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The properties of steam explained here, including the ability of steam under pressure to carry, and then give up, large amounts of energy. Topics include saturated steam tables, dryness fraction and flash steam.
A better understanding of the properties of steam may be achieved by understanding the general molecular and atomic structure of matter, and applying this knowledge to ice, water and steam.
A molecule is the smallest amount of any element or compound substance still possessing all the chemical properties of that substance which can exist. Molecules themselves are made up of even smaller particles called atoms, which define the basic elements such as hydrogen and oxygen
The specific combinations of these atomic elements provide compound substances. One such compound is represented by the chemical formula 2O, having molecules made up of two atoms of hydrogen and one atom of oxygen.
The reason water is so plentiful on the earth is because hydrogen and oxygen are amongst the most abundant elements in the universe. Carbon is another element of significant abundance, and is a key component in all organic matter.
Most mineral substances can exist in the three physical states (solid, liquid and vapour) which are referred to as phases. In the case of H2O, the terms ice, water and steam are used to denote the three phases respectively.
The molecular arrangement of ice, water, and steam is still not fully understood, but it is convenient to 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.
All the three phases of a particular substance can only coexist in equilibrium at a certain temperature and pressure, and this is known as its triple point.
The triple point of H2O, where the three phases of ice, water and steam are in equilibrium, occurs at a temperature of 273.16 K and an absolute pressure of 0.006 112 bar. This pressure is very close to a perfect vacuum. If the pressure is reduced further at this temperature, the ice, instead of melting, sublimates directly into steam.
In ice, the molecules are locked together in an orderly lattice type structure and can only vibrate.
In the solid phase, the movement of molecules in the lattice is a vibration about a mean bonded position where the molecules are less than one molecular diameter apart.
The continued addition of heat causes the vibration to increase to such an extent that some molecules will eventually break away from their neighbours, and the solid starts to melt to a liquid state. At atmospheric pressure, melting occurs at 0 °C. Changes in pressure have very little effect on the melting temperature, and for most practical purposes, 0 °C can be taken as the melting point. However, it has been shown that the melting point of ice falls by 0.0072 °C for each additional atmosphere of pressure. For example, a pressure of 13.9 bar g would be needed to reduce the melting temperature by 0.1 °C.
Heat that breaks the lattice bonds to produce the phase change while not increasing the temperature of the ice, is referred to as enthalpy of melting or heat of fusion. This phase change phenomenon is reversible when freezing occurs with the same amount of heat being released back to the surroundings.
For 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.
In the liquid phase, the molecules are free to move, but are still less than one molecular diameter apart due to mutual attraction, and collisions occur frequently. More heat increases molecular agitation and collision, raising the temperature of the liquid up to its boiling temperature.
Enthalpy of water, liquid enthalpy or sensible heat (hf) of water
This is the heat energy required to raise the temperature of water from a datum point of 0 °C to its current temperature.
At this reference state of 0 °C, the enthalpy of water has been arbitrarily set to zero. The enthalpy of all other states can then be identified, relative to this easily accessible reference state.
Sensible heat was the term once used, because the heat added to the water produced a change in temperature. However, the accepted terms these days are liquid enthalpy or enthalpy of water.
At atmospheric pressure (0 bar g), water boils at 100 °C, and 419 kJ of energy are required to heat 1 kg of water from 0 °C to its boiling temperature of 100 °C. It is from these figures that the value for the specific heat capacity of water (Cp) of 4.19 kJ/kg °C is derived for most calculations between 0 °C and 100 °C.
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. As the water is heated to its boiling point, bubbles of steam form within it and rise to break through the surface.
Considering the molecular arrangement of liquids and vapours, it is logical that the density of steam is much less than that of water, because the steam molecules are further apart from one another. The space immediately above the water surface thus becomes filled with less dense steam molecules.
When the number of molecules leaving the liquid surface is more than those re-entering, the water freely evaporates. At this point it has reached boiling point or its saturation temperature, as it is saturated with heat energy.
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.
At 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 without a change of phase.
Therefore, increasing the pressure ef fect ively increases both the enthalpy of water, and the saturation temperature. The relationship between the saturation temperature and the pressure is known as the steam saturation curve (see Figure 2.2.1).
Water and steam can coexist at any pressure on this curve, both being at the saturation temperature. Steam at a condition above the saturation curve is known as superheated steam:
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. 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, in turn allowing the saturation temperature to rise, as the temperature of saturated steam correlates to its pressure.
Enthalpy of evaporation or latent heat (hfg)
This is the amount of heat required to change the state of water at its boiling temperature, into steam. It involves no change in the temperature of the steam / water mixture, and all the energy is used to change the state from liquid (water) to vapour (saturated steam).
The old term latent heat is based on the fact that although heat was added, there was no change in temperature. However, the accepted term is now enthalpy of evaporation.
