Activity 31 a Quick Review of the Properties of Water Answers

Concrete and chemic properties of pure water

For broader coverage of this topic, see Water.

Water

Names
IUPAC name Water
Systematic IUPAC name Oxidane
Other names Hydrogen hydroxide (HH or HOH), hydrogen oxide, dihydrogen monoxide (DHMO) (systematic proper noun[1]), dihydrogen oxide, hydric acid, hydrohydroxic acrid, hydroxic acid, hydrol,[2] μ-oxido dihydrogen, κ1-hydroxyl hydrogen(0)
Identifiers
CAS Number	7732-18-5 Y
3D model (JSmol)	Interactive paradigm
Beilstein Reference	3587155
ChEBI	CHEBI:15377 Y
ChEMBL	ChEMBL1098659 Y
ChemSpider	937 Y
Gmelin Reference	117
PubChem CID	962
RTECS number	ZC0110000
UNII	059QF0KO0R Y
InChI InChI=1S/H2O/h1H2 Y Central: XLYOFNOQVPJJNP-UHFFFAOYSA-N Y
SMILES O
Properties
Chemical formula	H 2 O
Molar mass	18.01528(33) g/mol
Advent	White crystalline solid, almost colorless liquid with a hint of blue, colorless gas[3]
Aroma	None
Density	Liquid:[4] 0.9998396 g/mL at 0 °C 0.9970474 thousand/mL at 25 °C 0.961893 1000/mL at 95 °C Solid:[v] 0.9167 g/ml at 0 °C
Melting point	0.00 °C (32.00 °F; 273.15 1000) [b]
Boiling indicate	99.98 °C (211.96 °F; 373.13 K)[fifteen] [b]
Solubility in water	N/A
Solubility	Poorly soluble in haloalkanes, aliphatic and aromatic hydrocarbons, ethers.[6] Improved solubility in carboxylates, alcohols, ketones, amines. Miscible with methanol, ethanol, propanol, isopropanol, acetone, glycerol, 1,4-dioxane, tetrahydrofuran, sulfolane, acetaldehyde, dimethylformamide, dimethoxyethane, dimethyl sulfoxide, acetonitrile. Partially miscible with Diethyl ether, Methyl Ethyl Ketone, Dichloromethane, Ethyl Acetate, Bromine.
Vapor pressure	3.1690 kilopascals or 0.031276 atm at 25 °C[7]
Acidity (pK a)	13.995[viii] [nine] [a]
Basicity (pK b)	13.995
Cohabit acid	Hydronium H3O+ (pKa = 0)
Conjugate base	Hydroxide OH– (pKb = 0)
Thermal conductivity	0.6065 W/(m·Thousand)[12]
Refractive alphabetize (due north D)	ane.3330 (twenty °C)[thirteen]
Viscosity	0.890 mPa·due south (0.890 cP)[14]
Structure
Crystal structure	Hexagonal
Point grouping	C2v
Molecular shape	Aptitude
Dipole moment	ane.8546 D[16]
Thermochemistry
Heat capacity (C)	75.385 ± 0.05 J/(mol·Yard)[17]
Std tooth entropy (S o 298)	69.95 ± 0.03 J/(mol·1000)[17]
Std enthalpy of formation (Δf H ⦵ 298)	−285.83 ± 0.04 kJ/mol[six] [17]
Gibbs free energy (Δf G˚)	−237.24 kJ/mol[6]
Hazards
Occupational safety and health (OHS/OSH):
Main hazards	Drowning Barrage (equally snowfall) Water intoxication (see also Dihydrogen monoxide parody)
GHS labelling:[eighteen]
Hazard statements	[ ? ]
NFPA 704 (fire diamond)	0 0 0
Flash point	Non-combustible
Safety data sheet (SDS)	SDS
Related compounds
Other cations	Hydrogen sulfide Hydrogen selenide Hydrogen telluride Hydrogen polonide Hydrogen peroxide
Related solvents	Acetone Methanol
Supplementary data page
Water (data page)
Except where otherwise noted, data are given for materials in their standard land (at 25 °C [77 °F], 100 kPa).
Yverify (what is Y N  ?)
Infobox references

Chemical chemical compound

Water (H
_two O) is a polar inorganic compound that is at room temperature a tasteless and odorless liquid, which is nearly colorless apart from an inherent hint of blue. Information technology is by far the most studied chemical chemical compound^[nineteen] and is described as the "universal solvent"^[20] and the "solvent of life".^[21] It is the most abundant substance on the surface of Earth^[22] and the merely common substance to exist as a solid, liquid, and gas on Earth'southward surface.^[23] It is as well the 3rd nearly abundant molecule in the universe (behind molecular hydrogen and carbon monoxide).^[22]

Water molecules form hydrogen bonds with each other and are strongly polar. This polarity allows information technology to dissociate ions in salts and bond to other polar substances such as alcohols and acids, thus dissolving them. Its hydrogen bonding causes its many unique backdrop, such as having a solid course less dumbo than its liquid form,^[c] a relatively high boiling indicate of 100 °C for its molar mass, and a high heat capacity.

Water is amphoteric, meaning that it can showroom properties of an acid or a base, depending on the pH of the solution that it is in; it readily produces both H ⁺
and OH ⁻
ions.^[c] Related to its amphoteric character, it undergoes cocky-ionization. The product of the activities, or approximately, the concentrations of H ⁺
and OH ⁻
is a constant, so their respective concentrations are inversely proportional to each other.^[24]

Concrete properties [edit]

Water is the chemic substance with chemical formula H
₂ O; i molecule of water has two hydrogen atoms covalently bonded to a single oxygen atom.^[25] Water is a tasteless, odorless liquid at ambient temperature and pressure. Liquid water has weak assimilation bands at wavelengths of around 750 nm which cause it to appear to have a blue colour.^[3] This tin can easily be observed in a water-filled bath or wash-basin whose lining is white. Large ice crystals, as in glaciers, too appear blue.

Under standard conditions, water is primarily a liquid, unlike other analogous hydrides of the oxygen family, which are generally gaseous. This unique property of water is due to hydrogen 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).^[26] Even so, these bonds are strong plenty to create many of the peculiar properties of water, some of which arrive integral to life.

Water, ice, and vapour [edit]

Within the Earth's atmosphere and surface, the liquid phase is the most common and is the grade that is mostly denoted by the discussion "water". The solid phase of water is known as water ice and commonly takes the construction of hard, amalgamated crystals, such as ice cubes, or loosely accumulated granular crystals, like snow. Aside from common hexagonal crystalline ice, other crystalline and amorphous phases of water ice are known. The gaseous phase of water is known equally water vapor (or steam). Visible steam and clouds are formed from infinitesimal aerosol of h2o suspended in the air.

H2o besides forms a supercritical fluid. The critical temperature is 647 Chiliad and the critical pressure level is 22.064 MPa. In nature, this only rarely occurs in extremely hostile conditions. A likely example of naturally occurring supercritical h2o is in the hottest parts of deep water hydrothermal vents, in which water is heated to the critical temperature by volcanic plumes 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).^[27]

Heat capacity and heats of vaporization and fusion [edit]

Rut of vaporization of water from melting to critical temperature

Water has a very high specific heat capacity of 4184 J/(kg·K) at 25 °C—the second-highest amidst all the heteroatomic species (after ammonia), equally well as a high estrus of vaporization (forty.65 kJ/mol or 2257 kJ/kg at the normal humid point), both of which are a effect of the extensive hydrogen bonding betwixt its molecules. These two unusual backdrop allow water to moderate Earth's climate by buffering large fluctuations in temperature. Near of the additional energy stored in the climate system since 1970 has accumulated in the oceans.^[28]

The specific enthalpy of fusion (more usually known as latent heat) of water is 333.55 kJ/kg at 0 °C: the aforementioned 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 h2o by near eighty °C. Of common substances, only that of ammonia is college. This property confers resistance to melting on the ice of glaciers and drift water ice. Before and since the advent of mechanical refrigeration, ice was and still is in common use for retarding nutrient spoilage.

