Ice
Ice is water in a solid state, one of the three natural states of water that are part of the four states of aggregation of matter. It is recognized by its temperature, its snowy white color and its buoyancy. Pure water freezes at 0 °C when under one atmosphere of pressure.
Ice is the common name for water in a solid state; Other denominations are snow, frost, hail, etc.
Volume and density of water when it solidifies
Water is one of the few substances that increases in volume when frozen (therefore its density decreases); that is, it expands. This property prevents the oceans of the Earth's polar regions from freezing in their entire volume, since the ice floats in the water and is what is exposed to changes in temperature in the atmosphere. The typical density of ice at 0 °C is usually taken as 0.916 g/cm³; or as 916.8 kg/m³.
Crystal structure
Ice occurs in 12 different crystalline structures or phases. At the usual pressures in the terrestrial medium (in the neighborhood of atmospheric pressure), the stable phase is usually denoted as phase I according to Tamman's terminology. This phase I presents two interrelated variants: hexagonal ice, denoted Ih, and cubic ice, Ic. Hexagonal ice is the most common phase, and the best known: its hexagonal structure can be seen reflected in ice crystals, which always have a hexagonal base. Cubic ice Ic is obtained by deposition of water vapor at temperatures below –130 °C, so it is not as common; Even so, at about –38 °C and 200 MPa of pressure, a situation to be expected in the polar caps, both structures are in thermodynamic equilibrium.
Ih ice has a hexagonal structure in which each oxygen atom in a water molecule has four other hydrogen atoms as its nearest neighbors, located at the vertices of a regular tetrahedron whose center is the oxygen atom in interest. This tetrahedral unit is common to all other phases of ice, and is due to the fact that the angle between hydrogen atoms in the free water molecule H-O-H is 104.52º, instead of 90º. The tetrahedral angle between O-O-O is 109.47º. For terrestrial temperatures of interest, the distance between oxygen atoms O-O is 0.276 nm and between O-H is 0.0985 nm. The union between intramolecular atoms is by simple covalent bonds and therefore very stable, while the intermolecular union is produced by relatively weak hydrogen bonds, which explains the relatively low melting temperature of ice. The most relevant lattice parameters are the hexagonal side a=0.451 nm, and the height of the hexagonal prism c=0.7357 nm. These values may vary slightly with temperature, but the relationship between them, c/a=1.628, remains practically stable and very close to the optimal value of c/a=1.633, theorized for solid spheres in contact forming the same hexagonal structure. The stability of the parameter c/a explains the fact that the thermal expansion of ice occurs in an isotropic manner. For its part, the fact that ice Ih has a hexagonal structure explains the anisotropy usually observed in its mechanical properties: Young's modulus, for example, which is around E=9-10GPa for pure crystals, presents radial isotropy, and varies considerably depending on the direction of strain; the mechanical resistance, located in the neighborhood of 1MPa for pure crystals in the basal direction, can reach 7MPa in certain configurations. The presence of impurities in the network is prula, except for some specific substances such as ammonium fluoride, NH4F. There can be four crystalline defects: vacant, interstitial, ionic or Bjerrum, the last two being exclusive to ice and being related to the rotation of hydrogens from a water molecule in the lattice.
In any case, the structure Ih of ice is not very compact —which explains its lower density with respect to the liquid phase— especially when compared to similar structures in other crystalline materials such as metals. The packing factor is 0.34, much lower than the 0.74 typical for metals. This is explained by the repulsion of hydrogen and oxygen atoms as the lattice is compacted. In fact, this repulsion leads to the fact that, when the pressure on the hexagonal network is high enough, this structure ceases to be stable and others appear to replace it.
Indeed, the rest of the crystalline phases are produced at much higher pressures, and until 1900 they were unknown. In fact, they do not exist on Earth, since the terrestrial polar caps are too thin to allow the appearance of stable phases other than ice Ih. However, the situation is different in the large icy moons of the solar system such as Europa or Triton, where it is postulated that the pressures in the core are high enough to ensure the appearance of stable phases other than Ih, which at said pressures it would be unstable. The best known high pressure crystalline phases are phases II and III; in the laboratory only phases II, III, V and VI have been studied, while the rest remain basically unknown.
