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Phases of ice
Variations in pressure and temperature give rise to different phases of ice, which have varying properties and molecular geometries. Currently, twenty-two crystalline phases have been observed, including ice Ih, Ic, ..., XXI . In modern history, phases have been discovered through scientific research with various techniques including pressurization, force application, nucleation agents, and others.
On Earth, most ice is found in the hexagonal Ice Ih phase. Less common phases may be found in the atmosphere and underground due to more extreme pressures and temperatures. Some phases are manufactured for nano scale uses due to their properties. In space, amorphous ice is the most common form as confirmed by observation. Thus, it is theorized to be the most common phase in the universe. Various other phases could be found naturally in astronomical objects.
Most liquids under increased pressure freeze at higher temperatures because the pressure helps to hold the molecules together. However, the strong hydrogen bonds in water make it different: for some pressures higher than 0.10 MPa (1 atm), water freezes at a temperature below 0 °C. Subjected to higher pressures and varying temperatures, ice can form in nineteen separate known crystalline phases. With care, at least fifteen of these phases (one of the known exceptions being ice X) can be recovered at ambient pressure and low temperature in metastable form. The types are differentiated by their crystalline structure, proton ordering, and density. There are also two metastable phases of ice under pressure, both fully hydrogen-disordered; these are Ice IV and Ice XII.
The accepted crystal structure of ordinary ice was first proposed by Linus Pauling in 1935. The structure of ice Ih is the wurtzite lattice, roughly one of crinkled planes composed of tessellating hexagonal rings, with an oxygen atom on each vertex, and the edges of the rings formed by hydrogen bonds. The planes alternate in an ABAB pattern, with B planes being reflections of the A planes along the same axes as the planes themselves. The distance between oxygen atoms along each bond is about 275 pm and is the same between any two bonded oxygen atoms in the lattice. The angle between bonds in the crystal lattice is very close to the tetrahedral angle of 109.5°, which is also quite close to the angle between hydrogen atoms in the water molecule (in the gas phase), which is 105°.
This tetrahedral bonding angle of the water molecule essentially accounts for the unusually low density of the crystal lattice – it is beneficial for the lattice to be arranged with tetrahedral angles even though there is an energy penalty in the increased volume of the crystal lattice. As a result, the large hexagonal rings leave almost enough room for another water molecule to exist inside. This gives naturally occurring ice its rare property of being less dense than its liquid form. The tetrahedral-angled hydrogen-bonded hexagonal rings are also the mechanism that causes liquid water to be densest at 4 °C. Close to 0 °C, tiny hexagonal ice Ih-like lattices form in liquid water, with greater frequency closer to 0 °C. This effect decreases the density of the water, causing it to be densest at 4 °C when the structures form infrequently.
In the best-known form of ice, ice Ih, the crystal structure is characterized by the oxygen atoms forming hexagonal symmetry with near tetrahedral bonding angles. This structure is stable down to −268 °C (5 K; −450 °F), as evidenced by x-ray diffraction and extremely high resolution thermal expansion measurements. Ice Ih is also stable under applied pressures of up to about 210 megapascals (2,100 atm) where it transitions into ice III or ice II.
While most forms of ice are crystalline, several amorphous (or "vitreous") forms of ice also exist. Such ice is an amorphous solid form of water, which lacks long-range order in its molecular arrangement. Amorphous ice is produced either by rapid cooling of liquid water to its glass transition temperature (about 136 K or −137 °C) in milliseconds (so the molecules do not have enough time to form a crystal lattice), or by compressing ordinary ice at low temperatures. The most common form on Earth, low-density ice, is usually formed in the laboratory by a slow accumulation of water vapor molecules (physical vapor deposition) onto a very smooth metal crystal surface under 120 K. In outer space it is expected to be formed in a similar manner on a variety of cold substrates, such as dust particles. By contrast, hyperquenched glassy water is formed by spraying a fine mist of water droplets into a liquid such as propane around 80 K, or by hyperquenching fine micrometer-sized droplets on a sample-holder kept at liquid nitrogen temperature, 77 K, in a vacuum. Cooling rates above 104 K/s are required to prevent crystallization of the droplets. At liquid nitrogen temperature, 77 K, hyperquenched glassy water is kinetically stable and can be stored for many years.
Amorphous ices have the property of suppressing long-range density fluctuations and are, therefore, nearly hyperuniform. Classification analysis suggests that low and high density amorphous ices are glasses.