The specific heat chapters of water ice at −x °C is 2030 J/(kg·Grand)^[29] and the heat capacity of steam at 100 °C is 2080 J/(kg·M).^[30]

Density of h2o and water ice [edit]

Density of ice and water as a role of temperature

The density of water is about i gram per cubic centimetre (62 lb/cu ft): this relationship was originally used to ascertain the gram.^[31] The density varies with temperature, but not linearly: every bit the temperature increases, the density rises to a peak at 3.98 °C (39.16 °F) and and so decreases;^[32] this is unusual.^[d] Regular, hexagonal ice is also less dumbo than liquid water—upon freezing, the density of water decreases past almost 9%.^{[35] [due east]}

These peculiar effects are due to the highly directional bonding of h2o molecules via the hydrogen bonds: at depression temperature a comparatively depression-density, low-free energy structure results, while the breaking of hydrogen bonds with increasing temperature in the range 0–4 °C allows for a denser molecular packing.^[32] H2o about the humid point is about 4% less dense than water at 4 °C (39 °F).^{[35] [f]}

Under increasing pressure, water ice undergoes a number of transitions to other polymorphs with higher density than liquid water, such as ice II, ice III, high-density amorphous water ice (HDA), and very-high-density amorphous ice (VHDA).^{[36] [37]}

Temperature distribution in a lake in summertime and winter

The unusual density curve and lower density of ice than of water is essential for much of the life on world—if h2o were most dumbo at the freezing signal, and 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 lesser upwards, and all life in it would be killed.^[35] Furthermore, given that water is a practiced thermal insulator (due to its heat capacity), some frozen lakes might not completely thaw in summer.^[35] As information technology is, the inversion of the density bend leads to a stable layering for surface temperatures below 4 °C, and with the layer of ice that floats on superlative insulating the water below,^[38] even e.g., Lake Baikal in key Siberia freezes but to almost 1 m thickness in winter. In full general, for deep enough lakes, the temperature at the bottom stays abiding at most 4 °C (39 °F) throughout the year (come across diagram).^[35]

Density of saltwater and ice [edit]

The density of saltwater depends on the dissolved salt content equally well as the temperature. Water ice still floats in the oceans, otherwise, they would freeze from the bottom upward. However, the salt content of oceans lowers the freezing point past about ane.9 °C^[39] (run across here for explanation) and lowers the temperature of the density maximum of water to the sometime freezing point at 0 °C. This is why, in bounding main water, the downwardly convection of colder water is not blocked past an expansion of water as it becomes colder most the freezing point. The oceans' cold water near the freezing point continues to sink. And then creatures that alive at the bottom of cold oceans like the Arctic Ocean by and large live in water four °C colder than at the bottom of frozen-over fresh h2o lakes and rivers.

As the surface of saltwater begins to freeze (at −1.9 °C^[39] for normal salinity seawater, iii.5%) the ice that forms is substantially salt-complimentary, with nigh the aforementioned density as freshwater ice. This ice floats on the surface, and the salt that is "frozen out" adds to the salinity and density of the seawater but below it, in a process known as brine rejection. This denser saltwater sinks by convection and the replacing seawater is subject to the same procedure. This produces substantially freshwater water ice at −1.9 °C^[39] 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 procedure 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 organisation of currents called the thermohaline circulation.

Miscibility and condensation [edit]

Red line shows saturation

H2o is miscible with many liquids, including ethanol in all proportions. Water and near oils are immiscible normally forming layers according to increasing density from the top. This can be predicted past comparing the polarity. Water beingness a relatively polar compound will tend to exist miscible with liquids of high polarity such as ethanol and acetone, whereas compounds with low polarity volition tend to be immiscible and poorly soluble such as with hydrocarbons.

Every bit a gas, water vapor is completely miscible with air. On the other hand, the maximum water vapor 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's partial force per unit area is ii% of atmospheric pressure and the air is cooled from 25 °C, starting at about 22 °C, water will start to condense, defining the dew point, and creating fog or dew. The opposite process accounts for the fog burning off in the morn. 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 stage change then condenses out every bit minute water aerosol, unremarkably referred to equally steam.

A saturated gas or i with 100% relative humidity is when the vapor force per unit area of water in the air is at equilibrium with vapor pressure due to (liquid) water; water (or ice, if cool plenty) will fail to lose mass through evaporation when exposed to saturated air. Considering 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 force per unit area, is much more useful. Vapor force per unit area in a higher place 100% relative humidity is called supersaturated and can occur if the air is rapidly cooled, for case, by ascent suddenly in an updraft.^[g]

Vapor pressure [edit]

Vapor pressure diagrams of water

Compressibility [edit]

The compressibility of water is a function of force per unit area and temperature. At 0 °C, at the limit of zero pressure, the compressibility is 5.1×10^−x Pa⁻¹ . At the zero-pressure level limit, the compressibility reaches a minimum of 4.iv×10⁻¹⁰ Pa^−one around 45 °C before increasing again with increasing temperature. As the pressure is increased, the compressibility decreases, being 3.9×ten^−ten Pa⁻¹ at 0 °C and 100 megapascals (1,000 bar).^[40]

The bulk modulus of water is most ii.2 GPa.^[41] The low compressibility of non-gasses, and of h2o in particular, leads to their often being assumed as incompressible. The depression compressibility of water ways that even in the deep oceans at four km depth, where pressures are xl MPa, there is only a 1.eight% decrease in volume.^[41]

The majority modulus of water ice ranges from xi.three GPa at 0 K up to eight.6 GPa at 273 K.^[42] The large change in the compressibility of ice as a function of temperature is the consequence of its relatively big thermal expansion coefficient compared to other common solids.

Triple point [edit]

The solid/liquid/vapour triple point of liquid water, water ice I_h and water vapor in the lower left portion of a water phase diagram.

The temperature and pressure at which ordinary solid, liquid, and gaseous water coexist in equilibrium is a triple point of h2o. Since 1954, this bespeak had been used to define the base unit of temperature, the kelvin,^{[43] [44]} merely, starting in 2019, the kelvin is at present defined using the Boltzmann constant, rather than the triple signal of water.^[45]

Due to the existence of many polymorphs (forms) of ice, water has other triple points, which take either three polymorphs of water ice or two polymorphs of water ice and liquid in equilibrium.^[44] 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.^{[46] [47] [48]}

The various triple points of water

Phases in stable equilibrium	Pressure	Temperature
liquid water, water ice Ih, and water vapor	611.657 Pa[49]	273.16 K (0.01 °C)
liquid water, ice Ih, and water ice Iii	209.ix MPa	251 K (−22 °C)
liquid h2o, ice 3, and ice 5	350.1 MPa	−17.0 °C
liquid water, water ice V, and ice Half dozen	632.iv MPa	0.sixteen °C
water ice Ih, Ice II, and ice III	213 MPa	−35 °C
ice II, ice III, and ice V	344 MPa	−24 °C
water ice II, ice 5, and ice Half dozen	626 MPa	−70 °C

Melting indicate [edit]

The melting point of ice is 0 °C (32 °F; 273 K) at standard pressure; however, pure liquid water can be supercooled well below that temperature without freezing if the liquid is not mechanically disturbed. It tin remain in a fluid state down to its homogeneous nucleation point of almost 231 M (−42 °C; −44 °F).^[50] The melting bespeak of ordinary hexagonal ice falls slightly nether moderately high pressures, by 0.0073 °C (0.0131 °F)/atm^[h] or about 0.five °C (0.90 °F)/70 atm^{[i] [51]} as the stabilization free energy of hydrogen bonding is exceeded by intermolecular repulsion, but every bit ice transforms into its polymorphs (come across crystalline states of ice) above 209.9 MPa (2,072 atm), the melting betoken increases markedly with pressure, i.e., reaching 355 1000 (82 °C) at 2.216 GPa (21,870 atm) (triple bespeak of Ice Vii^[52]).

