Water (molecule)




Water (H2O)
IUPAC name Water
Other names Aqua
Hydrogen oxide
Hydrogen hydroxide
Hydrate
Oxidane
Hydroxic acid
Dihydrogen monoxide
Hydroxyl acid
Dihydrogen oxide
Hydrohydroxic acid
μ-Oxido dihydrogen
Light Water
Identifiers
CAS number 7732-18-5
RTECS number ZC0110000
Properties
Molecular formula H2O or HOH
Molar mass 18.01524 g/mol
Appearance transparent, almost
colorless liquid with
a slight hint of blue[1]
Density 1000 kg·m−3, liquid (4 °C)
917 kg·m−3, solid
Melting point

0 K)[2]

Boiling point

100 °C, 212 °F (373.15 K)[2]

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

Water (solvent. Because of this, water in nature and in use is rarely pure, and may have some properties different from those in the laboratory. However, there are many compounds that are essentially, if not completely, insoluble in water. Water is the only common substance found naturally in all three common states of matter—for other substances, see Chemical properties.

Forms of water

See the Water#Overview of types of water

Water can take many forms. The MPa), water molecules assume a supercritical condition, in which liquid-like clusters float within a vapor-like phase.

neutrons.

Physics and chemistry of water

Water is the ambient temperature and pressure, and appears colorless in small quantities, although it has its own intrinsic very light blue hue. Ice also appears colorless, and water vapor is essentially invisible as a gas.[3] Water is primarily a liquid under standard conditions, which is not predicted from its relationship to other analogous hydrides of the hydroxide ion (OH−(aq)). Water is in standard temperature and pressure, and is the only pure substance found naturally on Earth to be so.

Water, ice and vapor

Heat capacity and heat of vaporization

Main article: Enthalpy of vaporization

Water has the second highest hydrogen bonding between its molecules. These two unusual properties allow water to moderate Earth's climate by buffering large fluctuations in temperature.

Density of water and ice

Temp (°C)Density (g/cm³)
300.9956502
250.9970479
220.9977735
200.9982071
150.9991026
100.9997026
40.9999720
00.9998395
−100.998117
−200.993547
−300.983854
The density of water in grams per cubic centimeter
at various temperatures in degrees Celsius [4]
The values below 0 °C refer to supercooled water.

The solid form of most substances is more hydrogen bond.

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

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

Nevertheless, the unusual expansion of freezing water (in ordinary natural settings in relevant biological systems), due to the freezes and forms ice. Since downward convection of colder water is blocked by the density change, any large body of fresh water frozen in winter will have the coldest water near the surface, away from the riverbed or lakebed. This accounts for various little known phenomena of ice characteristics as they relate to ice in lakes and "ice falling out of lakes" as described by early 20th century scientist Horatio D. Craft.

Water will freeze at 0 °C (32 °F, 273 K), however, it can be crystal homogeneous nucleation at almost 231 K (−42 °C) [5].

Density of saltwater and ice

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

As the surface of salt water begins to freeze (at −1.9 °C for normal salinity brine rejection. This more dense saltwater sinks by convection and the replacing seawater is subject to the same process. This provides essentially freshwater ice at −1.9 °C on the surface. The increased density of the seawater beneath the forming ice causes it to sink towards the bottom.

Miscibility and condensation

Main article: Humidity

Water is miscible with many liquids, for example ethanol in all proportions, forming a single homogeneous liquid. On the other hand water and most oils are immiscible usually forming layers according to increasing density from the top.

  As 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 dew. The reverse process accounts for the fog burning off in the morning. If one raises the humidity at room temperature, say 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 condenses out as steam. A gas in this context is referred to as saturated or 100% relative humidity, when the vapor pressure of water in the air is at the 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 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. Water vapor pressure above 100% relative humidity is called super-saturated and can occur if air is rapidly cooled, say by rising suddenly in an updraft.[7]

Vapor Pressures of Water

Temperature (°C) Pressure (torr)
0 4.58
5 6.54
10 9.21
12 10.52
14 11.99
16 13.63
17 14.53
18 15.48
19 16.48
20 17.54
21 18.65
22 19.83
23 21.07
24 22.38
25 23.76

[8]

Compressibility

The compressibility of water is a function of pressure and temperature. At 0 °C in the limit of zero pressure the compressibility is 5.1×10-5 bar−1.[9] In the zero pressure limit the compressibility reaches a minimum of 4.4×10-5 bar−1 around 45 °C before increasing again with increasing temperature. As the pressure is increased the compressibility decreases, being 3.9×10-5 bar−1 at 0 °C and 1000 bar. The m depth, where pressures are 4×107 Pa, there is only a 1.8% decrease in volume.[10]

