are
strong bonds between the atoms that make
up the molecules and
weaker bonds between the molecules themselves.
The
spacing between particles is at a minimum.
Liquid
In a liquid state the substance has no fixed form;
it takes up the shape of its
container,
but does not expand to fill it.
Particles in a liquid constantly ‘jiggle’
about and the spacing between particles is slightly greater than that in the
solid.
Particles are free to move from one place to another in the liquid and
there is no regular structure.
Gas
In a gas the particles are very far apart
and move about in rapid random
motion.
A gas has no shape and completely
fills any container.
A gas which is
below its critical temperature
is called a vapour.
What causes a substance to
change state?
At any time the state in which
a substance exists
depends on its composition
and its temperature.
In general, substances change state
from solid to liquid to gas as the temperature rises.
This is because the
temperature is a measure of the average kinetic energy of the
particles that make up the substance.
It is the increase in
energy possessed by the substance that
causes the change of state.
As a solid is heated,
it absorbs energy,
which causes increased vibrations
of the particles within it.
Eventually, the kinetic energy of individual particles becomes so great as to overcome the inter-molecular forces holding the lattice together.
The lattice breaks up in a process
called melting or fusion.
Further heating results in
more energy being given to the
liquid particles until the
last residual forces holding the particles
in close contact are overcome
and the particles separate widely,
move randomly at high speeds and become
a gas. This is vaporisation.
What is ‘heat capacity’?
The amount of energy given to a
material to cause a 1 °C rise in temperature
is called the heat capacity.
There are two measures of heat capacity:
> The Specific Heat Capacity is defined as:
• The amount of heat required to
raise the temperature of 1 kg of a
substance by 1 kelvin (or 1 °C)
• The units are joule per kilogram per kelvin
(J kg−1 K−1).
> The Molar Heat Capacity is the
amount of heat required to raise
the temperature of
1 mole of the substance by 1 K (or 1 °C)
• The units are joule per mole per kelvin (J mol−1 K−1).
In general the more complex the chemical structure, the higher will be the heat capacity.
This is because there are more
ways for the molecule to vibrate
in a complex molecule.
Each mode of vibration can store kinetic energy.
Therefore it takes more energy to heat it up by 1 degree.
Draw a graph to demonstrate
the concept of latent heat and
explain what happens when a
solid is heated.
The diagram shows the rise in temperature as a solid is heated at a steady rate.
During the phase changes
(fusion and vaporisation)
the temperature remains constant
even though heat is still being applied.
This heat is used internally to
overcome the attraction
of the inter-molecular bonds.
Work needs to be done to separate
the particles against the
forces of attraction.
This heat is known as latent heat
and the term comes from the
Latin verb ‘latere’ – to lie hidden.
Latent heat refers to the amount of
heat energy absorbed (or released)
by a substance as it changes
phase at a given temperature.
Latent heat of fusion
When going from solid to liquid
(or vice versa)
it is called the latent heat of fusion.
When going from liquid to gas (or vice versa) it is called the latent heat of vaporisation.
> The specific latent heat of fusion of a substance is the heat required to convert 1 kg of solid at its melting point into liquid at the same temperature. The unit is joule per kilogram (J kg−1).
Specific heat of vaporisation
> The specific latent heat of vaporisation of a substance is the heat required to convert 1 kg of liquid at its melting point into vapour at the same temperature.
The unit is joule per kilogram (J kg−1).
The specific latent heat of vaporisation
is greater than that of fusion
(seen on the graph)
because more energy is
required to overcome the intermolecular
bonds here to liberate a gas.
A change of phase can be:
From solid to liquid = Fusion (Melting)
From liquid to vapour = Vaporisation
From vapour to liquid = Condensation
From liquid to solid = Freezing
The latent heat is given out to the surroundings when condensation or freezing occurs.
During melting or freezing the solid phase and the liquid phase are in equilibrium.
During boiling the liquid and vapour phase are in equilibrium.
What clinical examples can you
give where the concept of latent
heat is important?
