The History of Steam
dwells on the fundamental constituents of the universe, the forces
they exert on one another, and the results produced by these
forces. Here, we understand water molecules, what happens when
they are heated, pressurized and turned into steam.
- The Bell curve –
understanding the phase diagram of steam
......Enthalpy of Evaporation
......Producing steam at
......Enthalpy of Saturated Steam (hg)
Steam tables and how to read them
7 Geography of a process
11 Steam Table
understand steam we need to first understand what happens to water
when it evaporates or boils. When we look at water very closely,
we see that it is made of two atoms of hydrogen and one of oxygen.
We consider the molecules are bonded together by electrical
charges (referred to as the hydrogen bond). The degree of
excitation of the molecules determines the physical state (or
phase) of water – solid, liquid or gas, ie, ice, water or
Evaporation in an open vessel
energy of the particles in a liquid is governed by the
temperature. The higher the temperature, the higher the average
energy. Within that average, some particles have higher and some
have lower energy.
Some of the more energetic particles on
the surface of the liquid can be moving fast enough to escape from
the attractive forces holding the liquid together. They evaporate.
Our particles here are the molecules of H2O and we consider the
molecules as bonded together by electrical charges (referred to as
the hydrogen bond). The degree of excitation of the molecules
determines the physical state (or phase) of the substance. Steam
molecules are far more excited than the same molecules in a liquid
phase (water) or solid phase (ice).
The diagram shows a
small region of a liquid near its surface. Molecules in the water
jostle with each other and new molecules gain enough energy to
escape from the surface. Notice that evaporation only takes place
on the surface of the liquid. If you look at water which is just
evaporating in the sun, you don't see any bubbles. Water molecules
are simply breaking away from the surface layer.
evaporation is a surface phenomenon - since the vapor pressure is
low and since the pressure inside the liquid is equal to
atmospheric pressure plus the liquid pressure, bubbles of water
vapor cannot form. But at the boiling point, the saturated vapor
pressure is equal to atmospheric pressure, bubbles form, and the
vaporization becomes a volume phenomena.
of a liquid in a closed container
Now imagine what happens
if the liquid is in a closed container. Common sense tells you
that water in a sealed bottle doesn't seem to evaporate - or at
least, it doesn't disappear over time.
But there is
constant evaporation from the surface. Particles continue to break
away from the surface of the liquid - but this time they are
trapped in the space above the liquid. As the gaseous particles
bounce around, some of them will hit the surface of the liquid
again, and be trapped there. There will rapidly be an equilibrium
set up in which the number of particles leaving the surface is
exactly balanced by the number rejoining it. In this equilibrium,
there will be a fixed number of the gaseous particles in the space
above the liquid. When these particles hit the walls of the
container, they exert a pressure. This pressure is called the
saturated vapour pressure (also known as saturation vapour
pressure) of the liquid.
The effect of temperature on
the equilibrium between liquid and vapour
If you increase
the temperature, you are increasing the average energy of the
particles present. That means that more of them are likely to have
enough energy to escape from the surface of the liquid. That will
tend to increase the saturated vapour pressure.
water is heated at atmospheric pressure, its temperature rises
until it reaches 100°C, the highest temperature at which water
can exist as a liquid at this pressure. At this point, the water
is saturated with energy. It cannot absorb any more heat while
remaining a liquid. Additional heat does not raise the
temperature, but converts the water to steam, ie, a 'phase change'
occurs. We call this the “saturation” point
or, Boiling point and with any further addition of energy,
some of the water will boil off as steam
boiling point is the temperature at which the vapor pressure of
the substance equals the ambient pressure. So, boiling point is
dependent on the pressure. At heights, where the atm pressure is
much lower, the boiling point is also lower. The boiling point
increases with increased ambient pressure up to the critical
point, where the gas and liquid properties become identical. The
boiling point cannot be increased beyond the critical point.
Likewise, the boiling point decreases with decreasing ambient
pressure until the triple point. The boiling point cannot be
reduced below the triple point.
