Ice

Ice is water that is frozen into a solid state, typically forming at or below temperatures of 0 °C, 32 °F, or 273.15 K. It occurs naturally on Earth, on other planets, in Oort cloud objects, and as interstellar ice. As a naturally occurring crystalline inorganic solid with an ordered structure, ice is considered to be a mineral. Depending on the presence of impurities such as particles of soil or bubbles of air, it can appear transparent or a more or less opaque bluish-white color.
Virtually all of the ice on Earth is of a hexagonal crystalline structure denoted as ice I_h. Depending on temperature and pressure, at least nineteen phases can exist. The most common phase transition to ice I_h occurs when liquid water is cooled below at standard atmospheric pressure. When water is cooled rapidly, up to three types of amorphous ice can form. Interstellar ice is overwhelmingly low-density amorphous ice, which likely makes LDA ice the most abundant type in the universe. When cooled slowly, correlated proton tunneling occurs below giving rise to macroscopic quantum phenomena.
Ice is abundant on the Earth's surface, particularly in the polar regions and above the snow line, where it can aggregate from snow to form glaciers and ice sheets. As snowflakes and hail, ice is a common form of precipitation, and it may also be deposited directly by water vapor as frost. The transition from ice to water is melting and from ice directly to water vapor is sublimation. These processes play a key role in Earth's water cycle and climate. In the recent decades, ice volume on Earth has been decreasing due to climate change. The largest declines have occurred in the Arctic and in the mountains located outside of the polar regions. The loss of grounded ice is the primary contributor to sea level rise.
Humans have been using ice for various purposes for thousands of years. Some historic structures designed to hold ice to provide cooling are over 2,000 years old. Before the invention of refrigeration technology, the only way to safely store food without modifying it through preservatives was to use ice. Sufficiently solid surface ice makes waterways accessible to land transport during winter, and dedicated ice roads may be maintained. Ice also plays a major role in winter sports.

Physical properties

Ice possesses a regular crystalline structure based on the molecule of water, which consists of a single oxygen atom covalently bonded to two hydrogen atoms, or H–O–H. However, many of the physical properties of water and ice are controlled by the formation of hydrogen bonds between adjacent oxygen and hydrogen atoms; while it is a weak bond, it is nonetheless critical in controlling the structure of both water and ice.
An unusual property of water is that its solid form—ice frozen at atmospheric pressure—is approximately 8.3% less dense than its liquid form; this is equivalent to a volumetric expansion of 9%. The density of ice is 0.9167–0.9168 g/cm³ at 0 °C and standard atmospheric pressure, whereas water has a density of 0.9998–0.999863 g/cm³ at the same temperature and pressure. Liquid water is densest, essentially 1.00 g/cm³, at 4 °C and begins to lose its density as the water molecules begin to form the hexagonal crystals of ice as the freezing point is reached. This is due to hydrogen bonding dominating the intermolecular forces, which results in a packing of molecules less compact in the solid. The density of ice increases slightly with decreasing temperature and has a value of 0.9340 g/cm³ at −180 °C.
When water freezes, it increases in volume. The effect of expansion during freezing can be dramatic, and ice expansion is a basic cause of freeze-thaw weathering of rock in nature and damage to building foundations and roadways from frost heaving. It is also a common cause of the flooding of houses when water pipes burst due to the pressure of expanding water when it freezes.
Because ice is less dense than liquid water, it floats, and this prevents bottom-up freezing of the bodies of water. Instead, a sheltered environment for animal and plant life is formed beneath the floating ice, which protects the underside from short-term weather extremes such as wind chill. Sufficiently thin floating ice allows light to pass through, supporting the photosynthesis of bacterial and algal colonies. When sea water freezes, the ice is riddled with brine-filled channels which sustain sympagic organisms such as bacteria, algae, copepods and annelids. In turn, they provide food for animals such as krill and specialized fish like the bald notothen, fed upon in turn by larger animals such as emperor penguins and minke whales.
When ice melts, it absorbs as much energy as it would take to heat an equivalent mass of water by. During the melting process, the temperature remains constant at. While melting, any energy added breaks the hydrogen bonds between ice molecules. Energy becomes available to increase the thermal energy only after enough hydrogen bonds are broken that the ice can be considered liquid water. The amount of energy consumed in breaking hydrogen bonds in the transition from ice to water is known as the heat of fusion.
As with water, ice absorbs light at the red end of the spectrum preferentially as the result of an overtone of an oxygen–hydrogen bond stretch. Compared with water, this absorption is shifted toward slightly lower energies. Thus, ice appears blue, with a slightly greener tint than liquid water. Since absorption is cumulative, the color effect intensifies with increasing thickness or if internal reflections cause the light to take a longer path through the ice. Other colors can appear in the presence of light absorbing impurities, where the impurity is dictating the color rather than the ice itself. For instance, icebergs containing impurities can appear brown, grey or green.
Because ice in natural environments is usually close to its melting temperature, its hardness shows pronounced temperature variations. At its melting point, ice has a Mohs hardness of 2 or less, but the hardness increases to about 4 at a temperature of and to 6 at a temperature of, the vaporization point of solid carbon dioxide.