Triple point

The various triple points of water[11]
Phases in stable equilibrium Pressure Temperature
liquid water, ice I, and water vapour 611.73 Pa 273.16 K
liquid water, ice Ih, and ice III 209.9 MPa 251 K (-22 °C)
liquid water, ice Ih, and gaseous water 612 Pa 0.01 °C
liquid water, ice III, and ice V 350.1 MPa -17.0 °C
liquid water, ice V, and ice 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 Pa. This pressure is quite low, about 1/166 of the normal sea level barometric pressure of 101,325 Pa. The atmospheric surface pressure on planet Mars is remarkably close to the triple point pressure, and the zero-elevation or "sea level" of Mars is defined by the height at which the atmospheric pressure corresponds to the triple point of water.  


Although it is commonly named as "the triple point of water", the stable combination of liquid water, phase diagram of water. 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.[12][11][13]

Mpemba effect

The frost.

Hot ice

Hot ice is the name given to another surprising phenomenon in which water at room temperature can be turned into ice that remains at room temperature by supplying an electric field on the order of 106 volts per meter.[14]

The effect of such electric fields has been suggested as an explanation of cloud formation. The first time cloud ice forms around a clay particle, it requires a temperature of −10 °C, but subsequent freezing around the same clay particle requires a temperature of just −5 °C, suggesting some kind of structural change.[15]

Surface tension

Water drops are stable, due to the high surface tension of water, 72.8 mN/m, the highest of the non-metallic liquids. This can be seen when small quantities of water are put on a surface such as glass: the water stays together as drops. This property is important for life. For example, when water is carried through xylem up stems in plants the strong intermolecular attractions hold the water column together. Strong cohesive properties hold the water column together, and strong adhesive properties stick the water to the xylem, and prevent tension rupture caused by transpiration pull. Other liquids with lower surface tension would have a higher tendency to "rip", forming vacuum or air pockets and rendering the xylem water transport inoperative.

Electrical properties

Pure water containing no ions is an excellent ions in aqueous solution by which an electric current can flow.

Water can be split into its constituent elements, hydrogen and oxygen, by passing a current through it. This process is called nanosiemens per meter of conductance).

Electrical conductivity

Pure water has a low proton conductor).

Dipolar nature of water

  An important feature of water is its hydrogen bonding, and explains many of the properties of water. Certain molecules, such as carbon dioxide, also have a difference in electronegativity between the atoms but the difference is that the shape of carbon dioxide is symmetrically aligned and so the opposing charges cancel one another out. This phenomenon of water can be seen if you hold an electrical source near a thin stream of water falling vertically, causing the stream to bend towards the electrical source.

Although hydrogen bonding is a relatively weak attraction compared to the covalent bonds within the water molecule itself, it is responsible for a number of water's physical properties. One such property is its relatively high specific heat capacity. This high heat capacity makes water a good heat storage medium.

Hydrogen bonding also gives water its unusual behavior when freezing. When cooled to near freezing point, the presence of hydrogen bonds means that the molecules, as they rearrange to minimize their energy, form the hexagonal crystal structure of ice that is actually of lower density: hence the solid form, ice, will float in water. In other words, water expands as it freezes, whereas almost all other materials shrink on solidification.

An interesting consequence of the solid having a lower density than the liquid is that ice will melt if sufficient pressure is applied. With increasing pressure the melting point temperature drops and when the melting point temperature is lower than the ambient temperature the ice begins to melt. A significant increase of pressure is required to lower the melting point temperature —the pressure exerted by an ice skater on the ice would only reduce the melting point by approximately 0.09 °C (0.16 °F).

Electronegative Polarity

Water has a partial negative charge (σ-) near the oxygen atom due to the unshared pairs of electrons, and partial positive charges (σ+) near the hydrogen atoms. In water, this happens because the oxygen atom is more electrons, drawing them closer (along with their negative charge) and making the area around the oxygen atom more negative than the area around both of the hydrogen atoms.

Adhesion

  Water sticks to itself (cohesion) because it is polar. Water also has high adhesion properties because of its polar nature. On extremely clean/smooth glass the water may form a thin film because the molecular forces between glass and water molecules (adhesive forces) are stronger than the cohesive forces. 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 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. Their importance in biology has been extensively studied by V. Adrian Parsegian of the National Institute of Health.[16] They are particularly important when cells are dehydrated by exposure to dry atmospheres or to extracellular freezing.

