 | |||
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 | |||
| Names | |||
|---|---|---|---|
| Preferred IUPAC name Water | |||
| Systematic IUPAC name Oxidane (not in common use)[3] | |||
Other names
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| Identifiers | |||
3D model (JSmol) | |||
| 3587155 | |||
| ChEBI | |||
| ChEMBL | |||
| ChemSpider |
| ||
| DrugBank | |||
| EC Number |
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| 117 | |||
| KEGG | |||
| RTECS number |
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| UNII | |||
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| Properties | |||
| H 2O | |||
| Molar mass | 18.01528(33) g/mol | ||
| Appearance | Almost colorless or white crystalline solid, almost colorless liquid, with ahint of blue, colorless gas[4] | ||
| Odor | Odorless | ||
| Density | |||
| Melting point | 0.00 °C (32.00 °F; 273.15 K)[b] | ||
| Boiling point | 99.98 °C (211.96 °F; 373.13 K)[17][b] | ||
| Solubility | Poorly soluble inhaloalkanes,aliphatic andaromatic hydrocarbons,ethers.[8] Improved solubility incarboxylates,alcohols,ketones,amines. Miscible withmethanol,ethanol,propanol,isopropanol,acetone,glycerol,1,4-dioxane,tetrahydrofuran,sulfolane,acetaldehyde,dimethylformamide,dimethoxyethane,dimethyl sulfoxide,acetonitrile. Partially miscible withdiethyl ether,methyl ethyl ketone,dichloromethane,ethyl acetate,bromine. | ||
| Vapor pressure | 3.1690 kilopascals or 0.031276 atm at 25 °C[9] | ||
| Acidity (pKa) | 13.995[10][11][a] | ||
| Basicity (pKb) | 13.995 | ||
| Conjugate acid | Hydronium H3O+ (pKa = 0) | ||
| Conjugate base | Hydroxide OH– (pKb = 0) | ||
| Thermal conductivity | 0.6065 W/(m·K)[14] | ||
Refractive index (nD) | 1.3330 (20 °C)[15] | ||
| Viscosity | 0.890 mPa·s (0.890cP)[16] | ||
| Structure | |||
| Hexagonal | |||
| C2v | |||
| Bent | |||
| 1.8546D[18] | |||
| Thermochemistry | |||
| 75.385 ± 0.05 J/(mol·K)[17] | |||
Std molar entropy(S⦵298) | 69.95 ± 0.03 J/(mol·K)[17] | ||
Std enthalpy of formation(ΔfH⦵298) | −285.83 ± 0.04 kJ/mol[8][17] | ||
Gibbs free energy(ΔfG⦵) | −237.24 kJ/mol[8] | ||
| Hazards | |||
| Occupational safety and health (OHS/OSH): | |||
Main hazards | Drowning Avalanche (as snow) Water intoxication | ||
| NFPA 704 (fire diamond) | |||
| Flash point | Non-flammable | ||
| Related compounds | |||
Otheranions | |||
Relatedsolvents | |||
| Supplementary data page | |||
| Water (data page) | |||
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |||
Water (H2O) is apolar inorganic compound that is atroom temperature a tasteless and odorlessliquid, which is nearly colorless apart froman inherent hint of blue. It is by far the most studied chemical compound[20] and is described as the "universalsolvent"[21] and the "solvent of life".[22] It is the most abundant substance on the surface ofEarth[23] and the only common substance to exist as asolid, liquid, andgas on Earth's surface.[24] It is also the third most abundant molecule in the universe (behindmolecular hydrogen andcarbon monoxide).[23]
Water molecules formhydrogen bonds with each other and are strongly polar. This polarity allows it to dissociateions in salts and bond to other polar substances such as alcohols and acids, thus dissolving them. Its hydrogen bonding causes its many unique properties, such as having a solid form less dense than its liquid form, a relatively highboiling point of 100 °C for itsmolar mass, and a highheat capacity.
Water isamphoteric, meaning that it can exhibit properties of anacid or abase, depending on the pH of the solution that it is in; it readily produces bothH+
andOH−
ions.[c] Related to its amphoteric character, it undergoesself-ionization. The product of theactivities, or approximately, the concentrations ofH+
andOH−
is a constant, so their respective concentrations are inversely proportional to each other.[25]
Water is thechemical substance withchemical formulaH
2O; onemolecule of water has twohydrogenatomscovalentlybonded to a singleoxygen atom.[26] Water is a tasteless, odorless liquid atambient temperature and pressure. Liquid water has weakabsorption bands at wavelengths of around 750 nm which cause it to appear to have a blue color.[4] This can easily be observed in a water-filled bath or wash-basin whose lining is white. Large ice crystals, as inglaciers, also appear blue.