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Phases of ice
Variations in pressure and temperature give rise to different phases of ice, which have varying properties and molecular geometries. Currently, twenty-two crystalline phases have been observed, including ice Ih, Ic, ..., XXI . In modern history, phases have been discovered through scientific research with various techniques including pressurization, force application, nucleation agents, and others.
On Earth, most ice is found in the hexagonal Ice Ih phase. Less common phases may be found in the atmosphere and underground due to more extreme pressures and temperatures. Some phases are manufactured for nano scale uses due to their properties. In space, amorphous ice is the most common form as confirmed by observation. Thus, it is theorized to be the most common phase in the universe. Various other phases could be found naturally in astronomical objects.
Most liquids under increased pressure freeze at higher temperatures because the pressure helps to hold the molecules together. However, the strong hydrogen bonds in water make it different: for some pressures higher than 0.10 MPa (1 atm), water freezes at a temperature below 0 °C. Subjected to higher pressures and varying temperatures, ice can form in nineteen separate known crystalline phases. With care, at least fifteen of these phases (one of the known exceptions being ice X) can be recovered at ambient pressure and low temperature in metastable form. The types are differentiated by their crystalline structure, proton ordering, and density. There are also two metastable phases of ice under pressure, both fully hydrogen-disordered; these are Ice IV and Ice XII.
The accepted crystal structure of ordinary ice was first proposed by Linus Pauling in 1935. The structure of ice Ih is the wurtzite lattice, roughly one of crinkled planes composed of tessellating hexagonal rings, with an oxygen atom on each vertex, and the edges of the rings formed by hydrogen bonds. The planes alternate in an ABAB pattern, with B planes being reflections of the A planes along the same axes as the planes themselves. The distance between oxygen atoms along each bond is about 275 pm and is the same between any two bonded oxygen atoms in the lattice. The angle between bonds in the crystal lattice is very close to the tetrahedral angle of 109.5°, which is also quite close to the angle between hydrogen atoms in the water molecule (in the gas phase), which is 105°.
This tetrahedral bonding angle of the water molecule essentially accounts for the unusually low density of the crystal lattice – it is beneficial for the lattice to be arranged with tetrahedral angles even though there is an energy penalty in the increased volume of the crystal lattice. As a result, the large hexagonal rings leave almost enough room for another water molecule to exist inside. This gives naturally occurring ice its rare property of being less dense than its liquid form. The tetrahedral-angled hydrogen-bonded hexagonal rings are also the mechanism that causes liquid water to be densest at 4 °C. Close to 0 °C, tiny hexagonal ice Ih-like lattices form in liquid water, with greater frequency closer to 0 °C. This effect decreases the density of the water, causing it to be densest at 4 °C when the structures form infrequently.
In the best-known form of ice, ice Ih, the crystal structure is characterized by the oxygen atoms forming hexagonal symmetry with near tetrahedral bonding angles. This structure is stable down to −268 °C (5 K; −450 °F), as evidenced by x-ray diffraction and extremely high resolution thermal expansion measurements. Ice Ih is also stable under applied pressures of up to about 210 megapascals (2,100 atm) where it transitions into ice III or ice II.
While most forms of ice are crystalline, several amorphous (or "vitreous") forms of ice also exist. Such ice is an amorphous solid form of water, which lacks long-range order in its molecular arrangement. Amorphous ice is produced either by rapid cooling of liquid water to its glass transition temperature (about 136 K or −137 °C) in milliseconds (so the molecules do not have enough time to form a crystal lattice), or by compressing ordinary ice at low temperatures. The most common form on Earth, low-density ice, is usually formed in the laboratory by a slow accumulation of water vapor molecules (physical vapor deposition) onto a very smooth metal crystal surface under 120 K. In outer space it is expected to be formed in a similar manner on a variety of cold substrates, such as dust particles. By contrast, hyperquenched glassy water is formed by spraying a fine mist of water droplets into a liquid such as propane around 80 K, or by hyperquenching fine micrometer-sized droplets on a sample-holder kept at liquid nitrogen temperature, 77 K, in a vacuum. Cooling rates above 104 K/s are required to prevent crystallization of the droplets. At liquid nitrogen temperature, 77 K, hyperquenched glassy water is kinetically stable and can be stored for many years.
Amorphous ices have the property of suppressing long-range density fluctuations and are, therefore, nearly hyperuniform. Classification analysis suggests that low and high density amorphous ices are glasses.