Like the phase change from ice to water, the process of evaporation is also reversible. The same amount of heat that produced the steam is released back to its surroundings during condensation, when steam meets any surface at a lower temperature.
This may be considered as the useful portion of heat in the steam for heating purposes, as it is that portion of the total heat in the steam that is extracted when the steam condenses back to water.
Enthalpy of saturated steam, or total heat of saturated steam
This is the total energy in saturated steam, and is simply the sum of the enthalpy of water and the enthalpy of evaporation.
The enthalpy (and other properties) of saturated steam can easily be referenced using the tabulated results of previous experiments, known as steam tables.
The steam tables list the properties of steam at varying pressures. They are the results of actual tests carried out on steam. Table 2.2.1 shows the properties of dry saturated steam at atmospheric pressure - 0 bar g.
At atmospheric pressure (0 bar g), water boils at 100 °C, and 419 kJ 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 bar g and 100 °C is 419 kJ/kg, as shown in the steam tables (see Table 2.2.2).
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.
Note: Because of the pressure/volume relationship of steam, (volume is reduced as pressure is increased) it is usually generated in the boiler at a pressure of at least 7 bar g. The generation of steam at higher pressures enables the steam distribution pipes to be kept to a reasonable size.
As the steam pressure increases, the density of the steam will also increase. As the specific volume is inversely related to the density, the specific volume will decrease with increasing pressure.
Figure 2.2.2 shows the relationship of specific volume to pressure. This highlights that the greatest change in specific volume occurs at lower pressures, whereas at the higher end of the pressure scale there is much less change in specific volume.
The extract from the steam tables shown in Table 2.2.2 shows specific volume, and other data related to saturated steam.
At 7 bar g, the saturation temperature of water is 170 °C.
More heat energy is required to raise its temperature to saturation point at 7 bar g than would be needed if the water were at atmospheric pressure. The table gives a value of 721 kJ to raise 1 kg of water from 0 °C to its saturation temperature of 170 °C.
The heat energy (enthalpy of evaporation) needed by the water at 7 bar g to change it into steam is actually less than the heat energy required at atmospheric pressure. This is because the specific enthalpy of evaporation decreases as the steam pressure increases.
However, as the specific volume also decreases with increasing pressure, the amount of heat energy transferred in the same volume actually increases with steam pressure.
Steam with a temperature equal to the boiling point at that pressure is known as dry saturated steam. However, to produce 100% dry steam in an industrial boiler designed to produce saturated steam is rarely possible, and the steam will usually contain droplets of water.
In practice, because of turbulence and splashing, as bubbles of steam break through the water surface, the steam space contains a mixture of water droplets and steam.
Steam produced in any shell-type boiler (see Block 3), where the heat is supplied only to the water and where the steam remains in contact with the water surface, may typically contain around 5% water by mass.
If the water content of the steam is 5% by mass, then the steam is said to be 95% dry and has a dryness fraction of 0.95.
The actual enthalpy of evaporation of wet steam is the product of the dryness fraction (c) and the specific enthalpy (hfg) from the steam tables. Wet steam will have lower usable heat energy than dry saturated steam.
Because the specific volume of water is several orders of magnitude lower than that of steam, the droplets of water in wet steam will occupy negligible space. Therefore the specific volume of wet steam will be less than dry steam:
Where vg is the specific volume of dry saturated steam.
Steam at a pressure of 6 bar g having a dryness fraction of 0.94 will only contain 94% of the enthalpy of evaporation of dry saturated steam at 6 bar g. The following calculations use figures from steam tables:
The steam phase diagram
The data provided in the steam tables can also be expressed in a graphical form. Figure 2.2.3 illustrates the relationship between the enthalpy and temperature of the various states of water and steam; this is known as a phase diagram.
As water is heated from 0 °C to its saturation temperature, its condition follows the saturated water line until it has received all of its liquid enthalpy, hf, (A - B).
If further heat continues to be added, the water changes phase to a water/vapour mixture and continues to increase in enthalpy while remaining at saturation temperature, hfg, (B - C).
As the water/vapour mixture increases in dryness, its condition moves from the saturated liquid line to the saturated vapour line. Therefore at a point exactly halfway between these two states, the dryness fraction (c) is 0.5. Similarly, on the saturated steam line, the steam is 100% dry.
Once it has received all of its enthalpy of evaporation, it reaches the saturated steam line. If it continues to be heated after this point the pressure remains constant but the temperature of the steam will begin to rise as superheat is imparted (C - D).
The saturated water and saturated steam lines enclose a region in which a water/vapour mixture exists - wet steam. In the region to the left of the saturated water line only water exists, and in the region to the right of the saturated steam line only superheated steam exists.