Understandard conditions, water is primarily a liquid, unlike other analogoushydrides of the oxygen family, which are generally gaseous. This unique property of water is due tohydrogen bonding. The molecules of water are constantly moving concerning each other, and the hydrogen bonds are continually breaking and reforming at timescales faster than 200 femtoseconds (2 × 10−13 seconds).[27] However, these bonds are strong enough to create many of the peculiar properties of water, some of which make it integral to life.
Within the Earth's atmosphere and surface, theliquid phase is the most common and is the form that is generally denoted by the word "water". Thesolid phase of water is known asice and commonly takes the structure of hard, amalgamatedcrystals, such asice cubes, or loosely accumulatedgranular crystals, likesnow. Aside fromcommon hexagonal crystalline ice, other crystalline and amorphousphases of ice are known. Thegaseous phase of water is known aswater vapor (orsteam). Visible steam and clouds are formed from minute droplets of water suspended in the air.
Water also forms asupercritical fluid. Thecritical temperature is 647K and thecritical pressure is 22.064MPa. In nature, this only rarely occurs in extremely hostile conditions. A likely example of naturally occurring supercritical water is in the hottest parts of deep waterhydrothermal vents, in which water is heated to the critical temperature byvolcanicplumes and the critical pressure is caused by the weight of the ocean at the extreme depths where the vents are located. This pressure is reached at a depth of about 2200 meters: much less than the mean depth of the ocean (3800 meters).[28]
Water has a very highspecific heat capacity of 4184 J/(kg·K) at 20 °C (4182 J/(kg·K) at 25 °C)—the second-highest among all the heteroatomic species (afterammonia), as well as a highheat of vaporization (40.65 kJ/mol or 2257 kJ/kg at the normal boiling point), both of which are a result of the extensivehydrogen bonding between its molecules. These unusual properties allow water to moderate Earth'sclimate by buffering large fluctuations in temperature.Most of the additional energy stored in the climate system since 1970 has accumulated in the oceans.[29]
The specificenthalpy of fusion (more commonly known as latent heat) of water is 333.55 kJ/kg at 0 °C: the same amount of energy is required to melt ice as to warm ice from −160 °C up to its melting point or to heat the same amount of water by about 80 °C. Of common substances, only that of ammonia is higher. This property confers resistance to melting on the ice ofglaciers anddrift ice. Before and since the advent of mechanicalrefrigeration, ice was and still is in common use for retarding food spoilage.
The specific heat capacity of ice at −10 °C is 2030 J/(kg·K)[30] and the heat capacity of steam at 100 °C is 2080 J/(kg·K).[31]
.svg)
Thedensity of water is about 1 gram per cubic centimetre (62 lb/cu ft): this relationship was originally used to define the gram.[32] The density varies with temperature, but not linearly: as the temperature increases, the density rises to a peak at 3.98 °C (39.16 °F) and then decreases;[33] the initial increase is unusual because most liquids undergothermal expansion so that the density only decreases as a function of temperature. The increase observed for water from 0 °C (32 °F) to 3.98 °C (39.16 °F) and for a few other liquids[d] is described asnegative thermal expansion. Regular,hexagonal ice is also less dense than liquid water—upon freezing, the density of water decreases by about 9%.[36][e]
These peculiar effects are due to the highly directional bonding of water molecules via the hydrogen bonds: ice and liquid water at low temperature have comparatively low-density, low-energy open lattice structures. The breaking of hydrogen bonds on melting with increasing temperature in the range 0–4 °C allows for a denser molecular packing in which some of the lattice cavities are filled by water molecules.[33][37] Above 4 °C, however, thermal expansion becomes the dominant effect,[37] and water near the boiling point (100 °C) is about 4% less dense than water at 4 °C (39 °F).[36][f]
Under increasing pressure, ice undergoes a number of transitions to otherpolymorphs with higher density than liquid water, such asice II,ice III,high-density amorphous ice (HDA), andvery-high-density amorphous ice (VHDA).[38][39]
The unusual density curve and lower density of ice than of water is essential for much of the life on earth—if water were most dense at the freezing point, then in winter the cooling at the surface would lead to convective mixing. Once 0 °C are reached, the water body would freeze from the bottom up, and all life in it would be killed.[36] Furthermore, given that water is a good thermal insulator (due to its heat capacity), some frozen lakes might not completely thaw in summer.[36] As it is, the inversion of the density curve leads to a stable layering for surface temperatures below 4 °C, and with the layer of ice that floats on top insulating the water below,[40] even e.g.,Lake Baikal in centralSiberia freezes only to about 1 m thickness in winter. In general, for deep enough lakes, the temperature at the bottom stays constant at about 4 °C (39 °F) throughout the year (see diagram).[36]
.png)
The density of saltwater depends on the dissolved salt content as well as the temperature. 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 1.9 °C[41] (due tofreezing-point depression of a solvent containing a solute) and lowers the temperature of the density maximum of water to the former freezing point at 0 °C. This is why, in ocean water, the downward convection of colder water isnot 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. So creatures that live at the bottom of cold oceans like theArctic Ocean generally live in water 4 °C colder than at the bottom of frozen-overfresh water lakes and rivers.