Electrical properties [edit]

Conductivity [edit]

Pure h2o containing no exogenous ions is an first-class electronic insulator, merely non fifty-fifty "deionized" water is completely free of ions. Water undergoes autoionization in the liquid state when ii water molecules form one hydroxide anion (OH ⁻
) and one hydronium cation (H
₃ O ⁺
). Because of autoionization, at ambient temperatures pure liquid h2o has a like 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 accuse 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.^[53]

Because h2o is such a good solvent, information technology well-nigh always has some solute dissolved in it, often a salt. If h2o has even a tiny corporeality of such an impurity, so the ions can carry charges dorsum and forth, allowing the water to conduct electricity far more readily.

It is known that the theoretical maximum electric resistivity for h2o is approximately 18.2 MΩ·cm (182 kΩ·m) at 25 °C.^[54] This effigy agrees well with what is typically seen on reverse osmosis, ultra-filtered and deionized ultra-pure water systems used, for instance, in semiconductor manufacturing plants. A salt or acrid contaminant level exceeding even 100 parts per trillion (ppt) in otherwise ultra-pure water begins to noticeably lower its resistivity by upward to several kΩ·chiliad.^{[ commendation needed ]}

In pure water, sensitive equipment can observe a very slight electrical conductivity of 0.05501 ± 0.0001 μS/cm at 25.00 °C.^[54] Water can too be electrolyzed into oxygen and hydrogen gases but in the absenteeism of dissolved ions this is a very ho-hum procedure, equally very little current is conducted. In ice, the primary charge carriers are protons (meet proton conductor).^[55] Water ice was previously thought to have a small but measurable conductivity of 1×10 ⁻¹⁰  S/cm, simply this conductivity is now thought to be virtually entirely from surface defects, and without those, ice is an insulator with an immeasurably small conductivity.^[32]

Polarity and hydrogen bonding [edit]

A diagram showing the partial charges on the atoms in a h2o molecule

An important feature of h2o is its polar nature. The structure has a aptitude molecular geometry for the ii hydrogens from the oxygen vertex. The oxygen atom likewise has ii lone pairs of electrons. One result unremarkably ascribed to the lone pairs is that the H–O–H gas-phase bend angle is 104.48°,^[56] which is smaller than the typical tetrahedral angle of 109.47°. The solitary pairs are closer to the oxygen atom than the electrons sigma bonded to the hydrogens, so they require more space. The increased repulsion of the solitary pairs forces the O–H bonds closer to each other. ^[57]

Some other issue of its construction is that h2o is a polar molecule. Due to the difference in electronegativity, a bond dipole moment points from each H to the O, making the oxygen partially negative and each hydrogen partially positive. A big molecular dipole, points from a region between the two hydrogen atoms to the oxygen atom. The charge differences crusade h2o molecules to aggregate (the relatively positive areas being attracted to the relatively negative areas). This allure, hydrogen bonding, explains many of the properties of water, such every bit its solvent properties.^[58]

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 h2o's physical properties. These backdrop include its relatively high melting and boiling point temperatures: more free energy is required to suspension the hydrogen bonds between water molecules. In dissimilarity, hydrogen sulfide (H
_two S), has much weaker hydrogen bonding due to sulfur's lower electronegativity. H
₂ S is a gas at room temperature, despite hydrogen sulfide having nigh twice the molar mass of water. The extra bonding between water molecules also gives liquid h2o a big specific estrus capacity. This high heat capacity makes water a good heat storage medium (coolant) and heat shield.

Cohesion and adhesion [edit]

Water molecules stay close to each other (cohesion), due to the commonage activity of hydrogen bonds between water molecules. These hydrogen bonds are constantly breaking, with new bonds being formed with dissimilar 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.^[59]

H2o besides has loftier adhesion properties because of its polar nature. On clean, shine drinking glass the water may course a thin film because the molecular forces betwixt glass and h2o molecules (adhesive forces) are stronger than the cohesive forces.^{[ citation needed ]} In biological cells and organelles, water is in contact with membrane and protein surfaces that are hydrophilic; that is, surfaces that have a strong attraction to water. Irving Langmuir observed a strong repulsive forcefulness between hydrophilic surfaces. To dehydrate hydrophilic surfaces—to remove the strongly held layers of h2o of hydration—requires doing substantial work against these forces, chosen hydration forces. These forces are very large just decrease speedily over a nanometer or less.^[threescore] They are of import in biology, peculiarly when cells are dehydrated by exposure to dry atmospheres or to extracellular freezing.^[61]

Surface tension [edit]

This paper prune is nether the water level, which has risen gently and smoothly. Surface tension prevents the clip from submerging and the water from flood the glass edges.

Temperature dependence of the surface tension of pure water

Water has an unusually loftier surface tension of 71.99 mN/m at 25 °C^[62] which is caused by the strength of the hydrogen bonding between h2o molecules.^[63] This allows insects to walk on water.^[63]

Capillary action [edit]

Because water has strong cohesive and adhesive forces, it exhibits capillary action.^[64] Strong cohesion from hydrogen bonding and adhesion allows copse to transport water more than than 100 m upward.^[63]

H2o every bit a solvent [edit]

Water is an excellent solvent due to its high dielectric constant.^[65] Substances that mix well and dissolve in water are known as hydrophilic ("water-loving") substances, while those that do not mix well with water are known as hydrophobic ("water-fearing") substances.^[66] The ability of a substance to dissolve in water is determined by whether or not the substance can friction match or meliorate the potent attractive forces that water molecules generate between other water molecules. If a substance has properties that practice non allow information technology to overcome these strong intermolecular forces, the molecules are precipitated out from the water. Contrary to the common misconception, water and hydrophobic substances exercise not "repel", and the hydration of a hydrophobic surface is energetically, only non entropically, favorable.

When an ionic or polar compound enters water, it is surrounded past h2o molecules (hydration). The relatively small size of water molecules (~ 3 angstroms) allows many water molecules to surround one molecule of solute. 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 full general, ionic and polar substances such every bit acids, alcohols, and salts are relatively soluble in water, and nonpolar substances such equally fats and oils are not. Nonpolar molecules stay together in water considering it is energetically more favorable for the water molecules to hydrogen bond to each other than to appoint in van der Waals interactions with non-polar molecules.

An example of an ionic solute is table table salt; the sodium chloride, NaCl, separates into Na ⁺
cations and Cl ⁻
anions, each being surrounded past h2o molecules. The ions are then hands transported away from their crystalline lattice into solution. An example of a nonionic solute is table sugar. The water dipoles brand hydrogen bonds with the polar regions of the sugar molecule (OH groups) and let it to be carried away into solution.