1 Use of ethyl chloride to provide local anaesthesia: vaporisation of the liquid phase of the
agent from the surface
of the skin causes cooling,
rendering the sprayed area numb.
2 Use of volatile anaesthetic agents:
vaporisation of the agent leads to cooling,
which reduces the subsequent
rate of vaporisation of the remaining agent
and hence reduces the SVP.
Vaporisers, therefore, have temperature
compensation mechanisms
(e.g. bimetallic strip that adjusts the splitting ratio)
incorporated into their design to avoid fluctuations in the delivery of the volatile agent.
3 Use of nitrous oxide cylinder
and liquid oxygen contained
with a vacuum insulate evaporator:
vaporisation of these agents
leads to cooling,
which reduces the subsequent
rate of vaporisation of the
remaining agent and
hence reduces the SVP.
4
Loss of latent heat from the patient
through warming and humidification of inspired gases.
5
Loss of heat from the patient through evaporation.
6
A steam burn causes far more tissue damage than a burn from boiling water.
What is the triple point of water?
The states of a substance at different
temperatures and pressures may be
represented in a phase diagram.
A typical phase diagram for a pure molecular
substance is shown below.
The diagram shows the three states of
matter – solid, liquid and gas.
Above the triple point (see later),
at a given pressure the substance
progresses from solid to liquid to gas.
The boundaries between the three
main areas of the phase diagram represent
the conditions of temperature and pressure
that two of the three phases are in equilibrium.
Thus along the solid/liquid
boundary a solid is melting to a liquid.
A small increase in pressure or a
small decrease in temperature will
cause the liquid to
change back to solid and vice versa.
At the boundary itself there is a
dynamic equilibrium where as many molecules
of solid change to a liquid as molecules of liquid change to a solid.
Similarly along the liquid/vapour boundary,
liquid is in equilibrium with vapour.
The solid/vapour boundary in the bottom left represents the equilibrium
between solid and vapour. This is sublimation, where a solid changes to a
vapour without going through the liquid phase.
Solid carbon dioxide (dry ice)
changes directly to the CO2 vapour at normal atmospheric pressure.
The point at which the three
boundaries meet is called the triple point.
This is the point at which
all three phases are in equilibrium.
For water, the triple point is
273.16 K (0.01 °C) at 611.73 Pa pressure (0.006 atm).
The triple point of water is used
as the upper fixed point in the
definition of the kelvin scale of temperature.
How does boiling point change with pressure?
The boiling point of a substance
increases with pressure.
This is shown by the positive slope on the liquid/vapour boundary.
Similarly the melting point of
most substances increases with pressure.
Water is an exception;
its melting point decreases with pressure.
Its phase diagram actually has a negative slope
for the solid/liquid boundary.
What is evaporation?
Particles of a liquid can move
to the vapour state at any temperature.
This is known as evaporation.
In a liquid the particles have kinetic energy
and move about randomly in
proximity to each other.
Not all particles have the
same amount of kinetic energy.
Some will be moving slower
and some will move faster.
The distribution of energies of
these particles is shown below:
Fig. 62.4 Distribution of energy in particles in a liquid
A particle in the body of a liquid a will feel the attractive force of all the surrounding particles and will have no net force exerted upon it.
However a particle near or at the
surface of the liquid will have
particles below it and to its side but none above it.
Consequently, it will experience a net force inwards
towards the body of the liquid; this is the origin of surface tension.
Evaporation occurs when a particle
has sufficient energy to escape surface
forces and move into the surroundings.
At a low temperature only a small
proportion of the particles will have sufficient energy to do this
(see lowtemperature
energy curve above).
At higher temperatures a greater proportion
of particles have the required energy so evaporation is strongly dependent on temperature.
Evaporation occurs only at the surface of a liquid.
The particles that escape exert a
pressure in the surroundings called the vapour pressure.
Explain the concept of saturated vapour pressure.
If a liquid is placed into a closed container,
its molecules will evaporate
and
eventually a dynamic equilibrium will form
between the number of particles
escaping the liquid and the
number rejoining in a given time.