As thesurface of saltwater begins to freeze (at −1.9 °C[41] for normal salinityseawater, 3.5%) the ice that forms is essentially salt-free, with about the same density as freshwater ice. This ice floats on the surface, and the salt that is "frozen out" adds to thesalinity and density of the seawater just below it, in a process known asbrine rejection. This denser saltwater sinks by convection and the replacing seawater is subject to the same process. This produces essentially freshwater ice at −1.9 °C[41] on the surface. The increased density of the seawater beneath the forming ice causes it to sink towards the bottom. On a large scale, the process of brine rejection and sinking cold salty water results in ocean currents forming to transport such water away from the Poles, leading to a global system of currents called thethermohaline circulation.
Water ismiscible with many liquids, includingethanol in all proportions. Water and mostoils are immiscible, usually forming layers according to increasing density from the top. This can be predicted by comparing thepolarity. Water being a relatively polar compound will tend to be miscible with liquids of high polarity such as ethanol andacetone, whereas compounds with low polarity will tend to be immiscible and poorlysoluble such as withhydrocarbons.
As a gas, water vapor is completely miscible with air. On the other hand, the maximum watervapor pressure that is thermodynamically stable with the liquid (or solid) at a given temperature is relatively low compared with total atmospheric pressure. For example, if the vapor'spartial pressure is 2% of atmospheric pressure and the air is cooled from 25 °C, starting at about 22 °C, water will start to condense, defining thedew point, and creatingfog ordew. The reverse process accounts for the fog burning off in the morning. If the humidity is increased at room temperature, for example, by running a hot shower or a bath, and the temperature stays about the same, the vapor soon reaches the pressure for phase change and then condenses out as minute water droplets, commonly referred to as steam.
A saturated gas or one with 100% relative humidity is when the vapor pressure of water in the air is at equilibrium with vapor pressure due to (liquid) water; water (or ice, if cool enough) will fail to lose mass through evaporation when exposed to saturated air. Because the amount of water vapor in the air is small, relative humidity, the ratio of the partial pressure due to the water vapor to the saturated partial vapor pressure, is much more useful. Vapor pressure above 100% relative humidity is called supersaturated and can occur if the air is rapidly cooled, for example, by rising suddenly in an updraft.[g]
Thecompressibility of water is a function of pressure and temperature. At 0 °C, at the limit of zero pressure, the compressibility is5.1×10−10 Pa−1. At the zero-pressure limit, the compressibility reaches a minimum of4.4×10−10 Pa−1 around 45 °C before increasing again with increasing temperature. As the pressure is increased, the compressibility decreases, being3.9×10−10 Pa−1 at 0 °C and 100 megapascals (1,000 bar).[42]
Thebulk modulus of water is about 2.2 GPa.[43] The low compressibility of non-gasses, and of water in particular, leads to their often being assumed as incompressible. The low compressibility of water means that even in the deep oceans at 4 kilometres (2.5 mi) depth, where pressures are 40 MPa, there is only a 1.8% decrease in volume.[43]
The bulk modulus of water ice ranges from 11.3 GPa at 0 K up to 8.6 GPa at 273 K.[44] The large change in the compressibility of ice as a function of temperature is the result of its relatively large thermal expansion coefficient compared to other common solids.