Breakthrough tunneling [edit]

The quantum tunneling dynamics in water was reported equally early as 1992. At that time it was known that at that place are motions which destroy and regenerate the weak hydrogen bond by internal rotations of the substituent water monomers.^[67] On eighteen March 2016, information technology was reported that the hydrogen bond can exist broken by breakthrough tunneling in the water hexamer. Unlike previously reported tunneling motions in water, this involved the concerted breaking of ii hydrogen bonds.^[68] Later in the same year, the discovery of the quantum tunneling of water molecules was reported.^[69]

Electromagnetic absorption [edit]

Water is relatively transparent to visible light, well-nigh ultraviolet low-cal, and far-cherry-red lite, simply information technology absorbs most ultraviolet calorie-free, infrared light, and microwaves. Near photoreceptors and photosynthetic pigments apply the portion of the light spectrum that is transmitted well through water. Microwave ovens accept advantage of h2o's opacity to microwave radiations to oestrus the water inside of foods. Water's light bluish color is caused by weak absorption in the red part of the visible spectrum.^{[iii] [70]}

Structure [edit]

A single water molecule can participate in a maximum of four hydrogen bonds because it can accept ii bonds using the lone pairs on oxygen and donate two hydrogen atoms. Other molecules like hydrogen fluoride, ammonia, and methanol tin as well grade hydrogen bonds. However, they practice not evidence dissonant thermodynamic, 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 at that place are steric effects in beefy residues. In water, intermolecular tetrahedral structures form due to the four hydrogen bonds, thereby forming an open construction and a three-dimensional bonding network, resulting in the anomalous decrease in density when cooled below 4 °C. This repeated, constantly reorganizing unit of measurement defines a three-dimensional network extending throughout the liquid. This view is based upon neutron handful studies and figurer 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 from Stockholm University suggested that water molecules in the liquid country typically bind non to iv only but two others; thus forming chains and rings. The term "string theory of water" (which is not to be confused with the cord theory of physics) was coined. These observations were based upon X-ray absorption spectroscopy that probed the local environment of individual oxygen atoms.^[71]

Molecular structure [edit]

The repulsive effects of the ii solitary pairs on the oxygen cantlet cause water to have a aptitude, non linear, molecular structure,^[72] assuasive it to be polar. The hydrogen–oxygen–hydrogen bending is 104.45°, which is less than the 109.47° for ideal sp³ hybridization. The valence bail theory explanation is that the oxygen atom's lone pairs are physically larger and therefore accept up more space than the oxygen atom'due south bonds to the hydrogen atoms.^[73] The molecular orbital theory caption (Bent's dominion) is that lowering the energy of the oxygen atom'southward nonbonding hybrid orbitals (past assigning them more s character and less p character) and correspondingly raising the free energy of the oxygen atom's hybrid orbitals bonded to the hydrogen atoms (by assigning them more p grapheme and less due south graphic symbol) 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'southward alone pairs while the energy of the oxygen atom's other two hybrid orbitals contributes just partially to the energy of the bonding orbitals (the remainder of the contribution coming from the hydrogen atoms' 1s orbitals).

Chemical properties [edit]

Cocky-ionization [edit]

In liquid water there is some self-ionization giving hydronium ions and hydroxide ions.

2 H
₂ O ⇌ H
₃ O ⁺
+ OH ⁻

The equilibrium abiding for this reaction, known as the ionic product of water, ${\displaystyle K_{\rm {w}}=[{\rm {H_{three}O^{+}}}][{\rm {OH^{-}}}]}$ , has a value of virtually x ⁻¹⁴ at 25 °C. At neutral pH, the concentration of the hydroxide ion (OH ⁻
) equals that of the (solvated) hydrogen ion (H ⁺
), with a value close to 10^−vii mol L⁻¹ at 25 °C.^[74] See data folio for values at other temperatures.

The thermodynamic equilibrium abiding is a caliber of thermodynamic activities of all products and reactants including water:

${\displaystyle K_{\rm {eq}}={\frac {a_{\rm {H_{3}O^{+}}}\cdot a_{\rm {OH^{-}}}}{a_{\rm {H_{2}O}}^{ii}}}}$

However for dilute solutions, the activeness of a solute such as H₃O⁺ or OH⁻ is approximated by its concentration, and the action of the solvent H₂O is approximated by 1, then that we obtain the simple ionic product ${\displaystyle K_{\rm {eq}}\approx K_{\rm {w}}=[{\rm {H_{three}O^{+}}}][{\rm {OH^{-}}}]}$

Geochemistry [edit]

The action of water on rock over long periods of time typically leads to weathering and water erosion, physical processes that convert solid rocks and minerals into soil and sediment, but under some conditions chemic reactions with water occur as well, resulting in metasomatism or mineral hydration, a type of chemical alteration of a stone which produces clay minerals. It also occurs when Portland cement hardens.

Water ice tin grade clathrate compounds, known as clathrate hydrates, with a variety of small molecules that can exist embedded in its spacious crystal lattice. The most notable of these is methane clathrate, 4 CH
₄ ·23H
₂ O, naturally found in big quantities on the ocean flooring.

Acidity in nature [edit]

Rain is generally mildly acidic, with a pH between 5.2 and 5.8 if not having whatever acid stronger than carbon dioxide.^[75] If high amounts of nitrogen and sulfur oxides are present in the air, they as well will dissolve into the cloud and raindrops, producing acid rain.

Isotopologues [edit]

Several isotopes of both hydrogen and oxygen exist, giving rise to several known isotopologues of water. Vienna Standard Hateful Ocean Water is the current international standard for water isotopes. Naturally occurring water is almost completely composed of the neutron-less hydrogen isotope protium. Simply 155 ppm include deuterium ( ²
H or D), a hydrogen isotope with one neutron, and fewer than twenty parts per quintillion include tritium ( ³
H or T), which has two neutrons. Oxygen also has three stable isotopes, with ¹⁶
O present in 99.76%, ¹⁷
O in 0.04%, and ^eighteen
O in 0.2% of water molecules.^[76]

Deuterium oxide, D
₂ O, is as well known as heavy water because of its higher density. It is used in nuclear reactors as a neutron moderator. Tritium is radioactive, decaying with a half-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 ane deuterium atom HDO occur naturally in ordinary water in depression concentrations (~0.03%) and D
_ii O in far lower amounts (0.000003%) and whatsoever such molecules are temporary as the atoms recombine.

The almost notable physical differences between H
_two O and D
₂ O, other than the simple difference in specific mass, involve properties that are affected by hydrogen bonding, such as freezing and humid, and other kinetic furnishings. This is considering the nucleus of deuterium is twice every bit heavy as protium, and this causes noticeable differences in bonding energies. The deviation in boiling points allows the isotopologues to be separated. The self-improvidence coefficient of H
_two O at 25 °C is 23% higher than the value of D
₂ O.^[77] Considering h2o molecules exchange hydrogen atoms with 1 another, hydrogen deuterium oxide (DOH) is much more common in low-purity heavy water than pure dideuterium monoxide D
_ii O.

Consumption of pure isolated D
₂ O may affect biochemical processes—ingestion of large amounts impairs kidney and central nervous system function. Minor quantities tin be consumed without any ill-furnishings; humans are by and large unaware of taste differences,^[78] but sometimes report a burning sensation^[79] or sweet flavor.^[fourscore] Very large amounts of heavy water must exist consumed for any toxicity to become apparent. Rats, nevertheless, are able to avert heavy h2o by smell, and it is toxic to many animals.^[81]

Light water refers to deuterium-depleted water (DDW), water in which the deuterium content has been reduced beneath the standard 155 ppm level.