Surface tension

Main article: Surface tension

 

Water has a high polythene; the water stays together as drops. Just as significantly, air trapped in surface disturbances forms bubbles, which sometimes last long enough to transfer gas molecules to the water. Another surface tension effect is capillary waves which are the surface ripples that form from around the impact of drops on water surfaces, and some times occur with strong subsurface currents flow to the water surface. The apparent elasticity caused by surface tension drives the waves.

Capillary action

Main article: Capillary action

Capillary action refers to the process of water moving up a narrow tube against the force of gravity. It occurs because water adheres to the sides of the tube, and then surface tension tends to straighten the surface making the surface rise, and more water is pulled up through cohesion. The process is repeated as the water flows up the tube until there is enough water that gravity can counteract the adhesive force.

Water as a solvent

  Water is also a good pushed out" from the water, and do not dissolve. Contrary to the common misconception, water and hydrophobic substances does 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 typically 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 general, ionic and polar substances such as van der Waals interactions with nonpolar molecules.

An example of an ionic solute is table 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.

Amphoteric nature of water

Chemically, water is equilibrium is disturbed, the solution becomes acidic (higher concentration of hydronium ions) or basic (higher concentration of hydroxide ions).

Water can act as either an acid or a base in reactions. According to the Brønsted-Lowry system, an acid is defined as a species which donates a proton (an H+ ion) in a reaction, and a base as one which receives a proton. When reacting with a stronger acid, water acts as a base; when reacting with a stronger base, it acts as an acid. For instance, it receives an H+ ion from HCl in the equilibrium:

HCl + H2O ⇌ H3O+ + Cl−

Here water is acting as a base, by receiving an H+ ion.

In the reaction with ammonia, NH3, water donates an H+ ion, and is thus acting as an acid:

NH3 + H2O ⇌ NH4+ + OH−

Acidity in nature

In theory, pure water has a acid rain problems.

Hydrogen bonding in water

A water molecule can form a maximum of four silica in their anomalous behaviour, even though one (water) is a liquid which has a hydrogen bonding network while the other (silica) has a covalent network with a very high melting point. One reason that water is well suited, and chosen, by life-forms, is that it exhibits its unique properties over a temperature regime that suits diverse biological processes, including hydration.

It is believed that hydrogen bond in water is largely due to electrostatic forces and some amount of covalency. The partial covalent nature of hydrogen bond predicted by Linus Pauling in the 1930s is yet to be proven unambiguously by experiments and theoretical calculations.

Quantum properties of molecular water

Although the molecular formula of water is generally considered to be a stable result in molecular thermodynamics, recent work started in 1995 has shown that at certain scales, water may act more like H3/2O than H2O at the quantum level.[17] This result could have significant ramifications at the level of, for example, the oxygen respectively. However, the time-scale of this response is only seen at the level of attoseconds (10-18 seconds), and so is only relevant in highly resolved kinetic and dynamical systems.[18][19]

Heavy Water and isotopologues of water

Hydrogen has three isotopes. The most common, making up more than 95% of water, has 1 proton and 0 neutrons. A second isotope, half-life of 12.32 years. T2O exists in nature only in tiny quantities, being produced primarily via cosmic ray-driven nuclear reactions in the atmosphere. D2O is stable, but differs from H2O in in that it is more dense - hence, "heavy water" - and in that several other physical properties are slightly different from those of common, Hydrogen-1 containing "light water". D2O occurs naturally in ordinary water in very low concentrations. Consumption of pure isolated D2O may affect biochemical processes - ingestion of large amounts impairs kidney and central nervous system function. However, very large amounts of heavy water must be consumed for any toxicity to be apparent, and smaller quantities can be consumed with no ill effects at all.

Transparency

Water's transparency is also an important property of the liquid. If water were not transparent, sunlight, essential to aquatic plants, would not reach into seas and oceans.

History

The properties of water have historically been used to define various triple point of water is a more commonly used standard point today.[20]

The first scientific decomposition of water into hydrogen and oxygen, by Joseph Louis Gay-Lussac and Alexander von Humboldt showed that water is composed of two parts hydrogen and one part oxygen (by volume).

heavy water in 1933.

polymerized form of water that was the subject of much scientific controversy during the late 1960s. The consensus now is that it does not exist.

Pseudoscience concept is water memory.

Systematic naming

The accepted IUPAC name of water is simply "water", although there are two other systematic names which can be used to describe the molecule.

The simplest and best systematic name of water is hydrogen oxide. This is analogous to related compounds such as dioxane and trioxane.