Thetemperature andpressure at which ordinary solid, liquid, and gaseous water coexist in equilibrium is atriple point of water. Since 1954, this point had been used to define the base unit of temperature, thekelvin,[45][46] but,starting in 2019, the kelvin is now defined using theBoltzmann constant, rather than the triple point of water.[47]
Due to the existence of manypolymorphs (forms) of ice, water has other triple points, which have either three polymorphs of ice or two polymorphs of ice and liquid in equilibrium.[46]Gustav Heinrich Johann Apollon Tammann in Göttingen produced data on several other triple points in the early 20th century. Kamb and others documented further triple points in the 1960s.[48][49][50]
| Phases in stable equilibrium | Pressure | Temperature |
|---|---|---|
| liquid water,ice Ih, and water vapor | 611.657 Pa[51] | 273.16 K (0.01 °C) |
| liquid water, ice Ih, andice III | 209.9 MPa | 251 K (−22 °C) |
| liquid water, ice III, andice V | 350.1 MPa | −17.0 °C |
| liquid water, ice V, andice VI | 632.4 MPa | 0.16 °C |
| ice Ih,Ice II, and ice III | 213 MPa | −35 °C |
| ice II, ice III, and ice V | 344 MPa | −24 °C |
| ice II, ice V, and ice VI | 626 MPa | −70 °C |
The melting point of ice is 0 °C (32 °F; 273 K) at standard pressure; however, pure liquid water can besupercooled well below that temperature without freezing if the liquid is not mechanically disturbed. It can remain in a fluid state down to its homogeneousnucleation point of about 231 K (−42 °C; −44 °F).[52] The melting point of ordinary hexagonal ice falls slightly under moderately high pressures, by 0.0073 °C (0.0131 °F)/atm[h] or about 0.5 °C (0.90 °F)/70 atm[i][53] as the stabilization energy of hydrogen bonding is exceeded by intermolecular repulsion, but as ice transforms into its polymorphs (seecrystalline states of ice) above 209.9 MPa (2,072 atm), the melting point increases markedlywith pressure, i.e., reaching 355 K (82 °C) at 2.216 GPa (21,870 atm) (triple point ofIce VII[54]).
Pure water containing no exogenousions is an excellent electronicinsulator, but not even "deionized" water is completely free of ions. Water undergoesautoionization in the liquid state when two water molecules form one hydroxide anion (OH−
) and one hydronium cation (H
3O+
). Because of autoionization, at ambient temperatures pure liquid water has a similar intrinsic charge carrier concentration to the semiconductor germanium and an intrinsic charge carrier concentration three orders of magnitude greater than the semiconductor silicon, hence, based on charge carrier concentration, water can not be considered to be a completely dielectric material or electrical insulator but to be a limited conductor of ionic charge.[55]
Because water is such a good solvent, it almost always has somesolute dissolved in it, often asalt. If water has even a tiny amount of such an impurity, then the ions can carry charges back and forth, allowing the water to conduct electricity far more readily.
It is known that the theoretical maximum electrical resistivity for water is approximately 18.2 MΩ·cm (182kΩ·m) at 25 °C.[56] This figure agrees well with what is typically seen onreverse osmosis,ultra-filtered and deionized ultra-pure water systems used, for instance, in semiconductor manufacturing plants. A salt or acid contaminant level exceeding even 100 parts per trillion (ppt) in otherwise ultra-pure water begins to noticeably lower its resistivity by up to several kΩ·m.[citation needed]
In pure water, sensitive equipment can detect a very slightelectrical conductivity of 0.05501 ± 0.0001μS/cm at 25.00 °C.[56] Water can also beelectrolyzed into oxygen and hydrogen gases but in the absence of dissolved ions this is a very slow process, as very little current is conducted. In ice, the primary charge carriers areprotons (seeproton conductor).[57] Ice was previously thought to have a small but measurable conductivity of 1×10−10 S/cm, but this conductivity is now thought to be almost entirely from surface defects, and without those, ice is an insulator with an immeasurably small conductivity.[33]
An important feature of water is its polar nature. The structure has abent molecular geometry for the two hydrogens from the oxygen vertex. The oxygen atom also has twolone pairs of electrons. One effect usually ascribed to the lone pairs is that the H–O–H gas-phase bend angle is 104.48°,[58] which is smaller than the typicaltetrahedral angle of 109.47°. The lone pairs are closer to the oxygen atom than the electronssigma bonded to the hydrogens, so they require more space. The increased repulsion of the lone pairs forces the O–H bonds closer to each other.[59]
Another consequence of itsstructure is that water is apolar molecule. Due to the difference inelectronegativity, abond dipole moment points from each H to the O, making the oxygen partially negative and each hydrogen partially positive. A large moleculardipole, points from a region between the two hydrogen atoms to the oxygen atom. The charge differences cause water molecules to aggregate (the relatively positive areas being attracted to the relatively negative areas). This attraction,hydrogen bonding, explains many of the properties of water, such as its solvent properties.[60]
Although hydrogen bonding is a relatively weak attraction compared to the covalent bonds within the water molecule itself, it is responsible for several of the water's physical properties. These properties include its relatively highmelting and boiling point temperatures: more energy is required to break the hydrogen bonds between water molecules. In contrast, hydrogen sulfide (H
2S), has much weaker hydrogen bonding due to sulfur's lower electronegativity.H
2S is a gas atroom temperature, despite hydrogen sulfide having nearly twice the molar mass of water. The extra bonding between water molecules also gives liquid water a largespecific heat capacity. This high heat capacity makes water a good heat storage medium (coolant) and heat shield.