In terms of intermolecular interactions, the
boiling point is the point at which the liquid molecules possess
enough heat energy to overcome the various intermolecular
attractions binding the molecules into the liquid. Therefore the
boiling point is also an indicator of the strength of these
Before heat is applied, there is normal
evaporation from the surface of the water. In the liquid phase,
the molecules are free to move, but are still very close due to
mutual attraction, and collisions occur frequently. We start
heating the water. More heat increases molecular agitation and
collision. As the temperature increases and the water approaches
its boiling condition, some molecules attain enough kinetic energy
to reach velocities that allow them to momentarily escape from the
liquid into the space above the surface, before falling back into
the liquid. Further heating causes greater excitation and the
number of molecules with enough energy to leave the liquid
increases. There is now enough energy to disrupt the attractive
forces between the molecules throughout the liquid.
of vapour (steam) form within the body of the liquid - those are
the bubbles you see when a liquid boils. These steam bubbles float
up (the density of steam is much less than that of water) and
break through the surface.[Note: The bubbles that precede real
boiling in the pot on the stove are either (formerly) dissolved
air or water vapor forming on the very hot bottom of the pot that
will be condensed before it can get to the top of the liquid.] The
space immediately above the water surface is filled with low
density steam molecules. There are now more molecules leaving the
liquid than entering it. The water is freely evaporating. This is
the boiling point.
Once the liquid starts to boil, the
temperature remains constant until all of the liquid has been
converted to a gas. This 'phase change' requires a tremendous
amount of additional energy input. The ability of steam to carry
this large amounts of thermal energy is the property that makes it
so desirable as a working fluid.
how much heat does hot water contain?
1kg of water
at 30°C has 30 Kcals of heat energy trapped in it; at 100°C
it has 100 Kcals of heat energy.
Understanding the Vapour pressure
or, Steam Saturation curve
For a pure substance there is a
definite relationship between saturation pressure and saturation
temperature. If the pressure remains constant, adding more heat
does not cause the temperature to rise any further but causes the
water to form saturated steam. The temperature of the boiling
water and saturated steam within the same system is the same, but
the heat energy per unit mass is much greater in the steam.
atmospheric pressure the saturation temperature is 100°C.
However, if the pressure is increased, this will allow the
addition of more heat and an increase in temperature. The higher
the pressure, higher the heat that can be added to the liquid
without a change of phase and therefore we reach higher saturation
Therefore, increasing the pressure
effectively increases both the enthalpy of water, and the
saturation temperature. The graphical representation of this
relationship between saturation temperature and pressure at
saturated conditions is called the vapor pressure curve or, Steam
saturation curve (given on page 5).
An egg on a mountain
and a pressure cooker
The normal boiling point of water is
100°C. But if you try to cook an egg in boiling water while
camping in the Himalayas at an elevation of 10,000 feet, you will
find that it takes longer for the egg to cook because water boils
at only 90°C at this elevation.
In theory, you
shouldn't be able to heat a liquid to temperatures above its
normal boiling point. However in a pressure cooker we can decrease
the amount of time it takes to cook food. In a typical pressure
cooker, water can remain a liquid at temperatures as high as
120°C, and food cooks in as little as one-third the normal
To explain why water boils at 90°C in the
mountains and 120°C in a pressure cooker, even though the
normal boiling point of water is 100°C, we have to understand
why a liquid boils. By definition, a liquid boils when the vapor
pressure of the gas escaping from the liquid is equal to the
pressure exerted on the liquid by its surroundings, as shown in
the figure below.
water in a kettle at MSL
The normal boiling point of
water at MSL (mean sea level) is 100°C because this is the
temperature at which the saturated vapor pressure of water is 1
atm. Under normal conditions, when the pressure of the atmosphere
is also approximately 1 atm (or 101325 Pa or 101.325 kPa or 760
mmHg) water boils at 100°C.
an egg on the Himalayas
At 10,000 feet above sea
level somewhere on the Himalaya Mountains, the pressure of the
atmosphere is only 0.69 atm, 526 mmHg. At these elevations, water
boils when its vapor pressure is 0.69 atm, which occurs at a
temperature of 90°C. Therefore it takes longer to cook an egg
up in the mountains.
On top of Mount Everest the pressure
is about 0.256 atm, 260 mbar (26 kPa) so the boiling point of
water is 69°C.
water in a pressure cooker
Pressure cookers are
equipped with a valve that lets gas escape when the pressure
inside the pot exceeds some fixed value. This valve is often ste
at 15 psi, which means that the water vapor inside the pot must
reach a pressure of 2 atm before it can escape. Because water
doesn't reach a vapor pressure of 1 atm until the temperature is
120°C, it boils in this container at 120°C. Therefore food
cooks faster in a pressure cooker as it attains a higher
temperature by increasing the pressure.
steam saturation curve is extended for higher pressure towards the
right as below:
blue saturation curve is where the steam/water mix is saturated,
or boils at the different conditions of pressure and temperature.
