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 Thermodynamics 3
If you don't know heat or cold in thermodynamics, you must at least have a unit for the amount of heat. This is the enthalpy h, which, however, is related to a specific temperature. Thus, the specific enthalpy is the heat, or
thermodynamic energy, contained in one kilogram of a substance at 0°C (273 K). For water, it is set equal to zero, meaning it is positive above this temperature and negative below it.
In the figure above, the temperature is shown dependent on the specific enthalpy at normal pressure. The temperature rises to 100°C when 417 kJ per kg of heat energy is added to the water. Then comes the
aforementioned pause in the temperature increase, during which the water evaporates. It is the area between the two vertical lines. The water has become 'wet steam'. If energy is added beyond 2,675 kJ/kg, the
temperature rises further. The product is called 'superheated steam'.
Incidentally, while one kilogram of water occupies 0.001 cubic meters of space, as steam only approximately 1.7 cubic meters, or 1,700 times that amount. If you wanted to boil water at room temperature at an absolute
pressure below 1 bar, less than 0.025 bar would be needed. We wouldn't want to subject our combustion engine to this enormous diversion of energy at just 20°C.
So, if you climb a 3,000-meter-high mountain, the air pressure there is only about 0.7 bar (700 hPa). This means water already boils at 90°C. Boiling means that the molecules no longer just 'drum' against the imaginary
wall to form another molecular compound; they break out, take up 1700 times the space like with water vapor too, and more than 30 times as much for R134a. You can guess that this requires an extra portion of energy.
Here we have a phase diagram of water, where 'phases' represent the individual states of matter, in this case water, and their mixed forms. The point at which the three curves converge is called the triple point. We were
already very close to it when we tried to boil water at 20°C with an absolute pressure close to zero.
| Curves as phase boundary lines. |
If we had created an even lower pressure in the experiment, the water (top center panel) would have turned into either steam (bottom right panel) or ice (far left panel) even with very small temperature changes around 0°C.
The dashed line describes the progression of the curve from zero to the triple point for all other substances except water. For example, for refrigerants, although with different numerical values on the axes.
| The point 221 bar/374°C -> 'Critical Point'. |
To the left of the lower left curve, we have ice, and to the right, steam. Along the curve (at this incredibly low pressure), both states can exist. To the left of the curve beyond the triple point, we have ice; to the right, we have
water. Below the curve, between the triple point and the critical point, we have steam again; above, we have water. Along the phase boundary lines, both states are possible. This stops with the critical point.
| Saturated steam = without water content at boiling temperature |
| Superheated steam = without water content higher than boiling temperature |
Beyond the critical point, the molecules have in any case separated from the molecular structure. The substance depicted by the curve then exists only in vapor form.
A good example of the cooling effect of condensation is Formula 1 racing. If the racing car has to stand still for a certain period with the engine running, such as during the start, the mechanics place dry ice in front of the
radiator(s) to increase their cooling effect. This is sufficient during a race, but not when stationary without airflow.
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