The structure of ice II is rhomboid. This ice forms at around 238K for pressures of 283 atmospheres, and its density is 1193 kg/m³ because it is a much more compact structure. Ice III is tetragonal, and appears at about 246 K and 276 atm, with a density of 1166 kg/m³. Ice V is monoclinic, occurring at 237.5 K and 480 atm, with a density of 1267 kg/m³. Ice VI is tetragonal, and appears at 237.5K for 777atm, with a density of 1360 kg/m³. All these phases are essentially brittle, although they show a great tendency to creep over time (creep) and some viscoelastic behavior.
Although initially believed to be nanocrystalline phases, apart from the crystalline phases mentioned above, ice can appear in two amorphous (vitreous) phases: low-density amorphous ice (940 kg/m³ at –196 °C and 1 atm) and high-density amorphous ice (1170 kg/m³, same conditions). The formation of amorphous ice is complicated, and is related to the solidification time given to the water; can be formed by condensation of vapor below –160 °C, by collapse of structure Ih under elevated pressure below –196 °C. In any case, except in certain very specific situations, they are not common phases on Earth.
Ice as a mineral
In mineralogy it is accepted as a valid mineral by the International Mineralogical Association, since it is a stable solid at temperatures below 0 °C. It is classified in group 4 of oxide minerals as it is an oxide of hydrogen, normally with abundant impurities.
Types of ice
As in most solids, in ice the molecules are arranged in an orderly array. However, depending on the pressure and temperature conditions, it is possible that they adopt different forms of ordering. Starting in 1900, Gustave Tamman and later in 1912 Percy Bridgman carried out experiments on ice applying different pressures and temperatures, and obtained different ices with greater densities than normal (later many more types of ice were found). All these forms of ice have more compact structures (different forms of an element existing in the same physical state), that is, several allotropic modifications or allotropes are formed.
The known types of ice are:
- MDA o Middle-density amorphous ice, is characterized by not having a crystallized structure and being more similar in molecular structure to liquid water than to any other type of known ice. It also has the same density as liquid water.
- Ice Ih (All ice formed in the terrestrial biosphere is Ih-type ice, except for a small amount of Ic ice. Ice crystals have hexagonal shape.
- Ice Ic (low temperature, face-centered cubic, density approximately 900 kg/m3).
- Ice II (low temperature, centered orthrombic, density approximately 1200 kg/m3).
- Ice III or Iii (low temperature, tetragonal, density approximately 1100 kg/m3).
- Ice V (high pressure, low temperature, centered base monoclinics, density approximately 1200 kg/m3).
- Ice VI (high pressure, low temperature, tetragonal, density approximately 1300 kg/m3).
- Ice VII (high temperature, high pressure, simple cubic, density approximately 1700 kg/m3).
- Ice VIII (high pressure, centered tetragonal, density approximately 1600 kg/m3).
- Ice IX (high pressure, tetragonal, density approximately 1200 kg/m3).
- Ice XII (high pressure, low temperature, tetragonal, density approximately 1300 kg/m3).
Eutectic Fusion
Under terrestrial conditions and during winter it is common to add salt to the ice so that it melts. In fact, what melts is not ice, but a compound of ice and salt called "eutectic". When the salt NaCl (Na+, Cl–) comes into contact with ice, the ions arrange themselves around the water molecules, which are polar (H2δ+, Oδ–) and forms a compound (H2O).(NaCl). For this rearrangement, only small movements of atoms are needed, and it is therefore done in the solid phase. When the exact proportions are respected (around 23% salt by mass), you have a product that behaves like a pure product (in particular, there is a constant melting temperature) and is described as "eutectic". The melting temperature of this eutectic is around –21 °C.
If the proportion of salt is lower than this ratio, a water-eutectic mixture is given, which melts at a higher temperature (between –21 °C and 0 °C). If the proportion of salt is higher, you have a salt-eutectic mixture that also melts at a higher temperature. You can draw a diagram, called a phase diagram, which represents the melting temperature as a function of the water-salt ratios.
Water-sal phase diagram atmospheric pressure; euthectic is formed in a water ratio of 0.32331 in mass (23.31 % of salt and 76.69 % of mass water).
The "arrangement" water + salt → eutectic can only occur at the points of contact between the ice and salt crystals, that is, on the surface of the ice. This forms a surface layer of eutectic that melts (if the temperature is higher than –21 °C). Because the salt is supersaturated, it dissolves in the molten eutectic and can react with ice on the liquid film. The phenomenon then propagates until there is no water or salt to form a new eutectic. Thus, in theory, it would be possible to prevent ice formation down to –21 °C. In practice, it is impossible to dose the amount of salt to be used.