Eg, 100°C at 0 kg/cm2g or 170°C at 7 kg/cm2g.
Above the blue line, the temperatures is above saturation
temperature and is called the degree of superheat of the steam.
Below the blue curve, we have hot water, not steam, and this is
called sub-saturated water.
is the Triple point ?
In physics, the triple point of a
substance is the temperature and pressure at which three phases
(gas, liquid, and solid) of that substance may coexist in
The triple point of water is
used to define the kelvin, the SI unit of thermodynamic
temperature. The number given for the temperature of the triple
point of water is an exact definition rather than a measured
The single combination of pressure and
temperature at which water, ice, and water vapour can coexist in a
stable equilibrium occurs at exactly 273.16 kelvins (0.01 °C)
and a pressure of 611.73 pascals (ca. 6 millibars, .006037 Atm).
At that point, it is possible to change all of the substance to
ice, water, or steam by making infinitesimally small changes in
pressure and temperature. (Note that the pressure referred to here
is the vapor pressure of the substance, not the total pressure of
the entire system.)
At this temperature of 0°C or 273K
the Enthalpy of water is considered to be 0. Enthalpy of all
states are easily found with reference to this triple point.
Bell curve - understanding
the phase diagram of steam.
The Bell curve
this graph, ABCD is a constant pressure line, say 3.5 barg. The
other lines on top of it are also constant pressure lines, just at
higher pressures. It is called a bell curve because of its
Assume we have 1kg of water held at a
constant pressure. This water will be at atmospheric
temperature, say 25°C at point M between points A and B lying
on the saturated liquid line.
Enthalpy of water (hf)
We start heating the water. The water shows a change in
temperature. Therefore this heat is called Enthalpy of water,
liquid enthalpy or sensible heat (hf) of water. Remember, enthalpy
is the amount of heat available from any medium in its current
state. It moves from point M slowly towards point B. Water starts
boiling at point B. Here is where the change of phase from liquid
to vapour starts.
Enthalpy of evaporation or latent heat
Our sample of water, now is a mixture of two phases,
water and steam (also called sub-saturated water) and moves along
line BC changing into vapour as more heat is applied.
moving from B to C say at point X, the temperature of the
two-phase mixture does not change, as the extra heat applied
to this mixture, is being absorbed by the water which evaporates
into steam. What this means is that the heat being applied is only
being used for a phase change to occur. This heat is the Enthalpy
of evaporation or latent heat (hfg).
Like the phase change
from ice to water, the process of evaporation is also reversible.
The same amount of heat (hfg) that water absorbs to change to
steam is released back to its surroundings during condensation,
when steam meets any surface at a lower temperature. Steam then
changes state from gas to liquid (water) by giving up its Enthalpy
of evaporation (hfg) to the process and is called condensate.
Thus, Enthalpy of evaporation is the most useful part of heat,
as it is the only part of heat which is extracted by the process
when steam condenses back to water.
At point C there is
no water (liquid) left. We are left with 100% saturated steam or,
100% dry steam. Our water sample has converted to 1 kg of steam.
This is the saturated vapour line.
Our aim is to use
steam as close to point C, on the saturated vapour line as steam
is driest at this point.
we continue to provide more heat to the steam we move from point C
to D. Steam is now 100% dry (dryness fraction =1) and it absorbs
the extra heat to show a change in temperature as there is no
phase change possible any more. This heat is also 'sensible' heat
as the temperature of the steam starts to rise. The steam now is
called superheated steam.
Producing Steam at Higher
pressure from a boiler
If the steam is able to flow from
the boiler at the same rate that it is produced, the addition of
further heat simply increases the rate of production. But, if the
steam is restrained from leaving the boiler, and the heat input
rate is maintained, the energy flowing into the boiler will be
greater than the energy flowing out. This excess energy raises the
pressure say, to 10 barg, in turn allowing the saturation
temperature to rise to 184°C, as the temperature of saturated
steam correlates to its pressure. Now we are on the line AB'C'D'.
Bell curve - higher pressure line
at pressures higher than atmospheric, more heat must be added to
water (hf ' ) before it can turn to steam. Increasing the pressure
increases the temperature of the phase change, and increases the
total amount of energy ( hg' ) the water/steam can carry. But, the
Enthalpy of evaporation ( hfg' ), which is the most useful part of
heat, actually reduces!!