Systematic nomenclature and humor

Main article: Dihydrogen monoxide hoax

Dihydrogen monoxide or DHMO is an overly pedantic systematic covalent name of water. This term has been used in parodies of chemical research that call for this "lethal chemical" to be banned. In reality, a more realistic systematic name would be hydrogen oxide, since the "di-" and "mon-" prefixes are superfluous. hydrogen peroxide, H2O2, is never called "dihydrogen dioxide".

Some overzealous material safety data sheets for water list the following: Caution: May cause drowning![citation needed]

Other systematic names for water include hydroxic acid or hydroxylic acid. Likewise, the systematic alkali name of water is hydrogen hydroxide—both acid and alkali names exist for water because it is able to react both as an acid or an alkali, depending on the strength of the acid or alkali it is reacted with (amphoteric). None of these names are used widely outside of DHMO sites.

See also

References

  1. ^ Braun, Charles L.; Sergei N. Smirnov (1993). "Why is water blue?" (HTML). J. Chem. Educ. 70 (8): 612.
  2. ^ a b Vienna Standard Mean Ocean Water (VSMOW), used for calibration, melts at 273.1500089(10) K (0.000089(10) °C, and boils at 373.1339 K (99.9839 °C)
  3. ^ Braun, Charles L.; Sergei N. Smirnov (1993). "Why is water blue?" (HTML). J. Chem. Educ. 70 (8): 612.
  4. ^ Lide, D. R. (Ed.) (1990). CRC Handbook of Chemistry and Physics (70th Edn.). Boca Raton (FL):CRC Press.
  5. ^ P. G. Debenedetti, P. G., and Stanley, H. E.; "Supercooled and Glassy Water", Physics Today 56 (6), p. 40–46 (2003).
  6. ^ The pressure due to water vapor in the air is called the partial pressure(Boyle's law).
  7. ^ ideal gas law.
  8. ^ Brown, Theodore L., H. Eugene LeMay, Jr., and Bruce E. Burston. Chemistry: The Central Science. 10th ed. Upper Saddle River, NJ: Pearson Education, Inc., 2006.
  9. ^ Fine, R.A. and Millero, F.J. (1973). "Compressibility of water as a function of temperature and pressure". Journal of Chemical Physics 59 (10). doi:10.1063/1.1679903.
  10. ^ a b R. Nave. Bulk Elastic Properties. HyperPhysics. Georgia State University. Retrieved on 2007-10-26.
  11. ^ a b Oliver Schlüter (2003-07-28). "Impact of High Pressure — Low Temperature Processes on Cellular Materials Related to Foods" (PDF). Technischen Universität Berlin.
  12. ^ Gustav Heinrich Johann Apollon Tammann (1925). "The States Of Aggregation". Constable And Company Limited.
  13. ^ William Cudmore McCullagh Lewis and James Rice (1922). A System of Physical Chemistry. Longmans, Green and co.. 
  14. ^ Choi, Eun-Mi; Yoon, Young-Hwan; Lee, Sangyoub; Kang, Heon. "Freezing Transition of Interfacial Water at Room Temperature under Electric Fields". Physical Review Letters 95 (8): 085701. doi:10.1103/PhysRevLett.95.085701.
  15. ^ Connolly PJ, Saunders CPR, Gallagher MW, Bower KN, Flynn MJ, Choularton TW, Whiteway J, Lawson RP (April 2005). "Aircraft observations of the influence of electric fields on the aggregation of ice crystals". Quarterly Journal of the Royal Meteorological Society, Part B 131 (608): 1695–1712.
  16. ^ Physical Forces Organizing Biomolecules (PDF)
  17. ^ Phil Schewe, James Riordon, and Ben Stein. "A Water Molecule's Chemical Formula is Really Not H2O", Physics News Update, 31 Jul 03. 
  18. ^ C. A. Chatzidimitriou-Dreismann, T. Abdul Redah, R. M. F. Streffer and J. Mayers (1997). "Anomalous Deep Inelastic Neutron Scattering from Liquid H2O-D2O: Evidence of Nuclear Quantum Entanglement". Physical Review Letters 79 (15): 2839. doi:10.1103/PhysRevLett.79.2839.
  19. ^ C. A. Chatzidimitriou-Dreismann, M. Vos, C. Kleiner and T. Abdul-Redah (2003). "Comparison of Electron and Neutron Compton Scattering from Entangled Protons in a Solid Polymer". Physical Review Letters 91 (5): 057403-4. doi:10.1103/PhysRevLett.91.057403.
  20. ^ http://home.comcast.net/~igpl/Temperature.html
  21. ^ Leigh, G. J. et al. 1998. Principles of chemical nomenclature: a guide to IUPAC recommendations, p. 99. Blackwell Science Ltd, UK. ISBN 0-86542-685-6
 
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