A gas can have a boiling point
The boiling point (Abbreviation: Sdp) or also Boiling point (Abbr .: Kp) of a basic substance is a pair of values in its phase diagram and consists of two variables: The saturation temperature (especially also Boiling temperature) and the saturation vapor pressure (especially Boiling pressure) at the phase boundary between gas and liquid. It is made up of the two state variables pressure and temperature when a substance changes from a liquid to a gaseous state of aggregation.
The boiling point represents the conditions that exist during the phase transition of a substance from the liquid to the gaseous phase, which is referred to as boiling or evaporation. In addition, it is identical to the condensation point for the reverse process of condensation, but only for pure substances. When a mixture of substances evaporates, the boiling behavior changes and a boiling range is observed instead of a single boiling point. In the case of a phase transition from the liquid to the gaseous phase below the boiling point, one speaks of evaporation.
In tables, the boiling temperatures are given at normal pressure, i.e. at 1013.25 hPa. This boiling point is called Normal boiling point, the specified boiling temperature as Normal boiling temperature (TBoil) designated. The term boiling point is often used as a short form for the Normal boiling temperature is used and is therefore mostly synonymous in common parlance, which, however, would reduce the boiling point to a single pair of values and is therefore formally incorrect.
With a pressure cooker, for example, you make use of the fact that the boiling temperature and the boiling pressure are interdependent. By increasing the pressure by usually one bar (1000 hPa), the boiling temperature of the water can be increased from 100 ° C to around 120 ° C. Both temperatures represent boiling temperatures, but only the value of 100 ° C is also the boiling temperature under normal pressure and thus the normal boiling temperature. Mixing the two terms is therefore unspecific, by no means self-evident and should be avoided.
Main articles: evaporation, heat of evaporation and delayed boiling
Below and above the boiling point, heating the liquid or gas only leads to an increase in temperature. The supplied energy is converted into kinetic energy of the particles. During the phase transition of the liquid to the gas, however, the temperature remains constant, provided that the pressure also remains constant. All thermal energy supplied is invested in the change of state.
Once the boiling point has been reached, the chemical-physical interactions between the particles are dissolved when further energy is supplied - the particles pass into the gas phase. The temperature of the liquid stagnates because the thermal energy supplied is used entirely to break the intermolecular bonds. The energy that is required for one mole of the substance is also known as the enthalpy of vaporization and its counterpart, which is not related to the amount of material, is the heat of vaporization. Only when all the particles are in the gas phase does the temperature of the system rise again.
Water, hydrogen peroxide or alkalis (e.g. caustic soda) without dust particles or gas bubbles can also be heated above the boiling point in clean vessels without boiling. The smallest disturbances, such as vibrations that result in thorough mixing, can lead to an explosive separation of the liquid from the vapor phase, which is known as Delayed boiling designated. Because of this, so-called liquids are added in chemistry that are at risk of delayed boiling Boiling stones made of clay or pumice stone, which are not attacked by the chemical, but facilitate the formation of small bubbles due to their porous structure so that there is no delay in boiling.
See also : Evaporation, gasification, evaporation, transpiration, Pictet-Trouton rule
Boiling point curve
All temperature-pressure value pairs at the gas-liquid phase boundary line in a phase diagram, taken together, result in the Boiling point curve, with a thermodynamic equilibrium prevailing on it. The boiling point curve is often referred to as Boiling curve, Boiling line, Boiling pressure curve or Boiling point curve. This curve is limited by two points:
- Triple point: Is the pressure-temperature value pair lower than the triple temperature or the triple pressure, only a transition between solid and gaseous state, i.e. sublimation or resublimation, is possible.
- Critical point: Is the pressure-temperature value pair higher than the critical temperature or the critical pressure, there is no longer any difference between the density of the liquid and that of the gaseous state, which is why they are no longer separated by a phase boundary line and the substance is therefore in this state as supercritical fluid designated.
The equilibrium of the boiling point curve is a dynamic equilibrium. Particles from a liquid constantly pass into the gas phase - they evaporate. On the other hand, these particles also re-enter the liquid phase - they condense. The numerical ratio of the particles emerging from the liquid phase and the particles re-entering it depends on both the temperature and the pressure: the higher the temperature, the more particles evaporate due to their higher speed (see Maxwell-Boltzmann distribution) . The more particles evaporate, the higher the vapor pressure and the more particles condense again. An equilibrium is established when exactly as many particles enter the gas phase as return to the liquid phase. Since the gas phase is saturated in this state, one also speaks of the saturation vapor pressure. The thermodynamic law from which the boiling point curve is quantitatively derived is known as the Clausius-Clapeyron equation. For water, this relationship between saturation vapor pressure and saturation temperature can also be determined using the approximation equations of the Magnus formula.
Change of equilibrium using the example of water
Starting point: water is in equilibrium with its gas phase at the boiling point Sdp (74 ° C, 333 hPa):
The reactions of the system to the changes in individual state variables result in a shift in the equilibrium position: the phase transition that reverses the disturbance takes place more intensely (see principle of the smallest constraint).