Occurrence [edit]

H2o is the about abundant substance on Globe and also the third most abundant molecule in the universe, after H
_ii and CO.^[22] 0.23 ppm of the earth'southward mass is water and 97.39% of the global h2o book of 1.38×10 ⁹ km³ is institute in the oceans.^[82]

Reactions [edit]

Acid-base reactions [edit]

H2o is amphoteric: it has the ability to act every bit either an acid or a base in chemical reactions.^[83] According to the Brønsted-Lowry definition, an acid is a proton (H ⁺
) donor and a base is a proton acceptor.^[84] When reacting with a stronger acid, h2o acts every bit a base; when reacting with a stronger base, it acts as an acid.^[84] For instance, h2o receives an H ⁺
ion from HCl when muriatic acid is formed:

HCl
(acrid) + H
₂ O
(base) ⇌ H
_iii O ⁺
+ Cl ⁻

In the reaction with ammonia, NH
₃ , water donates a H ⁺
ion, and is thus acting as an acid:

NH
₃
(base) + H
₂ O
(acid) ⇌ NH ⁺
₄ + OH ⁻

Considering the oxygen cantlet in h2o has two lone pairs, water often acts as a Lewis base of operations, or electron-pair donor, in reactions with Lewis acids, although it can too 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 of operations, meaning that information technology reacts preferentially with other hard species:

H ⁺

(Lewis acrid) + H
₂ O
(Lewis base of operations) → H
₃ O ⁺

Fe ^three+

(Lewis acid) + H
_two O
(Lewis base) → Fe(H
₂ O) ³⁺
₆

Cl ⁻

(Lewis base of operations) + H
_ii O
(Lewis acrid) → Cl(H
₂ O) ⁻
_{half dozen}

When a table salt of a weak acrid or of a weak base of operations is dissolved in water, water tin can partially hydrolyze the salt, producing the corresponding base or acid, which gives aqueous solutions of lather and baking soda their basic pH:

Na
₂ CO
₃ + H
₂ O ⇌ NaOH + NaHCO
₃

Ligand chemistry [edit]

Water'due south Lewis base character makes it a mutual ligand in transition metal complexes, examples of which include metal aquo complexes such equally Fe(H
₂ O) ²⁺
_half-dozen to perrhenic acrid, which contains 2 h2o molecules coordinated to a rhenium center. In solid hydrates, h2o can be either a ligand or merely lodged in the framework, or both. Thus, FeSO
₄ ·7H
_two O consists of [Fe_ii(H₂O)₆]²⁺ centers and ane "lattice water". H2o is typically a monodentate ligand, i.e., information technology forms only one bail with the central atom.^[85]

Some hydrogen-bonding contacts in FeSO₄ ^.7H₂O. This metal aquo complex crystallizes with one molecule of "lattice" water, which interacts with the sulfate and with the [Fe(H_iiO)_vi]²⁺ centers.

Organic chemistry [edit]

Every bit a difficult base, water reacts readily with organic carbocations; for example in a hydration reaction, a hydroxyl grouping (OH ⁻
) and an acidic proton are added to the two carbon atoms bonded together in the carbon-carbon double bond, resulting in an booze. When the addition of water to an organic molecule cleaves the molecule in two, hydrolysis is said to occur. Notable examples of hydrolysis are the saponification of fats and the digestion of proteins and polysaccharides. H2o tin also be a leaving grouping in S_Ntwo commutation and E2 elimination reactions; the latter is then known as a dehydration reaction.

H2o in redox reactions [edit]

Water contains hydrogen in the oxidation land +1 and oxygen in the oxidation country −2.^[86] It oxidizes chemicals such as hydrides, alkali metals, and some alkaline world metals.^{[87] [88]} 1 case of an brine metallic reacting with water is:^[89]

2 Na + two H
₂ O → H
₂ + ii Na ⁺
+ 2 OH ⁻

Some other reactive metals, such as aluminum and beryllium, are oxidized by water besides, but their oxides attach to the metal and course a passive protective layer.^[90] Note that the rusting of iron is a reaction betwixt fe and oxygen^[91] that is dissolved in water, not between iron and water.

Water can exist oxidized to emit oxygen gas, but very few oxidants react with h2o even if their reduction potential is greater than the potential of O
₂ /H
₂ O. Almost all such reactions require a goad.^[92] An example of the oxidation of h2o is:

4 AgF
₂ + 2 H
₂ O → iv AgF + 4 HF + O
₂

Electrolysis [edit]

H2o can be split into its constituent elements, hydrogen, and oxygen, past passing an electric current through it.^[93] This process is chosen electrolysis. The cathode half reaction is:

2 H ⁺
+ 2
e ⁻
→ H
_ii

The anode one-half reaction is:

2 H
_two O → O
₂ + iv H ⁺
+ 4
eastward ⁻

The gases produced bubble to the surface, where they can exist collected or ignited with a flame higher up the water if this was the intention. The required potential for the electrolysis of pure h2o is 1.23 V at 25 °C.^[93] The operating potential is actually 1.48 V or higher in practical electrolysis.

History [edit]

Henry Cavendish showed that h2o was composed of oxygen and hydrogen in 1781.^[94] The first decomposition of water into hydrogen and oxygen, past electrolysis, was done in 1800 by English chemist William Nicholson and Anthony Carlisle.^{[94] [95]} In 1805, Joseph Louis Gay-Lussac and Alexander von Humboldt showed that h2o is composed of two parts hydrogen and one part oxygen.^[96]

Gilbert Newton Lewis isolated the starting time sample of pure heavy water in 1933.^[97]

The properties of water take historically been used to define diverse temperature scales. Notably, the Kelvin, Celsius, Rankine, and Fahrenheit scales were, or currently are, defined past the freezing and boiling points of water. The less common scales of Delisle, Newton, Réaumur, and Rømer were divers similarly. The triple point of water is a more commonly used standard point today.

Nomenclature [edit]

The accustomed IUPAC proper name of h2o is oxidane or simply h2o,^[98] or its equivalent in different languages, although in that location are other systematic names which can exist used to describe the molecule. Oxidane is only intended to be used as the proper name of the mononuclear parent hydride used for naming derivatives of water by substituent nomenclature.^[99] These derivatives normally have other recommended names. For example, the name hydroxyl is recommended over oxidanyl for the –OH group. The proper name oxane is explicitly mentioned by the IUPAC equally beingness unsuitable for this purpose, since it is already the name of a cyclic ether as well known as tetrahydropyran.^{[100] [101]}

The simplest systematic name of h2o is hydrogen oxide. This is analogous to related compounds such as hydrogen peroxide, hydrogen sulfide, and deuterium oxide (heavy water). Using chemical classification for type I ionic binary compounds, water would have the name hydrogen monoxide,^[102] only this is not among the names published by the International Union of Pure and Applied Chemistry (IUPAC).^[98] Another name is dihydrogen monoxide, which is a rarely used name of water, and mostly used in the dihydrogen monoxide parody.

Other systematic names for water include hydroxic acrid, hydroxylic acid, and hydrogen hydroxide, using acrid and base names.^[j] None of these exotic names are used widely. The polarized course of the water molecule, H ⁺
OH ⁻
, is too called hydron hydroxide by IUPAC nomenclature.^[103]

Water substance is a term used for hydrogen oxide (H₂O) when one does not wish to specify whether 1 is speaking of liquid water, steam, some form of water ice, or a component in a mixture or mineral.

Run into too [edit]

• Chemical bonding of water
• Dihydrogen monoxide parody
• Double distilled water
• Electromagnetic absorption by water
• Fluid dynamics
• Difficult h2o
• Heavy water
• Hydrogen polyoxide
• Ice
• Optical properties of water and ice
• Steam
• Superheated water
• Viscosity § Water
• Water cluster
• Water (data folio)
• Water dimer
• Water model
• Water thread experiment

Footnotes [edit]

1. ^ A commonly quoted value of 15.vii used mainly in organic chemical science for the pK_a of water is wrong.^{[10] [eleven]}
2. ^ ^{a b} Vienna Standard Hateful Ocean Water (VSMOW), used for calibration, melts at 273.1500089(10) Chiliad (0.000089(10) °C, and boils at 373.1339 1000 (99.9839 °C). Other isotopic compositions melt or eddy at slightly different temperatures.
3. ^ ^{a b} H+ represents H
   ₃ O ⁺
   (H
   _two O)
   _{due north} and more circuitous ions that form.
4. ^ Negative thermal expansion is also observed in molten silica.^[33] Also, fairly pure silicon has a negative coefficient of thermal expansion for temperatures betwixt almost 18 and 120 kelvins.^[34]
5. ^ Other substances that expand on freezing are silicon (melting bespeak of 1,687 Grand (one,414 °C; ii,577 °F)), gallium (melting point of 303 K (30 °C; 86 °F), germanium (melting point of i,211 Chiliad (938 °C; i,720 °F)), and bismuth (melting indicate of 545 K (272 °C; 521 °F))
6. ^ (one-0.95865/i.00000) × 100% = iv.135%
7. ^ Adiabatic cooling resulting from the platonic gas constabulary.
8. ^ The source gives it as 0.0072°C/atm. However the writer defines an atmosphere equally one,000,000 dynes/cm² (a bar). Using the standard definition of atmosphere, i,013,250 dynes/cm², it works out to 0.0073°C/atm.
9. ^ Using the fact that 0.five/0.0073 = 68.5.
10. ^ Both acrid and base names exist for water because it is amphoteric (able to react both as an acid or an brine).

