Dihydrogen Monoxide Is An Example Of A Nonpolar Molecule, Water_(Molecule)

Water (H2O) Identifiers Properties Structure Hazards
IUPAC name Water
Other names AquaHydrogen oxideHydrogen hydroxideHydrateOxidaneHydroxic acidDihydrogen monoxideHydroxyl acidDihydrogen oxideHydrohydroxic acidμ-Oxido dihydrogenLight Water
CAS number 7732-18-5
RTECS number ZC0110000
Molecular formula H2O or HOH
Molar mass 18.01524g/mol
Appearance transparent, almostcolorless liquid witha slight hint of blue<1>
Density 1000kg·m−3, liquid (4°C) 917kg·m−3, solid
Melting point

0°C, 32°F (273.15K)<2>

Boiling point

100°C, 212°F (373.15K)<2>

Acidity (pKa) 15.74~35-36
Basicity (pKb) 15.74
Viscosity 0.001Pa·s at 20°C
Crystal structure HexagonalSee ice
Molecular shape non-linear bent
Dipole moment 1.85D
Main hazards water intoxication, drowning (see also Dihydrogen monoxide hoax)
NFPA 704 0
Related Compounds Related solvents acetonemethanol Related compounds water vaporiceheavy water Except where noted otherwise, data are given for materials in their standard state(at 25°C, 100kPa)Infobox disclaimer and references

Water (H2O, HOH) is the most abundant molecule on Earth”s surface, composing of about 70% of the Earth”s surface as liquid and solid state in addition to being found in the atmosphere as a vapor. It is in dynamic equilibrium between the liquid and vapor states at standard temperature and pressure. At room temperature, it is a nearly colorless with a hint of blue, tasteless, and odorless liquid. Many substances dissolve in water and it is commonly referred to as the universal solvent. Because of this, water in nature and in use is rarely pure, and may have some properties different from those in the laboratory. However, there are many compounds that are essentially, if not completely, insoluble in water. Water is the only common substance found naturally in all three common states of matter—for other substances, see Chemical properties.

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2 Physics and chemistry of water2.1 Water, ice and vapor2.3 Electrical properties2.4 Dipolar nature of water4 Systematic naming


Forms of water

See the Water#Overview of types of water

Water can take many forms. The solid state of water is commonly known as ice (while many other forms exist; see amorphous solid water); the gaseous state is known as water vapor (or steam, though this is actually incorrect, since steam is just condensing liquid water droplets), and the common liquid phase is generally taken as simply water. Above a certain critical temperature and pressure (647 K and 22.064 MPa), water molecules assume a supercritical condition, in which liquid-like clusters float within a vapor-like phase.

Heavy water is water in which the hydrogen is replaced by its heavier isotope, deuterium. It is chemically almost identical to normal water. Heavy water is used in the nuclear industry to slow down neutrons.

Physics and chemistry of water

Water is the chemical substance with chemical formula H2O: one molecule of water has two hydrogen atoms covalently bonded to a single oxygen atom. Water is a tasteless, odorless liquid at ambient temperature and pressure, and appears colorless in small quantities, although it has its own intrinsic very light blue hue. Ice also appears colorless, and water vapor is essentially invisible as a gas.<3>Water is primarily a liquid under standard conditions, which is not predicted from its relationship to other analogous hydrides of the oxygen family in the periodic table, which are gases such as hydrogen sulfide. Also the elements surrounding oxygen in the periodic table, nitrogen, fluorine, phosphorus, sulfur and chlorine, all combine with hydrogen to produce gases under standard conditions. The reason that oxygen hydride (water) forms a liquid is that it is more electronegative than all of these elements (other than fluorine). Oxygen attracts electrons much more strongly than hydrogen, resulting in a net positive charge on the hydrogen atoms, and a net negative charge on the oxygen atom. The presence of a charge on each of these atoms gives each water molecule a net dipole moment. Electrical attraction between water molecules due to this dipole pulls individual molecules closer together, making it more difficult to separate the molecules and therefore raising the boiling point. This attraction is known as hydrogen bonding. Water can be described as a polar liquid that dissociates disproportionately into the hydronium ion (H3O+(aq)) and an associated hydroxide ion (OH−(aq)).Water is in dynamic equilibrium between the liquid, gas and solid states at standard temperature and pressure, and is the only pure substance found naturally on Earth to be so.

