What happens when you heat steam

With the help of these diagrams numerous properties of water and its aggregate states - as well as other substances - and thus fundamental facts of the theory of heat and the structure of matter can be discussed. The process shown consists in gradually adding energy to one kilogram of ice at a temperature of -40 C until it has finally turned into 150 C hot water vapor. During this process, the pressure is kept constant at 1.01325 bar (the "normal pressure", which corresponds to the mean air pressure at sea level).
  • Required energies:
    • We start with 1 kg of ice at a temperature of -40 C. On the horizontal axis of the diagram is the temperature T Plotted in C, the energy on the vertical axis E., which we will feed into the system in small portions from now on. The zero point of the E.-Axis has no absolute meaning, it simply corresponds to the beginning of the process we are considering.
    • At the beginning, our kilogram of ice is in the state that is represented by the intersection of the axes: T = –40 C and (since no energy has yet been supplied) E. = 0.
    • If we now supply the system with energy, it heats up. This is shown by the slightly rising red line. The rise (slope) of this line corresponds to the so-called specific heat capacity. In general, it indicates the energy that has to be added to one kilogram of a substance in order to heat it by 1 C (which is the same as 1 K). The specific heat capacity of ice depends a bit on the temperature. In the temperature range shown, it is around 2000 J / (kg · K) or 2 kJ / (kg · K). In order to heat our kilogram of ice from the initial temperature of -40 C to 0 C, an energy of around 40 · 2 kJ = 80 kJ is required.
    • How much energy is 1 kJ? Which orders of magnitude of energies appear in the diagram? The amounts of energy at issue here are therefore considerable, even though they only concern one kilogram of water.
    • What happens when ice is heated? in the solid state of aggregation the molecules are rigidly bound to one another by electromagnetic forces. They are quite tightly packed next to each other and can neither move freely nor rotate relative to each other. But they can swing against each other. The temperature is a measure of the vibration energy that comes on average to a single molecule. The energy supplied to a solid mainly causes these vibrations to become stronger - it heats up. (More about this "getting stronger" below, in the diagram about the volumes).
    • If the ice has reached a temperature of 0 C, so melts it, i.e. it goes into the liquid physical state above. The temperature at which this happens is characteristic of each substance (and depends on the pressure). she will Melting temperature called. The diagram shows two important properties of the melting process:
      1. Melting takes place (at normal pressure) at a constant temperature of 0 C. That means that a further supply of energy initially Not leads to further warming!
      2. In order to completely melt our kilograms of ice, a considerable amount of energy is required, the so-called Heat of fusion of 333 kJ, i.e. about four times as much as was necessary to warm the ice from -40 C to 0 C.
      Melting is a gradual process in which solid ice and liquid water coexist. The more energy is supplied, the greater the proportion of liquid water, until finally there is no more ice. At the end of this process, the beginning (i.e. the left end) of the blue line is reached.
    • The energy added during melting breaks the rigid bonds between the molecules. The molecules still exert forces on one another, but can now move against one another. Therefore, a liquid - in contrast to a solid - adapts to the shape of a vessel in which it is located.
    • Melting has an inverse: that Freeze. It occurs when energy is withdrawn from a liquid, i.e. when our diagram goes from top to bottom. The rigid connections that are characteristic of solids are formed first from small "crystallization nuclei", which then "grow". This can lead to the phenomenon that the first of these crystallization nuclei form with a delay, i.e. that the Freezing temperature is below the melting temperature. In practice, however, the smallest impurities mean that this does not occur when substances freeze in everyday life, i.e. that the freezing temperature is the same as the melting temperature (0 C for water). The energy that has to be extracted from the water in order to freeze it completely is equal to the heat of fusion.
    • Let's continue adding energy to our kilogram of water! The blue line in the diagram shows that the temperature then rises again, as was the case with ice below the melting temperature. The fact that the blue line rises more steeply than the red indicates that the specific heat capacity of liquid water (i.e. the energy that has to be added to one kilogram of liquid water to heat it by 1 C) is higher than that of ice. It depends a bit on the temperature and is (at normal pressure) about 4200 J / (kg · K) or 4.2 kJ / (kg · K). More precisely, it varies between 4.18 kJ / (kg · K) and 4.22 kJ / (kg · K) depending on the temperature. In order to heat our kilogram of water from 0 C to 100 C, an energy of around 100 4.2 kJ = 420 kJ is necessary. Conversely, it can be said that one kilogram of water emits an energy of 420 kJ when it cools down from 100 C to the freezing temperature at 0 C.