References [edit]

Notes [edit]

1. ^ "naming molecular compounds". world wide web.iun.edu. Archived from the original on 24 September 2018. Retrieved one October 2018. Sometimes these compounds have generic or common names (e.grand., H2O is "h2o") and they likewise have systematic names (e.g., H2o, dihydrogen monoxide).
2. ^ "Definition of Hydrol". Merriam-Webster . Retrieved 21 April 2019.
3. ^ ^{a b c} Braun, Charles L.; Smirnov, Sergei N. (1993-08-01). "Why is water blue?" (PDF). Journal of Chemical Instruction. 70 (viii): 612. Bibcode:1993JChEd..70..612B. doi:ten.1021/ed070p612. ISSN 0021-9584.
4. ^ Riddick 1970, Table of Concrete Backdrop, Water 0b. pg 67-8.
5. ^ Lide 2003, Backdrop of Ice and Supercooled H2o in Section half-dozen.
6. ^ ^{a b c} Anatolievich, Kiper Ruslan. "Properties of substance: water".
7. ^ Lide 2003, Vapor Pressure of Water From 0 to 370° C in Sec. 6.
8. ^ Lide 2003, Chapter 8: Dissociation Constants of Inorganic Acids and Bases.
9. ^ Weingärtner et al. 2016, p. thirteen.
10. ^ "What is the pKa of H2o". University of California, Davis. 2015-08-09.
11. ^ Silverstein, Todd P.; Heller, Stephen T. (17 April 2017). "pKa Values in the Undergraduate Curriculum: What Is the Existent pKa of H2o?". Journal of Chemic Education. 94 (vi): 690–695. Bibcode:2017JChEd..94..690S. doi:10.1021/acs.jchemed.6b00623.
12. ^ Ramires, Maria L. Five.; Castro, Carlos A. Nieto de; Nagasaka, Yuchi; Nagashima, Akira; Assael, Marc J.; Wakeham, William A. (1995-05-01). "Standard Reference Information for the Thermal Conductivity of Water". Journal of Physical and Chemical Reference Data. 24 (three): 1377–1381. Bibcode:1995JPCRD..24.1377R. doi:10.1063/1.555963. ISSN 0047-2689.
13. ^ Lide 2003, 8—Concentrative Properties of Aqueous Solutions: Density, Refractive Index, Freezing Point Depression, and Viscosity.
14. ^ Lide 2003, 6.186.
15. ^ H2o in Linstrom, Peter J.; Mallard, William G. (eds.); NIST Chemical science WebBook, NIST Standard Reference Database Number 69, National Institute of Standards and Technology, Gaithersburg (MD), http://webbook.nist.gov
16. ^ Lide 2003, 9—Dipole Moments.
17. ^ ^{a b c} H2o in Linstrom, Peter J.; Mallard, William G. (eds.); NIST Chemical science WebBook, NIST Standard Reference Database Number 69, National Institute of Standards and Technology, Gaithersburg (MD), http://webbook.nist.gov
18. ^ GHS: PubChem 962
19. ^ Greenwood & Earnshaw 1997, p. 620.
20. ^ "Water, the Universal Solvent". U.S. Department of the Interior. usgs.gov (website). U.s. of America: USGS. October 22, 2019. Retrieved December fifteen, 2020.
21. ^ Reece et al. 2013, p. 48.
22. ^ ^{a b c} Weingärtner et al. 2016, p. 2.
23. ^ Reece et al. 2013, p. 44.
24. ^ "Autoprotolysis abiding". IUPAC Compendium of Chemical Terminology. IUPAC. 2009. doi:10.1351/goldbook.A00532. ISBN978-0-9678550-9-7.
25. ^ Campbell, Williamson & Heyden 2006.
26. ^ Smith, Jared D.; Christopher D. Cappa; Kevin R. Wilson; Ronald C. Cohen; Phillip L. Geissler; Richard J. Saykally (2005). "Unified clarification of temperature-dependent hydrogen bond rearrangements in liquid h2o" (PDF). Proc. Natl. Acad. Sci. USA. 102 (40): 14171–14174. Bibcode:2005PNAS..10214171S. doi:10.1073/pnas.0506899102. PMC1242322. PMID 16179387.
27. ^ Deguchi, Shigeru; Tsujii, Kaoru (2007-06-nineteen). "Supercritical water: a fascinating medium for soft matter". Soft Matter. iii (vii): 797–803. Bibcode:2007SMat....3..797D. doi:x.1039/b611584e. ISSN 1744-6848. PMID 32900070.
28. ^ Rhein, M.; Rintoul, South.R. (2013). "3: Observations: Ocean" (PDF). IPCC WGI AR5 (Report). p. 257. Ocean warming dominates the global energy alter inventory. Warming of the ocean accounts for nigh 93% of the increase in the World'due south energy inventory between 1971 and 2010 (loftier confidence), with the warming of the upper (0 to 700 m) sea accounting for about 64% of the total. Melting ice (including Arctic ocean water ice, ice sheets, and glaciers) and warming of the continents and atmosphere account for the remainder of the change in energy.
29. ^ Lide 2003, Chapter 6: Properties of Water ice and Supercooled Water.
30. ^ Lide 2003, 6. Properties of Water and Steam as a Function of Temperature and Pressure.
31. ^ "Prescript on weights and measures". Apr 7, 1795. Archived from the original on February 25, 2013. Retrieved August 3, 2016. 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.
32. ^ ^{a b c} Greenwood & Earnshaw 1997, p. 625.
33. ^ Vanquish, Scott Grand.; Debenedetti, Pablo Grand.; Panagiotopoulos, Athanassios Z. (2002). "Molecular structural order and anomalies in liquid silica" (PDF). Phys. Rev. E. 66 (1): 011202. arXiv:cond-mat/0203383. Bibcode:2002PhRvE..66a1202S. doi:10.1103/PhysRevE.66.011202. PMID 12241346. S2CID 6109212. Archived from the original (PDF) on 2016-06-04. Retrieved 2009-07-07 .
34. ^ Bullis, W. Murray (1990). "Chapter 6". In O'Mara, William C.; Herring, Robert B.; Hunt, Lee P. (eds.). Handbook of semiconductor silicon technology. Park Ridge, New Jersey: Noyes Publications. p. 431. ISBN0-8155-1237-6 . Retrieved 2010-07-xi .
35. ^ ^{a b c d e} Perlman, Howard. "H2o Density". The USGS H2o Science School . Retrieved 2016-06-03 .
36. ^ Loerting, Thomas; Salzmann, Christoph; Kohl, Ingrid; Mayer, Erwin; Hallbrucker, Andreas (2001-01-01). "A 2d distinct structural "state" of high-density amorphous water ice at 77 M and 1 bar". Physical Chemical science Chemic Physics. three (24): 5355–5357. Bibcode:2001PCCP....three.5355L. doi:10.1039/b108676f. ISSN 1463-9084.
37. ^ Greenwood & Earnshaw 1997, p. 624.
38. ^ Zumdahl & Zumdahl 2013, p. 493.
39. ^ ^{a b c} "Can the ocean freeze?". National Body of water Service. National Oceanic and Atmospheric Administration. Retrieved 2016-06-09 .
40. ^ Fine, R.A.; Millero, F.J. (1973). "Compressibility of water as a part of temperature and pressure". Journal of Chemical Physics. 59 (10): 5529. Bibcode:1973JChPh..59.5529F. doi:x.1063/1.1679903.
41. ^ ^{a b} Nave, R. "Majority Rubberband Properties". HyperPhysics. Georgia State University. Retrieved 2007-x-26 .
42. ^ Neumeier, J.J. (2018). "Elastic Constants, Bulk Modulus, and Compressibility of H₂O Ice Ih for the Temperature Range 50 K-273 G". Journal of Physical and Chemical Reference Data. 47 (3): 033101. Bibcode:2018JPCRD..47c3101N. doi:10.1063/1.5030640. S2CID 105357042.
43. ^ "Base unit definitions: Kelvin". National Plant of Standards and Applied science. Retrieved 9 August 2018.
44. ^ ^{a b} Weingärtner et al. 2016, p. 5.
45. ^ Proceedings of the 106th meeting (PDF). International Committee for Weights and Measures. Sèvres. 16–20 October 2017.
46. ^ Schlüter, Oliver (2003-07-28). "Impact of Loftier Pressure level — Low Temperature Processes on Cellular Materials Related to Foods" (PDF). Technischen Universität Berlin. Archived from the original (PDF) on 2008-03-09.
47. ^ Tammann, Gustav H.J.A (1925). "The States Of Aggregation". Constable And Company.
48. ^ Lewis & Rice 1922.
