A<unk>  molecule consists of two hydrogen atoms and one oxygen atom<unk>  () is a polar inorganic compound that is at room temperature a tasteless and odorless liquid, nearly colorless with a hint of blue. This simplest hydrogen chalcogenide is by far the most studied chemical compound and is described as the "universal solvent" for its ability to dissolve many substances. This allows it to be the "solvent of life": indeed,<unk>  as found in nature almost always includes various dissolved substances, and special steps are required to obtain chemically pure<unk> .<unk>  is the only common substance to exist as a solid, liquid, and gas in normal terrestrial conditions. States The three common states of matter Along with oxidane,<unk>  is one of the two official names for the chemical compound ; it is also the liquid phase of . The other two common states of matter of<unk>  are the solid phase, ice, and the gaseous phase,<unk>  vapor or steam. The addition or removal of heat can cause phase transitions: freezing<unk>  to ice), melting (ice to<unk> , vaporization<unk>  to vapor), condensation (vapor to<unk> , sublimation (ice to vapor) and deposition (vapor to ice). Density<unk>  differs from most liquids in that it becomes less dense as it freezes. In 1 atm pressure, it reaches its maximum density of at . The density of ice is , an expansion of 9%. This expansion can exert enormous pressure, bursting pipes and cracking rocks (see Frost weathering). In a lake or ocean,<unk>  at 4 °C (39.2 °F) sinks to the bottom, and ice forms on the surface, floating on the liquid<unk> . This ice insulates the<unk>  below, preventing it from freezing solid. Without this protection, most<unk> uatic organisms would perish during the winter. Magnetism<unk>  is a diamagnetic<unk> ial. Though interaction is weak, with superconducting magnets it can attain a notable interaction. Phase transitions At a pressure of one atmosphere (atm), ice melts or<unk>  freezes at 0 °C (32 °F) and<unk>  boils or vapor condenses at 100 °C (212 °F). However, even below the boiling point,<unk>  can change to vapor at its surface by evaporation (vaporization throughout the liquid is known as boiling). Sublimation and deposition also occur on surfaces. For example, frost is deposited on cold surfaces while snowflakes form by deposition on an aerosol particle or ice nucleus. In the process of freeze-drying, a food is frozen and then stored at low pressure so the ice on its surface sublimates. The melting and boiling points depend on pressure. A good approximation for the rate of change of the melting temperature with pressure is given by the Clausius–Clapeyron relation: : where and are the molar volumes of the liquid and solid phases, and is the molar latent heat of melting. In most substances, the volume increases when melting occurs, so the melting temperature increases with pressure. However, because ice is less dense than<unk> , the melting temperature decreases. In glaciers, pressure melting can occur under sufficiently thick volumes of ice, resulting in subglacial lakes. The Clausius-Clapeyron relation also applies to the boiling point, but with the liquid/gas transition the vapor phase has a much lower density than the liquid phase, so the boiling point increases with pressure.<unk>  can remain in a liquid state at high temperatures in the deep ocean or underground. For example, temperatures exceed in Old Faithful, a geyser in Yellowstone National Park. In hydrothermal vents, the temperature can exceed . At sea level, the boiling point of<unk>  is . As atmospheric pressure decreases with altitude, the boiling point decreases by 1 °C every 274 meters. High-altitude cooking takes longer than sea-level cooking. For example, at , cooking time must be increased by a fourth to achieve the desired result. (Conversely, a pressure cooker can be used to decrease cooking times by raising the boiling temperature.) In a vacuum,<unk>  will boil at room temperature. Triple and critical points Phase diagram of<unk>  (simplified) On a pressure/temperature phase diagram (see figure), there are curves separating solid from vapor, vapor from liquid, and liquid from solid. These meet at a single point called the triple point, where all three phases can coexist. The triple point is at a temperature of and a pressure of ; it is the lowest pressure at which liquid<unk>  can exist. Until 2019, the triple point was used to define the Kelvin temperature scale. The<unk> /vapor phase curve terminates at and . This is known as the critical point. At higher temperatures and pressures