    • The specific heat capacity of liquid water is greater than that of most other substances: a high energy input only leads to a low level of warming. That means that water is a very good one Heat storage is. Since the oceans are mainly made up of water, this is also important for the earth's climate.
    • What happens when liquid water is heated? in the liquid physical state the molecules exert electromagnetic forces on one another, but, as already mentioned - in contrast to the solid state of aggregation - they are not rigidly bound to one another. As in the solid body, they are located relatively close to one another - their distance is in the same order of magnitude as their size. Therefore, they can only move freely over very short distances and keep bumping into each other. These movements become faster with continued energy supply, which manifests itself as warming, i.e. as a rise in temperature.
    • If our kilogram of water has reached a temperature of 100 C, so evaporates it, i.e. it goes into the gaseous physical state above. The temperature at which this happens is characteristic of every substance (and depends very much on the pressure). she will Boiling temperature (or Evaporation temperature) called. The diagram shows two important properties of the evaporation process:
      1. Evaporation takes place (at normal pressure) at a constant temperature of 100 C. That means that a further supply of energy initially Not leads to further warming!
      2. In order to completely evaporate our kilograms of water (at normal pressure), a very large amount of energy is required, the so-called Heat of evaporation of 2257 kJ. It is almost seven times the size of the heat of fusion!
      Evaporation is a gradual process in which liquid water and gaseous water vapor coexist. The more energy is supplied, the greater the proportion of water vapor until finally there is no more liquid water. At the end of this process, the beginning (i.e. the left end) of the green line is reached.
    • The energy supplied during evaporation almost completely breaks the bonds between the molecules. In the gaseous state there are still forces between the molecules, but they are considerably smaller than the forces that act in the liquid state. In gases, the molecules can move freely over relatively long distances before they collide with their own kind. This is why a gas - in contrast to solids and liquids - fills the space available to it (provided it is not - like our atmosphere - kept close to the ground by gravity).
    • Evaporation has an inverse: that Condense. It occurs when energy is withdrawn from a gas, i.e. when our diagram goes from top to bottom. The Condensation temperature is equal to the boiling temperature. The energy that must be extracted from the water vapor in order to condense it completely (the Heat of condensation) is equal to the heat of vaporization.
    • If we continue to add energy to the water vapor, the temperature rises. The rise of the green line in the diagram corresponds to specific heat capacity of water vapor. It depends a little on the temperature. In the temperature range shown, it is around 2 kJ / (kg · K), so it is roughly the same as that of ice.
  • Volumes:
    • If the specified values ​​for the volume that our kilogram of water assumes in the various states of aggregation (at normal pressure), one thing is immediately noticeable: one kilogram of ice or liquid water has a volume of approximately 1 dm3, while one kilogram of water vapor at 100 C fills an area 1700 times as large! This means that every molecule in 100 C hot water vapor has a volume available that is 1700 times larger than before. The molecules are therefore about 12 times (third root from 1700!) As far apart as in liquid water or ice. You now have a considerably larger one free path - this property characterizes a very generally gas. The numbers for the volumes on the way from ice to water vapor very nicely illustrate the dissolution of the relatively tightly packed molecular lattice during evaporation. As the high value of the evaporation heat shows, this costs a lot of energy.
    • Let's go through the displayed volume values, starting with the ice at -40 C, one after the other: In the solid state, the volume increases slightly with increasing temperature. That is the phenomenon of Thermal expansion. How does it come about? We have already mentioned that the vibrations of the molecules "become stronger" when the energy is supplied. "Stronger vibrations" mean nothing else than that the amplitudes of the vibrations become larger, i.e. that the points around which the individual molecules vibrate move apart. The immediate consequence is an increase in volume.