When this state is reached,
the space above the liquid
is saturated with vapour particles.
These vapour particles will bounce
off the surfaces of the container
and of the liquid and in
doing so exert a pressure.
This pressure is called the
Saturated Vapour Pressure of the liquid (SVP).
Define SVP
Formally the SVP of a liquid is defined as:
• The pressure exerted by a vapour when in contact with and in equilibrium with its liquid phase within a closed system at a given temperature.
It is fairly obvious that heating the container,
i.e. raising the ambient
temperature,
will raise the energy of the system
and give more molecules
sufficient energy to escape
the surface of the liquid.
Consequently, there will be a greater number of molecules bouncing against the sides of the container,
resulting in a net increase in force,
which is reflected in an increase in SVP.
For a given substance in a closed system the SVP depends on temperature and nothing else
(i.e. it is not affected by atmospheric pressure).
The maximum
SVP of an open system at
sea level is 1 atmosphere.
How can the SVP of water at altitude cause hypoxia?
If atmospheric pressure at 5500 m
above the sea level
is 50 kPa and saturated water vapour pressure
is 6.3 kPa,
the above question can be answered using
the alveolar gas equation:
PAMBO2 = FAMBO2 × Atmospheric pressure at altitude
= 0.21 × 50 kPa
= 10.5 kPa
PiO2 = FAMBO2 × (Atmospheric pressure at altitude - SVP of water)
= 0.21 × (50 - 6.3)
= 9.17 kPa
PAO2 = PiO2 - PACO2/RQ
= 9.17 - 5/0.8
= 2.92 kPa
What are the effects of SVP on a vaporiser?
As vapour is used up and
removed from the vaporiser
chamber it is replaced
by further vaporisation of the volatile liquid.
The process of vaporisation
requires energy
(latent heat of vaporisation).
This energy requirement
causes the temperature within
the vaporiser to drop.
As SVP is dependent on temperature,
this causes the SVP of the vapour to also drop.
In order to prevent this fluctuating SVP resulting in fluctuations in the delivery of volatile agent,
vaporises have temperature-compensating mechanisms (e.g. bimetallic strip or heated vaporiser chambers).
What is boiling?
A liquid boils when the
SVP equals the surrounding ambient pressure.
This results in pockets,
or bubbles,
of vapour forming in the liquid
and the liquid boils.
These bubbles are less dense
than the surrounding liquid and so rise to
the surface, transferring energy to other liquid molecules they collide with on their way.
What are colligative properties?
Colligative properties are
those properties of a solution
that depend on the number of
dissolved particles in a given
mass of solvent
and not on the identities and properties of those particles
(i.e. they depend on the osmolality).
Freezing point –
1 mole of solute added to 1 kg of water will reduce its
freezing point by 1.86 °C.
This is why roads are gritted with salt during winter.
SVP – solute particles occupy space within the solvent, reducing the surface area available for vapourisation and hence reducing SVP.
Boiling point –
this is the temperature at which liquid and vapour phases are in equilibrium but because solute particles occupy space within the solvent,
the surface area available
for solvent particles to enter the vapour phase is
reduced.
In order to re-establish equilibrium,
the boiling point of the solution is
achieved at a higher temperature.
Osmotic pressure – this is increased.
What is Raoult’s law?
This law states that the
depression of SVP of a solvent is
proportional to the molar concentration
of the solute present.
Why can’t you ‘have a nice cup
of tea up Everest’?
Using the example of water in a saucepan at sea level, the ambient pressure
is 1 atmosphere or 101.3 kPa.
The water needs to be heated to 100 °C
to attain sufficient energy to boil at this pressure.
However, at the summit
of Mount Everest, 8848 m above sea level, the ambient pressure is only 0.30 atmospheres or 30 kPa.
It is easy to understand why up there, water
only needs to be heated to around 80 °C to have sufficient energy to escape
its liquid phase and boil and why you cannot have a ‘nice hot cup of tea’ at
the top of Everest.
In space, water cannot be boiled: it simply evaporates
because the ambient pressure is very close to zero.