Water molecules stay close to each other (cohesion), due to the collective action of hydrogen bonds between water molecules. These hydrogen bonds are constantly breaking, with new bonds being formed with different water molecules; but at any given time in a sample of liquid water, a large portion of the molecules are held together by such bonds.[61]
Water also has highadhesion properties because of its polar nature. On clean, smoothglass the water may form a thin film because the molecular forces between glass and water molecules (adhesive forces) are stronger than the cohesive forces.[citation needed] In biological cells andorganelles, water is in contact with membrane and protein surfaces that arehydrophilic; that is, surfaces that have a strong attraction to water.Irving Langmuir observed a strong repulsive force between hydrophilic surfaces. To dehydrate hydrophilic surfaces—to remove the strongly held layers of water of hydration—requires doing substantial work against these forces, called hydration forces. These forces are very large but decrease rapidly over a nanometer or less.[62] They are important in biology, particularly when cells are dehydrated by exposure to dry atmospheres or to extracellular freezing.[63]
Water has an unusually highsurface tension of 71.99 mN/m at 25 °C[64] which is caused by the strength of the hydrogen bonding between water molecules.[65] This allows insects to walk on water.[65]
Because water has strong cohesive and adhesive forces, it exhibits capillary action.[66] Strong cohesion from hydrogen bonding and adhesion allows trees to transport water more than 100 m upward.[65]
Water is an excellentsolvent due to its high dielectric constant.[67] Substances that mix well and dissolve in water are known ashydrophilic ("water-loving") substances, while those that do not mix well with water are known ashydrophobic ("water-fearing") substances.[68] The ability of a substance to dissolve in water is determined by whether or not the substance can match or better the strongattractive forces that water molecules generate between other water molecules. If a substance has properties that do not allow it to overcome these strong intermolecular forces, the molecules areprecipitated out from the water. Contrary to the common misconception, water and hydrophobic substances do not "repel", and the hydration of a hydrophobic surface is energetically, but not entropically, favorable.
When an ionic or polar compound enters water, it is surrounded by water molecules (hydration). The relatively small size of water molecules (~3 angstroms) allows many water molecules to surround one molecule ofsolute. The partially negative dipole ends of the water are attracted to positively charged components of the solute, and vice versa for the positive dipole ends.
In general, ionic and polar substances such asacids,alcohols, andsalts are relatively soluble in water, and nonpolar substances such as fats and oils are not. Nonpolar molecules stay together in water because it is energetically more favorable for the water molecules to hydrogen bond to each other than to engage invan der Waals interactions with non-polar molecules.
An example of an ionic solute istable salt; the sodium chloride, NaCl, separates intoNa+
cations andCl−
anions, each being surrounded by water molecules. The ions are then easily transported away from theircrystalline lattice into solution. An example of a nonionic solute istable sugar. The water dipoles make hydrogen bonds with the polar regions of the sugar molecule (OH groups) and allow it to be carried away into solution.
Thequantum tunneling dynamics in water was reported as early as 1992. At that time it was known that there are motions which destroy and regenerate the weakhydrogen bond by internal rotations of the substituent watermonomers.[69] On 18 March 2016, it was reported that the hydrogen bond can be broken by quantum tunneling in the waterhexamer. Unlike previously reported tunneling motions in water, this involved the concerted breaking of two hydrogen bonds.[70] Later in the same year, the discovery of the quantum tunneling of water molecules was reported.[71]
Water is relatively transparent tovisible light,near ultraviolet light, andfar-red light, but it absorbs mostultraviolet light,infrared light, andmicrowaves. Mostphotoreceptors andphotosynthetic pigments utilize the portion of the light spectrum that is transmitted well through water.Microwave ovens take advantage of water's opacity to microwave radiation to heat the water inside of foods. Water's light blue color is caused by weakabsorption in the red part of thevisible spectrum.[4][72]
A single water molecule can participate in a maximum of fourhydrogen bonds because it can accept two bonds using the lone pairs on oxygen and donate two hydrogen atoms. Other molecules likehydrogen fluoride, ammonia, andmethanol can also form hydrogen bonds. However, they do not show anomalousthermodynamic,kinetic, or structural properties like those observed in water because none of them can form four hydrogen bonds: either they cannot donate or accept hydrogen atoms, or there aresteric effects in bulky residues. In water, intermoleculartetrahedral structures form due to the four hydrogen bonds, thereby forming an open structure and a three-dimensional bonding network, resulting in the anomalous decrease in density when cooled below 4 °C. This repeated, constantly reorganising unit defines a three-dimensional network extending throughout the liquid. This view is based upon neutron scattering studies and computer simulations, and it makes sense in the light of the unambiguously tetrahedral arrangement of water molecules in ice structures.