Cryoscopic descent
Cryoscopic depression is the lowering of the melting point of a pure solvent by the presence of solutes. It is directly proportional to molality, which makes it more important for ionic solutes, such as those found in seawater, than for non-ionics. The phenomenon has important consequences in the case of seawater, because the response to intense cooling of ocean water, as occurs in the winter of polar regions, is the separation of a floating solid phase of pure water in the form of ice. This is how ice packs form around Antarctica or the Arctic Ocean, as a compact aggregate of pure water ice, with brine filling the interstices, and floating on a body of liquid water at less than 0 °C (up to a limit of –1.9 °C for a salinity of 3.5 %).
The color of ice
Sometimes the ice appears blue.
The white light of the Sun is actually made up of a mixture of colors, from red to violet, as seen when a ray of light is passed through a glass prism, or in a rainbow. Bluer light waves have more energy than yellow or red waves. Snow is white because all the light that hits it is reflected by a very thin layer on its surface. The tiny air bubbles that are trapped in the ice reflect light multiple times and all colors from red to violet escape, so the light we receive is white light. Ice appears blue when it has a very high consistency and air bubbles do not block light from passing through it. Without the "scattering" effect of the bubbles, light can penetrate the ice and is gradually absorbed on its way to the deeper parts. Red photons, which have lower energy than blue ones, penetrate less distance and are absorbed earlier. On average, red light absorption in ice is six times more efficient than blue light absorption; therefore, the more distance a beam of white light travels, it loses more and more red, yellow, green photons on its way… and it is the blue ones that “survive”. This is the reason for the blue color of pure ice, and of a glacier or iceberg.
In other words, the most compact ice, such as glacial ice, behaves in a special way when it receives light. When a ray of light hits, only the blue component of solar radiation has enough energy to penetrate the interior of the ice mass. Therefore, by absorbing the other colors, the glacial ice appears blue.
Other meanings
By extension, the name ice is used for other types of chemical compounds. Thus, for example, we speak of dry ice to refer to the solid state of carbon dioxide (carbon dioxide or CO2).
Erosive action
The expansion of water as it solidifies has important geological effects. Water that gets into tiny cracks in rocks on Earth's surface creates an enormous amount of pressure as it solidifies and splits or breaks the rocks. This ice action plays an important role in erosion. In addition, glaciers, through friction, polish the terrain where they circulate.
Ice formations without water
The solid phases of some other substances are also called ice, especially in the astrophysical context: dry ice is a commonly used term for solid carbon dioxide.
A magnetic monopole of ice can also be realized by isolating magnetic materials in which the magnetic moments mimic the position of protons in water ice and are governed by energetic constraints similar to the Bernal-Fowler rules, derived from geometric frustration in the configuration of a proton in water ice. These materials are called spin ice.
Volume of water when melted
When a floating ice mass melts, the water level does not rise. This is because at all times the volume of water that an ice mass can potentially generate as it melts is equal to the volume of ice that is submerged at that time.
The weight of a mass of ice is:
P=ρ ρ hVhielorg{displaystyle P=rho _{h}V_{hielo}g}
The Archimedean push it receives is:
E=ρ ρ aVsg{displaystyle E=rho _{a}V_{s}g}
Equating both terms making balance of forces we arrive at:
Vs=ρ ρ hρ ρ aVhielor{displaystyle V_{s}={frac {rho _{h}}{rho _{a}}}}V_{hielo}}}
Which relates the submerged volume to the total volume of existing ice.
To show that the volume of ice that melts is equal to the volume that is submerged first we apply conservation of mass:
Mhielor=ρ ρ hVhielor=ρ ρ aVderretidor{displaystyle M_{hielo}=rho _{h}V_{hielo}=rho _{a}V_{derretido}}}
Introducing the values calculated above we arrive at:
Vderretidor=ρ ρ hρ ρ a↓ ↓ ρ ρ aρ ρ hVs=Vs{displaystyle V_{derretido}={frac {rho _{h}}{rho _{a}}}{frac {rho _{a}{rho _{h}}}}}{V_{s}=V_{s}}}
Proving volume equivalence.
If the water level does not change... Why do they say that melting the poles will increase the level of the oceans?
We have seen that if a floating ice mass melts it will not affect the water level. However, if the melting ice is on land, yes. Antarctica, for example, is a continent completely covered by ice that rests on land. Therefore, if the ice melts and/or some pieces of ice break off into the ocean, the water level will rise.