    • The transition from ice to liquid water is interesting: the volume becomes a little smaller when it melts. At this point, water behaves rather atypically, as most other substances increase their volume slightly when they melt ..
    • Even in the liquid state of aggregation there is a Thermal expansion instead of. Only in a small temperature range between 0 C and 3.98 C does water show a anomaly: In contrast to other liquids, the volume decreases with increasing temperature (from 1,00016 dm3 to 1.00003 dm3 - So it's one very small effect). At temperatures above 3.98 C, water behaves "normally", i.e. it expands with increasing temperature.
    • Also our water vapor expands when heated - In principle, the volume it takes up can be of any size. This is a common property of gases. In order to understand this correctly, however, it is important to remember that the supply of energy always occurs during the process we are considering at constant pressure takes place. If such a process is to be realized experimentally for a gas, it must be given the opportunity to choose its volume itself, for example by locking it in a balloon with the air pressure kept constant (the force on the gas, however, should be negligibly small). If, on the other hand, a gas is locked in a rigid container, the volume is retained and any thermal expansion is prevented from the outset.
    • Most of the energy contained in water vapor (as in any gas) is contained in the Kinetic energy of the molecules. Forces between the molecules that are associated with binding energies (which can be calculated as negative) are still present, but only play a subordinate role. Accordingly, any energy that continues to be supplied directly heats the movement of the molecules. For this reason, the thermal properties of gases are more predictable than those of solids and liquids and can be explained more easily.
    • We also note that the states of aggregation are also subordinate with regard to behavior varying pressure differ from each other: solids and liquids are relatively "volume-stable", i.e. they change their volume only slightly when they are pressurized - they cannot be easily compressed. This behavior indicates that the molecules, which in these aggregate states are, so to speak, right next to each other, are very robust. They do not "squeeze" or "press together" easily. Gases, on the other hand, are easy to compress - compressing a gas simply involves reducing the distance between the molecules.
  • Densities:
  • Comparison of water vapor with ideal gas:
    • If the forces between the molecules in a gas are negligibly small, and their size does not matter compared to their typical distance, the gas can be described with the help of a simple model. We then speak of one ideal gas.
    • As mentioned, the rise of the green line represents the specific heat capacity of water vapor at constant pressure cp designated. The slope of the steepest of the three gray lines corresponds to the (theoretically predicted) specific heat capacity of an ideal three-atom gas with the same molar mass as water (namely 18 g / mol or, to be very precise, 18.01528 g / mol). It is almost parallel to the green line, which means that water vapor can be described quite well as an ideal three-atom gas.
    • The theory to:
      • The kinetic energy contained in a three-atom ideal gas is divided equally into the energy of the translational movement of the molecules (3 degrees of freedom) and the energy of the rotation of the molecules (also 3 degrees of freedom). According to Boltzmann's famous equipartition theorem (uniform distribution theorem), there is an average kinetic energy of k TSection/ 2, where k = 1.38·10–23 J / K is the Boltzmann constant and TSection the absolute temperature is (Tabs, in K = Tin C + 273.15). For a three-atom ideal gas, this results in a total kinetic energy of 3k TSection per molecule.
      • If energy is now supplied to the gas, this kinetic energy is increased on the one hand. In addition, the gas expands (we keep the pressure constant and must therefore give the gas the opportunity to determine its volume itself), doing expansion work to the extent of k ΔT per molecule with an increase in temperature of ΔT. (Note that ΔTSection = ΔT, that I TSection and T only differ by an additive constant). Overall, the gas must therefore be supplied with an energy that is on average 4k ΔT per molecule constitutes a temperature rise of ΔT to reach. Exactly this required energy (multiplied by 3.34271025, the number of molecules per kilogram of our gas) represents the rise of the steepest of the three gray lines. It agrees very well with the rise of the green line for water vapor. Small deviations of this idealized value from the actually measured value are due to the fact that the molecules exert (small) forces on each other even in the gaseous state. At very high temperatures (higher than those considered here), the vibrations of the atoms in the water molecule can also absorb a certain amount of energy, which also increases the specific heat capacity.