However, there is an alternative theory for the structure of water. In 2004, a controversial paper fromStockholm University suggested that water molecules in the liquid state typically bind not to four but only two others; thus forming chains and rings. The term "string theory of water" (which is not to be confused with thestring theory of physics) was coined. These observations were based upon X-ray absorption spectroscopy that probed the local environment of individual oxygen atoms.[73]
The repulsive effects of the two lone pairs on the oxygen atom cause water to have abent, notlinear, molecular structure,[74] allowing it to be polar. The hydrogen–oxygen–hydrogen angle is 104.45°, which is less than the 109.47° for idealsp3 hybridization. Thevalence bond theory explanation is that the oxygen atom's lone pairs are physically larger and therefore take up more space than the oxygen atom's bonds to the hydrogen atoms.[75] Themolecular orbital theory explanation (Bent's rule) is that lowering the energy of the oxygen atom's nonbonding hybrid orbitals (by assigning them more s character and less p character) and correspondingly raising the energy of the oxygen atom's hybrid orbitals bonded to the hydrogen atoms (by assigning them more p character and less s character) has the net effect of lowering the energy of the occupied molecular orbitals because the energy of the oxygen atom's nonbonding hybrid orbitals contributes completely to the energy of the oxygen atom's lone pairs while the energy of the oxygen atom's other two hybrid orbitals contributes only partially to the energy of the bonding orbitals (the remainder of the contribution coming from the hydrogen atoms' 1s orbitals).
In liquid water there is someself-ionization givinghydronium ions andhydroxide ions.
Theequilibrium constant for this reaction, known as theionic product of water,, has a value of about 10−14 at 25 °C. At neutralpH, the concentration of thehydroxide ion (OH−
) equals that of the (solvated) hydrogen ion (H+
), with a value close to 10−7 mol L−1 at 25 °C.[76] Seedata page for values at other temperatures.
The thermodynamic equilibrium constant is a quotient ofthermodynamic activities of all products and reactants including water:
However, for dilute solutions, the activity of a solute such as H3O+ or OH− is approximated by its concentration, and the activity of the solvent H2O is approximated by 1, so that we obtain the simple ionic product
The action of water on rock over long periods of time typically leads toweathering andwater erosion, physical processes that convert solid rocks and minerals into soil and sediment, but under some conditions chemical reactions with water occur as well, resulting inmetasomatism ormineral hydration, a type of chemical alteration of a rock which producesclay minerals. It also occurs whenPortland cement hardens.
Water ice can formclathrate compounds, known asclathrate hydrates, with a variety of small molecules that can be embedded in its spacious crystal lattice. The most notable of these ismethane clathrate, 4CH
4·23H
2O, naturally found in large quantities on the ocean floor.
Rain is generally mildly acidic, with a pH between 5.2 and 5.8 if not having any acid stronger than carbon dioxide.[77] If high amounts ofnitrogen andsulfur oxides are present in the air, they too will dissolve into the cloud and raindrops, producingacid rain.
Severalisotopes of both hydrogen and oxygen exist, giving rise to several knownisotopologues of water.Vienna Standard Mean Ocean Water is the current international standard for water isotopes. Naturally occurring water is almost completely composed of the neutron-less hydrogen isotopeprotium. Only 155ppm includedeuterium (2
H or D), a hydrogen isotope with one neutron, and fewer than 20 parts perquintillion includetritium (3
H or T), which has two neutrons. Oxygen also has three stable isotopes, with16
O present in 99.76%,17
O in 0.04%, and18
O in 0.2% of water molecules.[78]
Deuterium oxide,D
2O, is also known asheavy water because of its higher density. It is used innuclear reactors as aneutron moderator. Tritium isradioactive, decaying with ahalf-life of 4500 days;THO exists in nature only in minute quantities, being produced primarily via cosmic ray-induced nuclear reactions in the atmosphere. Water with one protium and one deuterium atomHDO occur naturally in ordinary water in low concentrations (~0.03%) andD
2O in far lower amounts (0.000003%) and any such molecules are temporary as the atoms recombine.