Phases

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, water freezes at a temperature below. Ice, water, and water vapour can coexist at the triple point, which is exactly at a pressure of 611.657 Pa. The kelvin was defined as of the difference between this triple point and absolute zero, though this definition changed in May 2019. Unlike most other solids, ice is difficult to superheat. In an experiment, ice at −3 °C was superheated to about 17 °C for about 250 picoseconds.
Subjected to higher pressures and varying temperatures, ice can form in nineteen separate known crystalline phases at various densities, along with hypothetical proposed phases of ice that have not been observed. With care, at least fifteen of these phases 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. Ice XII was discovered in 1996. In 2006, Ice XIII and Ice XIV were discovered. Ices XI, XIII, and XIV are hydrogen-ordered forms of ices I, V, and XII respectively. In 2009, ice XV was found at extremely high pressures and −143 °C. At even higher pressures, ice is predicted to become a metal; this has been variously estimated to occur at 1.55 TPa or 5.62 TPa.
As well as crystalline forms, solid water can exist in amorphous states as amorphous solid water of varying densities. In outer space, hexagonal crystalline ice is present in the ice volcanoes, but is extremely rare otherwise. Even icy moons like Ganymede are expected to mainly consist of other crystalline forms of ice. Water in the interstellar medium is dominated by amorphous ice, making it likely the most common form of water in the universe. Low-density ASW, also known as hyperquenched glassy water, may be responsible for noctilucent clouds on Earth and is usually formed by deposition of water vapor in cold or vacuum conditions. High-density ASW is formed by compression of ordinary ice I or LDA at GPa pressures. Very-high-density ASW is HDA slightly warmed to 160 K under 1–2 GPa pressures.
Ice from a theorized superionic water may possess two crystalline structures. At pressures in excess of such superionic ice would take on a body-centered cubic structure. However, at pressures in excess of the structure may shift to a more stable face-centered cubic lattice. It is speculated that superionic ice could compose the interior of ice giants such as Uranus and Neptune.

Friction properties

Ice is "slippery" because it has a low coefficient of friction. This subject was first scientifically investigated in the 19th century. The preferred explanation at the time was "pressure melting" – i.e. the blade of an ice skate, upon exerting pressure on the ice, would melt a thin layer, providing sufficient lubrication for the blade to glide across the ice. Yet, research in 1939 by Frank P. Bowden and T. P. Hughes found that skaters would experience a lot more friction than they actually do if it were the only explanation. Further, the optimum temperature for figure skating is and for ice hockey; yet, according to pressure melting theory, skating below would be outright impossible. Instead, Bowden and Hughes argued that heating and melting of the ice layer is caused by friction. However, this theory does not sufficiently explain why ice is slippery when standing still even at below-zero temperatures.
Subsequent research suggested that ice molecules at the interface cannot properly bond with the molecules of the mass of ice beneath. These molecules remain in a semi-liquid state, providing lubrication regardless of pressure against the ice exerted by any object. However, the significance of this hypothesis is disputed by experiments showing a high coefficient of friction for ice using atomic force microscopy. Thus, the mechanism controlling the frictional properties of ice is still an active area of scientific study. A comprehensive theory of ice friction must take into account all of the aforementioned mechanisms to estimate friction coefficient of ice against various materials as a function of temperature and sliding speed. Research in 2014 suggests that frictional heating is the most important process under most typical conditions.