      • The increases in the other two gray lines correspond to one- and two-atom gases with the same molar mass as water. In a single-atom gas, the molecules of which (= atoms) have no significant degrees of freedom of rotation, there is an average kinetic energy of 3 per atomk TSection/ 2, in a two-atom gas, the molecules of which have only two relevant degrees of freedom of rotation, there is an average kinetic energy of 5 per moleculek TSection/ 2. The work of expansion when heated under constant pressure is the same as with three-atom gas (k ΔT per molecule), so that these two gray lines have an energy input of 5k ΔT/ 2 or 7k ΔT/ 2 per molecule for an increase in temperature by ΔT represent.
Some interesting properties of water and matter in general were not mentioned here:
  • Above all, what is known from everyday experience would be here Evaporate should be mentioned, in which individual (particularly fast and therefore high-energy) molecules close to a water surface show the forces of attraction exerted by their neighbors (which are also reflected by Surface tension show) and thus escape the liquid. This process (in contrast to evaporation) takes place at any temperature, even without external energy input. As the most energetic molecules escape, the liquid becomes the Heat of evaporation withdrawn, which is equivalent to a cooling effect (which is clearly felt on wet skin, for example). An important term in this context is the Saturation vapor pressurewhich results when there is only water and water vapor in a vessel and an equilibrium has been established between the evaporation and the re-entry of molecules into the liquid.
  • These are other phenomena Sublimate and the Resublimate, i.e. the direct transition from the solid to the gaseous state and vice versa.
  • Become other pressures than the normal pressure of (slightly more than) 1 bar is permitted, the conditions change sometimes dramatically, such as the (schematic and simplified) Phase diagram the equilibrium of the water shows:

    The states occurring in the process discussed above are all on the red line. If we leave it, new phenomena appear:
    Under even more extreme conditions, the situation changes again. In the last few decades it has been discovered that ice can actually take on numerous different aggregate states, the explanation of which is the subject of current research based on the forces acting between the molecules.

Post Comment: In thermodynamics (especially in connection with its "first law") the inner energyU spoken of a substance. It includes both the disordered movement of the molecules (translational and rotational energies) and the binding energies. It changes when the substance is supplied or removed (for example through contact with a heat bath surrounding it), when mechanical work is performed on it or when it performs mechanical work on its environment. Therefore, much of the change in ΔU the inner energy spoken. However, you rarely find out - except when we are talking about ideal gases - how great the internal energy is overall (i.e. a U without Δ). This is partly due to the fact that a substance contains binding energies that are normally not considered in thermodynamics, e.g. the bonds between the core building blocks (nucleons) due to the nuclear forces. If chemical transformations are not considered, the binding energies that connect the atoms to form molecules are also irrelevant. So it is not at all easy to say how big "the inner energy" of a substance actually is. However, in order to get an impression of it, we set that state of our kilogram of water as a comparison state, that state in which the individual water molecules are far apart and do not move. So we ignore chemical energies and nuclear energies that are in the water. Let us assign the value to such a state U = 0 for the internal energy, then on the vertical axis of our diagram above, instead of the energy put into it from the starting point at -40 C E. the inner energy U be applied. The diagram then looks like this:

An imaginary extension of the green line to the left then intersects the temperature axis at absolute zero, i.e. at -273.15 C. If we ignore chemical reactions and core processes, the following picture emerges:
  • From a thermodynamic point of view, ice and liquid water below 100 C have a thermodynamic effect at normal pressure negative internal energy, i.e. the (negatively assessed) binding energies outweigh the (positive) kinetic energies of the disordered heat movement.
  • In contrast, water vapor at temperatures above 100 C has a thermodynamic effect at normal pressure positive internal energy, i.e. here the interaction energies of the water molecules can be neglected in a first approximation (which in turn is the reason why water vapor can be described quite well as an ideal gas).
  • This has an interesting consequence: If our kilograms of water vapor were left to their own devices at a temperature of over 100 C in weightlessness and without vessel walls, it could dilute itself indefinitely in terms of energy, i.e. its molecules could completely separate from one another and fly apart in different directions. This would not be possible in a state with negative internal energy (i.e. below the temperature axis). A chunk of ice left alone in space could lose individual molecules through sublimation, but it would become colder and thus bound more tightly.