The most notable physical differences betweenH
2O andD
2O, other than the simple difference in specific mass, involve properties that are affected by hydrogen bonding, such as freezing and boiling, and other kinetic effects. This is because the nucleus of deuterium is twice as heavy as protium, and this causes noticeable differences in bonding energies. The difference in boiling points allows the isotopologues to be separated. Theself-diffusion coefficient ofH
2O at 25 °C is 23% higher than the value ofD
2O.[79] Because water molecules exchange hydrogen atoms with one another, hydrogen deuterium oxide (DOH) is much more common in low-purity heavy water than pure dideuterium monoxideD
2O.
Consumption of pure isolatedD
2O may affect biochemical processes—ingestion of large amounts impairs kidney and central nervous system function. Small quantities can be consumed without any ill-effects; humans are generally unaware of taste differences,[80] but sometimes report a burning sensation[81] or sweet flavor.[82] Very large amounts of heavy water must be consumed for any toxicity to become apparent. Rats, however, are able to avoid heavy water by smell, and it is toxic to many animals.[83]
Light water refers to deuterium-depleted water (DDW), water in which the deuterium content has been reduced below the standard155 ppm level.
Water is the most abundant substance on Earth's surface and also the third most abundant molecule in the universe, afterH
2 andCO.[23] 0.23 ppm of the earth's mass is water and 97.39% of the global water volume of 1.38×109 km3 is found in the oceans.[84]
Water is far more prevalent in the outer Solar System, beyond a point called thefrost line, where the Sun's radiation is too weak to vaporize solid and liquid water (as well as other elements and chemical compounds with relatively low melting points, such asmethane andammonia). In the inner Solar System, planets, asteroids, and moons formed almost entirely of metals and silicates. Water has since been delivered to the inner Solar System via an as-yet unknown mechanism, theorized to be the impacts of asteroids or comets carrying water from the outer Solar System, where bodies contain much more water ice.[85] The difference between planetary bodies located inside and outside the frost line can be stark. Earth's mass is 0.000023% water, whileTethys, a moon of Saturn, is almost entirely made of water.[86]
Water isamphoteric: it has the ability to act as either anacid or abase in chemical reactions.[87] According to theBrønsted-Lowry definition, an acid is a proton (H+
) donor and a base is a proton acceptor.[88] When reacting with a stronger acid, water acts as a base; when reacting with a stronger base, it acts as an acid.[88] For instance, water receives anH+
ion from HCl whenhydrochloric acid is formed:
In the reaction withammonia,NH
3, water donates aH+
ion, and is thus acting as an acid:
Because the oxygen atom in water has twolone pairs, water often acts as aLewis base, or electron-pair donor, in reactions withLewis acids, although it can also react with Lewis bases, forming hydrogen bonds between the electron pair donors and the hydrogen atoms of water.HSAB theory describes water as both a weak hard acid and a weak hard base, meaning that it reacts preferentially with other hard species:
When a salt of a weak acid or of a weak base is dissolved in water, water can partiallyhydrolyze the salt, producing the corresponding base or acid, which gives aqueous solutions ofsoap andbaking soda their basic pH:
Water's Lewis base character makes it a commonligand intransition metal complexes, examples of which includemetal aquo complexes such asFe(H
2O)2+
6 toperrhenic acid, which contains two water molecules coordinated to arhenium center. In solidhydrates, water can be either a ligand or simply lodged in the framework, or both. Thus,FeSO
4·7H
2O consists of [Fe(H2O)6]2+ centers and one "lattice water". Water is typically amonodentate ligand, i.e., it forms only one bond with the central atom.[89]
As a hard base, water reacts readily with organiccarbocations; for example in ahydration reaction, a hydroxyl group (OH−
) and an acidic proton are added to the two carbon atoms bonded together in the carbon-carbon double bond, resulting in an alcohol. When the addition of water to an organic molecule cleaves the molecule in two,hydrolysis is said to occur. Notable examples of hydrolysis are thesaponification of fats and thedigestion of proteins andpolysaccharides. Water can also be aleaving group inSN2 substitution andE2 elimination reactions; the latter is then known as adehydration reaction.
Water contains hydrogen in theoxidation state +1 and oxygen in the oxidation state −2.[90] It oxidizes chemicals such ashydrides,alkali metals, and somealkaline earth metals.[91][92] One example of an alkali metal reacting with water is:[93]
Some other reactive metals, such asaluminium andberyllium, are oxidized by water as well, but their oxides adhere to the metal and form apassive protective layer.[94] Note that therusting ofiron is a reaction between iron and oxygen[95] that is dissolved in water, not between iron and water.