In between the
saturated liquid line and saturated vapour line - a mixture of
steam + water exists. This is wet steam. Only water exists to the
left of the saturated liquid line, and only superheated steam
exists in the region to the right of the saturated vapour
The critical point is the
highest temperature at which liquid can exist. This is the point
at which the saturated liquid and saturated vapour lines meet. The
enthalpy of evaporation keeps decreasing as the pressure increases
and it becomes zero at the critical point. Water changes directly
into saturated steam at the critical point. Above the critical
point only gas may exist. In gaseous state the molecules have an
almost unrestricted motion, and the volume increases without limit
as the pressure is reduced. For steam the critical point is at
374.15°C and 221.2 bara.
Enthalpy of saturated
steam, i.e. total heat of saturated steam
The total energy
in saturated steam is simply the sum of the enthalpy of water and
the enthalpy of evaporation.
= hf + hfg
= Total enthalpy or total heat of saturated steam (kJ/kg)
Liquid enthalpy (Sensible heat) (kJ/kg)
hfg = Enthalpy of
evaporation (Latent heat) (kJ/kg)
In tobacco drying, we
start providing heat to dry the moisture laden leaves.
stage: the process uses sensible heat to raise the temperature of
the tobacco. ie, the temperature of process rises.
stage: More heat is supplied, but the temperature of tobacco
doesn't rise. The water now uses this heat to change state to
steam (latent heat) and evaporates from the tobacco.
Tables and how to read them.
of the information just discussed has fortunately been compiled
into steam tables. This is the table for the properties of
Saturated Steam and Saturated Water. It shows us the correlation
of data for various pressures and temperatures.
make sure you are converting between absolute and gauge pressures
There are other tables which give you the
properties of superheated steam, but we are only concerned with
the Saturated steam and water tables, for now.
1 kg of water heated from 0°C to 100°C contains
100 kcals (100°C - 0°C) per kilogram. To go from 100°C
water to 100°C steam at atmospheric pressure requires another
540 kcals per kg. This extra heat which has to be added to convert
1 kg of water to 1 kg of steam is the 'Latent Heat of
Evaporation'. This huge energy that water absorbs while changing
state is the most important part of steam. This is the the part of
heat that steam transfers to process. Remember, while this Latent
heat is being added, the water and the steam released are both at
the same temperature, 100°C.
These drawings show how much heat is required to
generate one kg of steam at 10 kg/cm2g pressure. Note the extra
heat and higher temperature required to make water boil at 10
kg/cm2g pressure than at atmospheric pressure. Note, too, the
lesser amount of heat required to change water to steam at the
Suppose that we have water vapor (or steam) in a closed
container at 120°C and 1 atm. Since the temperature of the
system is above the normal boiling point of water, there is no
reason for the steam to condense to form a liquid. Nothing happens
as we slowly compress the container thereby raising the pressure
on the gas until the pressure reaches 2 atm. At this point, the
system is at the boiling point of water for that pressure, and
some of the gas will condense to form a liquid. As soon as the
pressure on the gas exceeds 2 atm, the vapor pressure of water at
120°C is no longer large enough for the liquid to boil. The
gas therefore condenses to form a liquid, as shown in the fig
Water boils at 100°C at an atmospheric pressure of
0 kg/cm2 g, and 100 kcals of energy are required to heat 1 kg of
water from 0°C to its saturation temperature of 100°C.
Therefore the specific enthalpy of water at 0 kg/cm2 g and 100°C
is 100 kcals/kg, as shown in the steam tables.
kcals/kg of energy are required to evaporate 1 kg of water at
100°C into 1 kg of steam at 100°C. Therefore at 0 kg/cm2 g
the specific enthalpy of steam is 640 kcals/kg, as shown in the
Enthalpy of steam, hg = 100 + 540 = 640
However, steam at atmospheric pressure is of a
limited practical use. This is because it cannot be conveyed under
its own pressure along a steam pipe to the point of use.
1. Why is steam generated at higher pressures in a boiler if
Enthalpy of evaporation (hfg) drops as pressure increases?
of the pressure/volume relationship of steam, (volume is reduced
as pressure is increased) generating steam at a pressure of at
least 7 barg enables the steam distribution pipes to be kept to a
2. Why does ice float on water?
most substances, the density decreases as it changes from the
solid to the liquid phase. However, H2O is an exception to this
rule as its density increases upon melting, which is why ice
floats on water.