49. ^ White potato, D. M. (2005). "Review of the vapour pressures of ice and supercooled water for atmospheric applications". Quarterly Journal of the Royal Meteorological Society. 131 (608): 1539–1565. Bibcode:2005QJRMS.131.1539M. doi:x.1256/qj.04.94.
50. ^ Debenedetti, P. G.; Stanley, H. Due east. (2003). "Supercooled and Glassy Water" (PDF). Physics Today. 56 (6): xl–46. Bibcode:2003PhT....56f..40D. doi:10.1063/1.1595053.
51. ^ Sharp 1988, p. 27.
52. ^ "Revised Release on the Pressure along the Melting and Sublimation Curves of Ordinary Water Substance" (PDF). IAPWS. September 2011. Retrieved 2013-02-nineteen .
53. ^ C. Southward. Fuller "Defect Interactions in Semiconductors" Affiliate 5 pp. 192-221 in "Semiconductors" North. B. Hannay Ed. Reinhold, New York 1959
54. ^ ^{a b} Light, Truman Due south.; Licht, Stuart; Bevilacqua, Anthony C.; Morash, Kenneth R. (2005-01-01). "The Fundamental Electrical conductivity and Resistivity of Water". Electrochemical and Solid-State Letters. 8 (ane): E16–E19. doi:10.1149/1.1836121. ISSN 1099-0062.
55. ^ Crofts, A. (1996). "Lecture 12: Proton Conduction, Stoichiometry". University of Illinois at Urbana-Champaign. Retrieved 2009-12-06 .
56. ^ Hoy, AR; Bunker, PR (1979). "A precise solution of the rotation bending Schrödinger equation for a triatomic molecule with application to the water molecule". Journal of Molecular Spectroscopy. 74 (one): i–eight. Bibcode:1979JMoSp..74....1H. doi:10.1016/0022-2852(79)90019-5.
57. ^ Zumdahl & Zumdahl 2013, p. 393.
58. ^ Campbell & Farrell 2007, pp. 37–38.
59. ^ Campbell & Reece 2009, p. 47.
60. ^ Chiavazzo, Eliodoro; Fasano, Matteo; Asinari, Pietro; Decuzzi, Paolo (2014). "Scaling behaviour for the water transport in nanoconfined geometries". Nature Communications. 5: 4565. Bibcode:2014NatCo...5.4565C. doi:10.1038/ncomms4565. PMC3988813. PMID 24699509.
61. ^ "Physical Forces Organizing Biomolecules" (PDF). Biophysical Social club. Archived from the original on August 7, 2007. {{cite web}}: CS1 maint: unfit URL (link)
62. ^ Lide 2003, Surface Tension of Mutual Liquids.
63. ^ ^{a b c} Reece et al. 2013, p. 46.
64. ^ Zumdahl & Zumdahl 2013, pp. 458–459.
65. ^ Greenwood & Earnshaw 1997, p. 627.
66. ^ Zumdahl & Zumdahl 2013, p. 518.
67. ^ Pugliano, N. (1992-11-01). "Vibration-Rotation-Tunneling Dynamics in Small H2o Clusters". Lawrence Berkeley Lab., CA (The states): six. doi:x.2172/6642535. OSTI 6642535.
68. ^ Richardson, Jeremy O.; Pérez, Cristóbal; Lobsiger, Simon; Reid, Adam A.; Temelso, Berhane; Shields, George C.; Kisiel, Zbigniew; Wales, David J.; Pate, Brooks H.; Althorpe, Stuart C. (2016-03-18). "Concerted hydrogen-bond breaking by quantum tunneling in the water hexamer prism". Scientific discipline. 351 (6279): 1310–1313. Bibcode:2016Sci...351.1310R. doi:10.1126/science.aae0012. ISSN 0036-8075. PMID 26989250.
69. ^ Kolesnikov, Alexander I. (2016-04-22). "Quantum Tunneling of Water in Beryl: A New State of the H2o Molecule". Physical Review Letters. 116 (16): 167802. Bibcode:2016PhRvL.116p7802K. doi:10.1103/PhysRevLett.116.167802. PMID 27152824.
70. ^ Pope; Fry (1996). "Assimilation spectrum (380-700nm) of pure h2o. II. Integrating cavity measurements". Applied Eyes. 36 (33): 8710–23. Bibcode:1997ApOpt..36.8710P. doi:ten.1364/ao.36.008710. PMID 18264420.
71. ^ Brawl, Philip (2008). "H2o—an enduring mystery". Nature. 452 (7185): 291–292. Bibcode:2008Natur.452..291B. doi:10.1038/452291a. PMID 18354466. S2CID 4365814.
72. ^ Gonick, Larry; Criddle, Craig (2005-05-03). "Chapter 3 Togetherness". The cartoon guide to chemistry (1st ed.). HarperResource. p. 59. ISBN9780060936778. Water, H_iiO, is similar. Information technology has ii electron pairs with goose egg fastened to them. They, as well, must be taken into account. Molecules like NH₃ and H₂O are called bent.
73. ^ Theodore L. Brown; et al. (2015). "9.2 The Vsepr Model". Chemical science : the central scientific discipline (13 ed.). p. 351. ISBN978-0-321-91041-7 . Retrieved 21 April 2019. Notice that the bond angles decrease equally 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 i 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 (Effigy 9.seven). Nonbonding electron pairs, therefore, take up more than space than bonding pairs; in essence, they human action as big and fatter balloons in our analogy of Figure nine.5. As a result, electron domains for nonbonding electron pairs exert greater repulsive forces on adjacent electron domains and tend to shrink bail angles
74. ^ Boyd 2000, p. 105.
75. ^ Boyd 2000, p. 106.
76. ^ "Guideline on the Use of Cardinal Physical Constants and Basic Constants of Water" (PDF). IAPWS. 2001.
77. ^ Hardy, Edme H.; Zygar, Astrid; Zeidler, Manfred D.; Holz, Manfred; Sacher, Frank D. (2001). "Isotope issue on the translational and rotational motion in liquid water and ammonia". J. Chem. Phys. 114 (7): 3174–3181. Bibcode:2001JChPh.114.3174H. doi:10.1063/1.1340584.
78. ^ Urey, Harold C.; et al. (15 Mar 1935). "Apropos the Taste of Heavy Water". Science. Vol. 81, no. 2098. New York: The Scientific discipline Printing. p. 273. Bibcode:1935Sci....81..273U. doi:ten.1126/science.81.2098.273-a.
79. ^ "Experimenter Drinks 'Heavy Water' at $v,000 a Quart". Popular Science Monthly. Vol. 126, no. four. New York: Popular Science Publishing. Apr 1935. p. 17. Retrieved seven Jan 2011.
80. ^ Müller, Grover C. (June 1937). "Is 'Heavy Water' the Fountain of Youth?". Pop Science Monthly. Vol. 130, no. 6. New York: Popular Science Publishing. pp. 22–23. Retrieved 7 Jan 2011.
81. ^ Miller, Inglis J., Jr.; Mooser, Gregory (Jul 1979). "Taste Responses to Deuterium Oxide". Physiology & Behavior. 23 (1): 69–74. doi:x.1016/0031-9384(79)90124-0. PMID 515218. S2CID 39474797.
82. ^ Weingärtner et al. 2016, p. 29.
83. ^ Zumdahl & Zumdahl 2013, p. 659.
84. ^ ^{a b} Zumdahl & Zumdahl 2013, p. 654.
85. ^ Zumdahl & Zumdahl 2013, p. 984.
86. ^ Zumdahl & Zumdahl 2013, p. 171.
87. ^ "Hydrides". Chemwiki. UC Davis. two Oct 2013. Retrieved 2016-06-25 .
88. ^ Zumdahl & Zumdahl 2013, pp. 932, 936.
89. ^ Zumdahl & Zumdahl 2013, p. 338.
90. ^ Zumdahl & Zumdahl 2013, p. 862.
91. ^ Zumdahl & Zumdahl 2013, p. 981.
92. ^ Charlot 2007, p. 275.
93. ^ ^{a b} Zumdahl & Zumdahl 2013, p. 866.
94. ^ ^{a b} Greenwood & Earnshaw 1997, p. 601.
95. ^ "Enterprise and electrolysis..." Majestic Society of Chemistry. August 2003. Retrieved 2016-06-24 .
96. ^ "Joseph Louis Gay-Lussac, French pharmacist (1778–1850)". 1902 Encyclopedia. Footnote 122-1. Retrieved 2016-05-26 .
97. ^ Lewis, G. Northward.; MacDonald, R. T. (1933). "Concentration of H2 Isotope". The Journal of Chemical Physics. 1 (6): 341. Bibcode:1933JChPh...one..341L. doi:10.1063/i.1749300.
98. ^ ^{a b} Leigh, Favre & Metanomski 1998, p. 34.
99. ^ IUPAC 2005, p. 85.
100. ^ Leigh, Favre & Metanomski 1998, p. 99.
101. ^ "Tetrahydropyran". Pubchem. National Institutes of Health. Retrieved 2016-07-31 .
102. ^ Leigh, Favre & Metanomski 1998, pp. 27–28.
103. ^ "Compound Summary for CID 22247451". Pubchem Compound Database. National Eye for Biotechnology Information.