Water can be oxidized to emit oxygen gas, but very few oxidants react with water even if their reduction potential is greater than the potential ofO
2/H
2O. Almost all such reactions require acatalyst.[96] An example of the oxidation of water is:
Water can be split into its constituent elements, hydrogen and oxygen, by passing an electric current through it.[97] This process is called electrolysis. The cathode half reaction is:
The anode half reaction is:
The gases produced bubble to the surface, where they can be collected or ignited with a flame above the water if this was the intention. The required potential for the electrolysis of pure water is 1.23 V at 25 °C.[97] The operating potential is actually 1.48 V or higher in practical electrolysis.
Henry Cavendish showed that water was composed of oxygen and hydrogen in 1781.[98] The first decomposition of water into hydrogen and oxygen, byelectrolysis, was done in 1800 by English chemistWilliam Nicholson andAnthony Carlisle.[98][99] In 1805,Joseph Louis Gay-Lussac andAlexander von Humboldt showed that water is composed of two parts hydrogen and one part oxygen.[100]
Gilbert Newton Lewis isolated the first sample of pureheavy water in 1933.[101]
The properties of water have historically been used to define varioustemperature scales. Notably, theKelvin,Celsius,Rankine, andFahrenheit scales were, or currently are, defined by the freezing and boiling points of water. The less common scales ofDelisle,Newton,Réaumur, andRømer were defined similarly. Thetriple point of water is a more commonly used standard point today.
The acceptedIUPAC name of water isoxidane or simplywater,[102] or its equivalent in different languages, although there are other systematic names which can be used to describe the molecule.Oxidane is only intended to be used as the name of the mononuclearparent hydride used for naming derivatives of water bysubstituent nomenclature.[103] These derivatives commonly have other recommended names. For example, the namehydroxyl is recommended overoxidanyl for the –OH group. The nameoxane is explicitly mentioned by the IUPAC as being unsuitable for this purpose, since it is already the name of a cyclic ether also known astetrahydropyran.[3][104]
The simplest systematic name of water ishydrogen oxide. This is analogous to related compounds such ashydrogen peroxide,hydrogen sulfide, anddeuterium oxide (heavy water). Using chemical nomenclature fortype I ionic binary compounds, water would take the namehydrogen monoxide,[105] but this is not among the names published by theInternational Union of Pure and Applied Chemistry (IUPAC).[102] Another name isdihydrogen monoxide, which is a rarely used name of water, and mostly used in thedihydrogen monoxide parody.
Other systematic names for water includehydroxic acid,hydroxylic acid, andhydrogen hydroxide, using acid and base names.[j] None of these exotic names are used widely. The polarized form of the water molecule,H+
OH−
, is also calledhydron hydroxide by IUPAC nomenclature.[106]
Water substance is a rare term used for H2O when one does not wish to specify the phase of matter (liquid water,water vapor, some form ofice, or a component in a mixture) though the termwater is also used with this general meaning.
Oxygen dihydride is another way of referring to water, but modern usage often restricts the term "hydride" to ionic compounds (which water is not).
Sometimes these compounds have generic or common names (e.g., H2O is "water") and they also have systematic names (e.g., H2O, dihydrogen monoxide).
Ocean warming dominates the global energy change inventory. Warming of the ocean accounts for about 93% of the increase in the Earth's energy inventory between 1971 and 2010 (high confidence), with the warming of the upper (0 to 700 m) ocean accounting for about 64% of the total. Melting ice (including Arctic sea ice, ice sheets, and glaciers) and warming of the continents and atmosphere account for the remainder of the change in energy.
Gramme, le poids absolu d'un volume d'eau pure égal au cube de la centième partie du mètre, et à la température de la glace fondante.
Water, H2O, is similar. It has two electron pairs with nothing attached to them. They, too, must be taken into account. Molecules like NH3 and H2O are calledbent.
Notice that the bond angles decrease as the number of nonbonding electron pairs increases. A bonding pair of electrons is attracted by both nuclei of the bonded atoms, but a nonbonding pair is attracted primarily by only one nucleus. Because a nonbonding pair experiences less nuclear attraction, its electron domain is spread out more in space than is the electron domain for a bonding pair (Figure 9.7). Nonbonding electron pairs, therefore, take up more space than bonding pairs; in essence, they act as large and fatter balloons in our analogy of Figure 9.5. As a result,electron domains for nonbonding electron pairs exert greater repulsive forces on adjacent electron domains and tend to compress bond angles