- If the system is cooled to 53 ° C, the vapor pressure of the gas phase is too high and water vapor condenses until the vapor pressure reaches the new equilibrium value of 143 hPa or there is no more gaseous water.
- If the system is heated to 95 ° C, the vapor pressure of the gas phase is too low and water evaporates until the vapor pressure reaches the new equilibrium value of 845 hPa or there is no more liquid water left.
- If the pressure is increased from 333 to 560 hPa while the temperature remains the same, the vapor pressure of the gas phase is too high and gaseous water condenses until the vapor pressure has the old equilibrium value of 333 hPa or there is no more water vapor left.
- If the pressure is reduced from 333 to 65 hPa while the temperature remains the same, the vapor pressure of the gas phase is too low and water evaporates until the vapor pressure has the old equilibrium value of 333 hPa or there is no more liquid water left.
Substance dependence of the boiling point
- The boiling point depends on the molar mass or molecular mass of the substance. The following applies: the greater the molar mass, the higher the boiling point. If one compares, for example, the series HCl (36 g / mol) - HBr (81 g / mol) - HI (128 g / mol) on the dark blue line, this relationship can be clearly seen. Explanation: The greater the mass of a particle, the more kinetic energy it needs to be able to pass into the gas phase.
- The boiling point is also dependent on the strength of the binding forces between the smallest particles in the liquid phase: the stronger the binding forces, the higher the boiling point, as these would first have to be overcome. This becomes clear when comparing HF and HCl, for example: In liquid hydrogen fluoride, the molecules form hydrogen bonds, while in liquid hydrogen chloride the weaker dipole-dipole interactions predominate. The same applies to the comparatively very high boiling point of water, which becomes clear when you compare this with carbon dioxide and take into account the influence of the molar masses.
- The observation that substances have a higher boiling point than similar substances with a higher molar mass is called a boiling point anomaly.
- Van der Waals interactions are even weaker than dipole-dipole interactions. For this reason, when comparing, all hydrogen compounds of the elements of main group IV have the lowest boiling points.
- The strength of the intermolecular binding forces also depends on the geometry of the molecules. See the boiling points of the homologous series of hydrocarbons or alcohols.
Examples of normal boiling points of pure substances
- The lowest normal boiling temperature of all elements with -269 ° C has helium, although it has a larger molar mass than hydrogen with a normal boiling temperature of -253 ° C. This is due to the fact that the hydrogen molecule is somewhat easier to polarize than helium and therefore also forms somewhat stronger van der Waals interactions.
- Tungsten has the highest normal boiling temperature at 5555 ° C.
- A group comparison of noble gases, non-metals, semimetals and metals shows that metals have a significantly higher boiling point than non-metals, since the metal bond (in addition to the ionic and atomic bonds) is the strongest bond. Exceptions:
- Mercury has a normal boiling temperature of 357 ° C, which is unusually low for metals
- Carbon has an extremely high boiling point of 4827 ° C for non-metals.
Carbon monoxide has one of the lowest normal boiling temperatures at –191.6 ° C, while metal carbides such as titanium (IV) carbide (TiC, 4820 ° C) and tungsten (IV) carbide (WC, 6000 ° C) have the highest.
A special feature is a modification of sulfur trioxide (SO3) before: here the normal boiling temperature of 44.8 ° C is lower than the normal melting temperature of 62.3 ° C.
If the critical pressure is below normal pressure, no normal boiling temperature can be specified. In order to still bring the liquid to the boil, this must be done under lower pressure. In this case, when specifying the boiling temperature, the boiling pressure must also be specified, which is another reason for strictly separating the terms normal boiling temperature and boiling point.
If the pressure of the triple point is above normal pressure, the normal sublimation temperature or a boiling temperature at a higher boiling pressure is specified instead of the normal boiling temperature. Example: sulfur hexafluoride SF6 sublimates under normal pressure at -63 ° C.
Many organic compounds, especially macromolecular ones, decompose when heated before the boiling point is reached. Their intermolecular bonds are stronger than the bonds within the molecule. You cannot specify a boiling point here, only the decomposition temperature. Example: sulfuric acid decomposes at 340 ° C before the boiling process begins.
Homogeneous multi-component systems
The boiling points of homogeneous mixtures such as alloys, gas mixtures or aqueous solutions have different boiling points and a different boiling behavior compared to the pure substances.