Bibliography [edit]

• Boyd, Claude E. (2000). "pH, Carbon Dioxide, and Alkalinity". Water Quality. Boston, Massachusetts: Springer. pp. 105–122. doi:10.1007/978-1-4615-4485-2_7. ISBN9781461544852.
• Campbell, Mary Grand.; Farrell, Shawn O. (2007). Biochemistry (6th ed.). Cengage Learning. ISBN978-0-495-39041-ane.
• Campbell, Neil A.; Reece, Jane B. (2009). Biology (8th ed.). Pearson. ISBN978-0-8053-6844-4.
• Campbell, Neil A.; Williamson, Brad; Heyden, Robin J. (2006). Biological science: Exploring Life. Boston, Massachusetts: Pearson Prentice Hall. ISBN978-0-13-250882-7.
• Charlot, M. (2007). Qualitative Inorganic Analysis. Read Books. ISBN978-1-4067-4789-8.
• Greenwood, Norman North.; Earnshaw, Alan (1997). Chemistry of the Elements (second ed.). Butterworth-Heinemann. ISBN978-0-08-037941-eight.
• International Matrimony of Pure and Applied Chemistry (2005-eleven-22). Classification of Inorganic Chemistry: IUPAC Recommendations 2005 (PDF). Royal Society of Chemistry. ISBN978-0-85404-438-2 . Retrieved 2016-07-31 .
• Leigh, G. J.; Favre, H. A; Metanomski, W. Five. (1998). Principles of chemical classification: a guide to IUPAC recommendations (PDF). Oxford: Blackwell Science. ISBN978-0-86542-685-6. OCLC 37341352. Archived from the original (PDF) on 2011-07-26.
• Lewis, William C.M.; Rice, James (1922). A System of Physical Chemical science. Longmans, Green and Co.
• Lide, David R. (2003-06-19). CRC Handbook of Chemistry and Physics, 84th Edition. CRC Handbook. CRC Printing. ISBN9780849304842.
• Reece, Jane B.; Urry, Lisa A.; Cain, Michael L.; Wasserman, Steven A.; Minorsky, Peter 5.; Jackson, Robert B. (2013-11-10). Campbell Biology (10th ed.). Boston, Mass.: Pearson. ISBN9780321775658.
• Riddick, John (1970). Organic Solvents Physical Properties and Methods of Purification . Techniques of Chemical science. Wiley-Interscience. ISBN978-0471927266.
• Precipitous, Robert Phillip (1988-eleven-25). Living Ice: Understanding Glaciers and Glaciation . Cambridge University Press. p. 27. ISBN978-0-521-33009-i.
• Weingärtner, Hermann; Teermann, Ilka; Borchers, Ulrich; Balsaa, Peter; Lutze, Holger V.; Schmidt, Torsten C.; Franck, Ernst Ulrich; Wiegand, Gabriele; Dahmen, Nicolaus; Schwedt, Georg; Frimmel, Fritz H.; Gordalla, Birgit C. (2016). "Water, one. Properties, Assay, and Hydrological Cycle". Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH Verlag GmbH & Co. KGaA. doi:ten.1002/14356007.a28_001.pub3. ISBN9783527306732.
• Zumdahl, Steven S.; Zumdahl, Susan A. (2013). Chemistry (9th ed.). Cengage Learning. ISBN978-1-13-361109-7.

Farther reading [edit]

• Ben-Naim, A. (2011), Molecular Theory of H2o and Aqueous Solutions, World Scientific

External links [edit]

• "Water Properties and Measurements". The states Geological Survey. May 2, 2016. Retrieved Baronial 31, 2016.
• Release on the IAPWS Formulation 1995 for the Thermodynamic Properties of Ordinary H2o Substance for General and Scientific Utilize (simpler conception)
• Online calculator using the IAPWS Supplementary Release on Properties of Liquid H2o at 0.1 MPa, September 2008
• Chaplin, Martin (2019). "Construction and Properties of Water in its Various States". Encyclopedia of Water. Wiley Online Library 2019. pp. one–19. doi:10.1002/9781119300762.wsts0002. ISBN9781119300755. S2CID 213738895.
• Calculation of vapor pressure, liquid density, dynamic liquid viscosity, and surface tension of water
• Water Density Calculator
• Why does ice float in my drink?, NASA

rayciatell.blogspot.com

Source: https://en.wikipedia.org/wiki/Properties_of_water