Water, ice and vapor

Heat capacity and heat of vaporization
Main article: Enthalpy of vaporization

Water has the second highest specific heat capacity of any known chemical compound, after ammonia, as well as a high heat of vaporization (40.65 kJ mol−1), both of which are a result of the extensive hydrogen bonding between its molecules. These two unusual properties allow water to moderate Earth”s climate by buffering large fluctuations in temperature.

Density of water and ice

Temp (°C)Density (g/cm³)
30 0.9956502
25 0.9970479
22 0.9977735
20 0.9982071
15 0.9991026
10 0.9997026
4 0.9999720
0 0.9998395
−10 0.998117
−20 0.993547
−30 0.983854
The density of water in grams per cubic centimeter at various temperatures in degrees Celsius <4> The values below 0 °C refer to supercooled water.

The solid form of most substances is more dense than the liquid phase; thus, a block of pure solid substance will sink in a tub of pure liquid substance. But, by contrast, a block of common ice will float in a tub of water because solid water is less dense than liquid water. This is an extremely important characteristic property of water. At room temperature, liquid water becomes denser with lowering temperature, just like other substances. But at 4 °C (3.98 more precisely), just above freezing, water reaches its maximum density, and as water cools further toward its freezing point, the liquid water, under standard conditions, expands to become less dense. The physical reason for this is related to the crystal structure of ordinary ice, known as hexagonal ice Ih. Water, lead, uranium, neon and silicon are some of the few materials which expand when they freeze; most other materials contract. It should be noted however, that not all forms of ice are less dense than liquid water. For example HDA and VHDA are both more dense than liquid phase pure water. Thus, the reason that the common form of ice is less dense than water is a bit non-intuitive and relies heavily on the unusual properties inherent to the hydrogen bond.

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Generally, water expands when it freezes because of its molecular structure, in tandem with the unusual elasticity of the hydrogen bond and the particular lowest energy hexagonal crystal conformation that it adopts under standard conditions. That is, when water cools, it tries to stack in a crystalline lattice configuration that stretches the rotational and vibrational components of the bond, so that the effect is that each molecule of water is pushed further from each of its neighboring molecules. This effectively reduces the density ρ of water when ice is formed under standard conditions.

The importance of this property cannot be overemphasized for its role on the ecosystem of Earth. For example, if water were more dense when frozen, lakes and oceans in a polar environment would eventually freeze solid (from top to bottom). This would happen because frozen ice would settle on the lake and riverbeds, and the necessary warming phenomenon (see below) could not occur in summer, as the warm surface layer would be less dense than the solid frozen layer below. It is a significant feature of nature that this does not occur naturally in the environment.

Nevertheless, the unusual expansion of freezing water (in ordinary natural settings in relevant biological systems), due to the hydrogen bond, from 4 °C above freezing to the freezing point offers an important advantage for freshwater life in winter. Water chilled at the surface increases in density and sinks, forming convection currents that cool the whole water body, but when the temperature of the lake water reaches 4 °C, water on the surface decreases in density as it chills further and remains as a surface layer which eventually freezes and forms ice. Since downward convection of colder water is blocked by the density change, any large body of fresh water frozen in winter will have the coldest water near the surface, away from the riverbed or lakebed. This accounts for various little known phenomena of ice characteristics as they relate to ice in lakes and “ice falling out of lakes” as described by early 20th century scientist Horatio D. Craft.

Water will freeze at 0 °C (32 °F, 273 K), however, it can be supercooled in a fluid state down to its crystal homogeneous nucleation at almost 231 K (−42 °C) <5>.

Density of saltwater and ice

The density of water is dependent on the dissolved salt content as well as the temperature of the water. Ice still floats in the oceans, otherwise they would freeze from the bottom up. However, the salt content of oceans lowers the freezing point by about 2 °C and lowers the temperature of the density maximum of water to the freezing point. That is why, in ocean water, the downward convection of colder water is not blocked by an expansion of water as it becomes colder near the freezing point. The oceans” cold water near the freezing point continues to sink. For this reason, any creature attempting to survive at the bottom of such cold water as the Arctic Ocean generally lives in water that is 4 °C colder than the temperature at the bottom of frozen-over fresh water lakes and rivers in the winter.

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As the surface of salt water begins to freeze (at −1.9 °C for normal salinity seawater, 3.5%) the ice that forms is essentially salt free with a density approximately equal to that of freshwater ice. This ice floats on the surface and the salt that is “frozen out” adds to the salinity and density of the seawater just below it, in a process known as brine rejection. This more dense saltwater sinks by convection and the replacing seawater is subject to the same process. This provides essentially freshwater ice at −1.9 °C on the surface. The increased density of the seawater beneath the forming ice causes it to sink towards the bottom.

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