the liquid and vapor phases form a continuous phase called a supercritical fluid. It can be gradually compressed or expanded between gas-like and liquid-like densities, its properties (which are quite different from those of ambient<unk>  are sensitive to density. For example, for suitable pressures and temperatures it can mix freely with nonpolar compounds, including most organic compounds. This makes it useful in a variety of applications including high-temperature electrochemistry and as an ecologically benign solvent or catalyst in chemical reactions involving organic compounds. In Earth's mantle, it acts as a solvent during mineral formation, dissolution and deposition. Phases of ice and<unk>  The normal form of ice on the surface of Earth is Ice Ih, a phase that forms crystals with hexagonal symmetry. Another with cubic crystalline symmetry, Ice Ic, can occur in the upper atmosphere. As the pressure increases, ice forms other crystal structures. As of 2019, 17 have been experimentally confirmed and several more are predicted theoretically. The 18th form of ice, ice XVIII, a face-centred-cubic, superionic ice phase, was discovered when a droplet of<unk>  was subject to a shock wave that raised the<unk> ’s pressure to millions of atmospheres and its temperature to thousands of degrees, resulting in a structure of rigid oxygen atoms in which hydrogen atoms flowed freely. When sandwiched between layers of graphene, ice forms a square lattice. The details of the chemical nature of liquid<unk>  are not well understood; some theories suggest that its unusual behaviour is due to the existence of 2 liquid states. Taste and odor Pure<unk>  is usually described as tasteless and odorless, although humans have specific sensors that can feel the presence of<unk>  in their mouths, and frogs are known to be able to smell it. However,<unk>  from ordinary sources (including bottled mineral<unk>  usually has many dissolved substances, that may give it varying tastes and odors. Humans and other animals have developed senses that enable them to evaluate the potability of<unk>  by avoiding<unk>  that is too salty or putrid. Color and appearance Pure<unk>  is visibly blue due to absorption of light in the region ca. 600 nm – 800 nm. The color can be easily observed in a glass of tap-water placed against a pure white background, in daylight. The principal absorption bands responsible for the color are overtones of the O–H stretching vibrations. The apparent intensity of the color increases with the depth of the<unk>  column, following Beer's law. This also applies, for example, with a swimming pool when the light source is sunlight reflected from the pool's white tiles. In nature, the color may also be modified from blue to green due to the presence of suspended solids or algae. In industry, near-infrared spectroscopy is used with<unk> ueous solutions as the greater intensity of the lower overtones of<unk>  means that glass cuvettes with short path-length may be employed. To observe the fundamental stretching absorption spectrum of<unk>  or of an<unk> ueous solution in the region around 3500 cm−1 (2.85 μm) a path length of about 25 μm is needed. Also, the cuvette must be both transparent around 3500 cm−1 and insoluble in<unk> ; calcium fluoride is one<unk> ial that is in common use for the cuvette windows with<unk> ueous solutions. The Raman-active fundamental vibrations may be observed with, for example, a 1 cm sample cell.<unk> uatic plants, algae, and other photosynthetic organisms can live in<unk>  up to hundreds of meters deep, because sunlight can reach them. Practically no sunlight reaches the parts of the oceans below of depth. The refractive index of liquid<unk>  (1.333 at ) is much higher than that of air (1.0), similar to those of alkanes and ethanol, but lower than those of glycerol (1.473), benzene (1.501), carbon disulfide (1.627), and common types of glass (1.4 to 1.6). The refraction index of ice (1.31) is lower than that of liquid<unk> . Polar molecule Tetrahedral structure of<unk>  In a<unk>  molecule, the hydrogen atoms form a 104.5° angle with the oxygen atom. The hydrogen atoms are close to two corners of a tetrahedron centered on the oxygen. At the other two corners are lone pairs of valence electrons that do not participate in the bonding. In a perfect tetrahedron, the atoms would form a 109.5° angle, but the repulsion between the lone pairs is greater than the repulsion between the hydrogen atoms. The O–H bond length is about 0.096
Water
Carbonated water
Tar water
Protein 
Gripe water
Saline 
Juice
Raw water