Boiling point increase
If a substance is dissolved in a solvent, the boiling point of the mixture increases in comparison to the pure solvent; one speaks of the dissolution effect in relation to the saturation vapor pressure. This is simply due to the fact that the particles of the dissolved substance hinder the transition of the solvent particles into the gas phase. According to Raoult's law by François Marie Raoult (1830-1901), this increase is ΔTSdp proportional to the amount of substance in the solute:
The individual symbols represent the following values:
- ΔTSdp - Boiling point increase
- Ke - ebullioscopic constant
- b - Molality of the solute
- K - molar increase in boiling point
- n - Amount of substance
As stated, the proportionality factor is either the ebullioscopic constant (also Boiling point constantKS.), i.e. the change in the boiling point of one kilogram of the solution compared to the pure solvent, whereby the amount of substance of the dissolved substance is one mole or the molar increase in boiling point, which is less common and does not make a statement about the mass.
For example, the boiling point of one kilogram of water rises by 0.51 K to 100.51 ° C if exactly one mole of any other substance is dissolved in it, provided the substance dissolves in water and is not volatile. If two moles are dissolved in one kilogram of water, the water only boils at (100 + 2 · 0.51 ° C) = 101.02 ° C.
It should be noted that salts dissociate in aqueous solution. Table salt, for example, breaks down into the ions Na+ and Cl-. The increase in boiling point is therefore (in dilute solutions) twice as high as initially expected.
A practical example: pasta water has a typical salt content of 10 g / kg. With a molar mass of 58.4 g / mol, this, together with the doubling mentioned above, corresponds to 0.34 mol / kg of ions. The salt content results in a boiling point increase of only about 0.17 K.
Raoult's law only applies to "ideal" solutions, that is, solutions in which a substance is only physically dissolved. In the case of “non-ideal” solutions, energetic phenomena (heating or cooling) occur during mixing, which are caused by the formation of hydrogen bonds or by protolysis. This results in deviations from Raoult's law. Only in a very strong dilution does the formula also apply to "non-ideal" solutions in approximation, which is why in the case of the ideal solution one should also use a infinitely dilute solution speaks. The increase in boiling point is also a colligative property and therefore depends on the number of particles in the dissolved substance, but not on its type. By changing the above formula, the increase in boiling point can also be used to determine molar mass, which is known as ebullioscopy.
The melting point is also dependent on the concentration of the dissolved substances, which is why one speaks of a lowering of the melting point. The cause of these effects is also an increase in the chemical potential. If you combine an increase in the boiling point and a decrease in the melting point, an overall expansion of the thermodynamic range of states of the liquid at the expense of the other states of aggregation can be seen.
If a mixture (= a uniform distribution process to be described by the increase in entropy) is heated, it begins to boil when the temperature has the boiling point of the component that has the lowest boiling point. When boiling, the particles of this component increasingly pass into the gas phase. This changes the composition of the mixture and its boiling point changes continuously. This rise in temperature does not end until the boiling point of the component with the highest boiling point has been reached. In this case, too, one speaks of a Boiling range (also boiling interval, boiling limit) of the mixture and no longer of one boiling point. The dependence of the state of aggregation and the composition of mixtures on the temperature is shown in boiling diagrams:
Example: If a liquid mixture contains equal parts nitrogen and oxygen, the boiling range is between –191.5 ° C and –183 ° C.
In the case of azeotropic substance mixtures, the boiling point of the substance mixture is higher or lower than the boiling point of the two pure substance components at a certain molar ratio. At this mixing ratio there is a boiling point and no boiling range.
Importance to living beings
Under the physical conditions on earth, the boiling behavior of water means that water exists in large quantities as a liquid. This is one of the basic requirements for the development of living beings.
With a lower air pressure or higher water temperatures, this would of course be different and would result in bodies of water evaporating within a very short period of time and thus an important condition for life in general, namely liquid water, would be encountered much less often. At a higher air pressure or a lower temperature, however, less and less water would be able to evaporate, and thus the prerequisite for precipitation, namely gaseous water in the atmosphere, would become increasingly rare, which would, for example, result in a restriction of fresh water resources.
- Analytical chemistry: the boiling point is a specific property of a substance. In this way, pure substances can be characterized based on their boiling point.
- Distillation or fractional distillation, a method for separating a mixture of substances based on the different boiling points of the individual components. The low-boiling substance is separated from the higher-boiling substance by evaporation.
- Ebullioscopy (lat. bulla = Boiling bubble, gr. skopein = consider) is a method for determining the molar mass by increasing the boiling point. Since increases in the boiling point are smaller than decreases in the freezing point, cryoscopy is usually preferred. Both methods use a special thermometer, which was developed in 1888 by Ernst Beckmann (1853-1923): the Beckmann thermometer. It has a scale that is only about 6 °, but it can also be read with an accuracy of 0.01 degrees. The zero point of the scale can be set to the required temperature.
- Pressure cooker: If the water in the airtight saucepan is heated to over 100 ° C, the boiling point and pressure of the water increase. This results in faster cooking.
- Altitude measurement: As the air pressure decreases with increasing altitude, the boiling point also decreases. As a rule of thumb, the boiling point is lowered by about one degree for every 300 m. By determining the boiling temperature of pure water, it is possible to estimate the respective altitude above mean sea level.
Categories: Threshold (temperature) | Substance property
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