The point at which the saturated water and saturated steam lines meet is known as the critical point. As the pressure increases towards the critical point the enthalpy of evaporation decreases, until it becomes zero at the critical point. This suggests that water changes directly into saturated steam at the critical point.
Above the critical point the steam may be considered as a gas. The gaseous state is the mostdiffuse state in which the molecules have an almost unrestricted motion, and the volume increases without limit as the pressure is reduced.
The critical point is the highest temperature at which water can exist. Any compression at constant temperature above the critical point will not produce a phase change.
Compression at constant temperature below the critical point however, will result in liquefaction of the vapour as it passes from the superheated region into the wet steam region.
The critical point occurs at 374.15 °C and 221.2 bar a for steam. Above this pressure the steamis termed supercritical and no well-defined boiling point applies.
The term ‘flash steam’ is traditionally used to describe steam issuing from condensate receiver vents and open-ended condensate discharge lines from steam traps. How can steam be formed from water without adding heat?
Flash steam occurs whenever water at high pressure (and a temperature higher than the saturation temperature of the low-pressure liquid) is allowed to drop to a lower pressure. Conversely, if the temperature of the high-pressure water is lower than the saturation temperature at the lower pressure, flash steam cannot be formed. In the case of condensate passing through a steam trap, it is usually the case that the upstream temperature is high enough to form flash steam.
See Figure 2.2.4.
Consider a kilogram of condensate at 5 bar g and a saturation temperature of 159 °C passing through a steam trap to a lower pressure of 0 bar g. The amount of energy in one kilogram of condensate at saturation temperature at 5 bar g is 671 kJ. In accordance with the first law of thermodynamics, the amount of energy contained in the fluid on the low-pressure side of the steam trap must equal that on the high-pressure side, and constitutes the principle of conservation of energy.
Consequently, the heat contained in one kilogram of low-pressure fluid is also 671 kJ. However, water at 0 bar g is only able to contain 419 kJ of heat, subsequently there appears to be an imbalance of heat on the low-pressure side of 671 – 419 = 252 kJ, which, in terms of the water, could be considered as excess heat.
This excess heat boils some of the condensate into what is known as flash steam and the boiling process is called flashing. Therefore, the one kilogram of condensate which existed as one kilogram of liquid water on the high pressure side of the steam trap now partly exists as both water and steam on the low-pressure side.
The amount of flash steam produced at the final pressure (P2) can be determined using Equation 2.2.5:
The case where the high pressure condensate temperature is higher than the low pressure saturation temperature.
Consider a quantity of water at a pressure of 5 bar g, containing 671 kJ/kg of heat energy at its saturation temperature of 159 °C. If the pressure was then reduced down to atmospheric pressure (0 bar g), the water could only exist at 100 °C and contain 419 kJ/kg of heat energy.
This difference of 671 - 419 = 252 kJ/kg of heat energy, would then produce flash steam at atmospheric pressure.
The proportion of flash steam produced can be thought of as the ratio of the excess energy to the enthalpy of evaporation at the final pressure.
The case where the high pressure condensate temperature is lower than the low pressure saturation temperature.
Consider the same conditions as in Example 2.2.3, with the exception that the high-pressure condensate temperature is at 90 °C, that is, sub-cooled below the atmospheric saturation temperature of 100 °C.
Note: It is not usually practical for such a large drop in condensate temperature from its saturation temperature (in this case 159 °C to 90 °C); it is simply being used to illustrate the point about flash steam not being produced under such circumstances.
In this case, the sub-saturated water table will show that the liquid enthalpy of one kilogram of condensate at 5 bar g and 90 °C is 377 kJ. As this enthalpy is less than the enthalpy of one kilogram of saturated water at atmospheric pressure (419 kJ), there is no excess heat available to produce flash steam. The condensate simply passes through the trap and remains in a liquid state at the same temperature but lower pressure, atmospheric pressure in this case.
See Figure 2.2.5.
The vapour pressure of water at 90 °C is 0.7 bar absolute. Should the lower condensate pressure have been less than this, flash steam would have been produced.
The principles of the conservation of energy and mass allow the flash steam phenomenon to be thought of from a different direction.
Consider the conditions in Example 2.2.3.
1 kg of condensate at 5 bar g and 159 °C produces 0.112 kg of flash steam at atmospheric pressure. This can be illustrated schematically in Figure 2.2.6. The total mass of flash and condensate remains at 1 kg.
The principle of energy conservation states that the total energy in the lower-pressure state must equal the total energy in the higher-pressure state. Therefore, the amount of heat in the flash steam and condensate must equal that in the initial condensate of 671 kJ.
Steam tables give the following information:
Therefore, according to the steam tables, the enthalpy expected in the lower-pressure state is the same as that in the higher-pressure state, thus proving the principle of conservation of energy.
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