LABORATORIUM KIMIA FISIKA Percobaan : VISKOSITAS Kelompok : VIII A Nama : 1. Clarissa Amalia NRP. 2313 030 015 2. Daniatus Syarh Hajj NRP. 2313 030 023 3. Aprise Mujiartono NRP. 2313 030 051 4. Fano Alfian Ardyansyah NRP. 2313 030 079 5. Khairul Anam NRP. 2313 030 097 Tanggal Percobaan : 23 September 2013 Tanggal Penyerahan : 30 September 2013 Dosen Pembimbing : Nurlaili Humaidah ST, MT Asisten Laboratorium : Dhaniar Rulandri W PROGRAM STUDI D3 TEKNIK KIMIA FAKULTAS TEKNOLOGI INDUSTRI INSTITUT TEKNOLOGI SEPULUH NOPEMBER SURABAYA 2013
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Temperature dependence of liquid viscosityFrom Wikipedia, the free encyclopedia
The temperature dependence of liquid viscosity is the phenomenon by which liquid viscosity tends todecrease (or, alternatively, its fluidity tends to increase) as its temperature increases. This can be observed, forexample, by watching how cooking oil appears to move more fluidly upon a frying pan after being heated by astove.
Contents
1 Physical causes
2 Models for shear viscosity
2.1 Exponential model
2.2 Arrhenius model
2.3 Williams-Landel-Ferry model2.4 Viscosity of water
3 Models for kinematic viscosity3.1 Walther formula
3.2 Wright model
3.3 Seeton model4 Notes
5 See also
Physical causes
A molecular view of liquids can be used for a qualitative picture of the process of decrease in the shear (or bulk)viscosity of a simple fluid with temperature. As the temperature increases, the time of interaction betweenneighbouring molecules of a liquid decreases because of the increased velocities of individual molecules. Themacroscopic effect is that the intermolecular force appears to decrease and so does the bulk (or shear)
viscosity.[citation needed] The actual process can be quite complex and is typically represented by simplified
mathematical or empirical models, some of which are discussed below.[1] The models are valid over limitedtemperature ranges and for selected materials.
Models for shear viscosity
Exponential model
An exponential model for the temperature-dependence of shear viscosity (μ) was first proposed by Reynolds in
1886.[2]
where T is temperature and and are coefficients. See first-order fluid and second-order fluid. This is anempirical model that usually works for a limited range of temperatures.
The model is based on the assumption that the fluid flow obeys the Arrhenius equation for molecular kinetics:
where T is temperature, is a coefficient, E is the activation energy and R is the universal gas constant. A first-order fluid is another name for a power-law fluid with exponential dependence of viscosity on temperature.
Williams-Landel-Ferry model
See also: Williams-Landel-Ferry equation
The Williams-Landel-Ferry model, or WLF for short, is usually used for polymer melts or other fluids thathave a glass transition temperature.
The model is:
where T-temperature, , , and are empiric parameters (only three of them are independent from
each other).
If one selects the parameter based on the glass transition temperature, then the parameters , become
very similar for the wide class of polymers. Typically, if is set to match the glass transition temperature ,
we get
17.44
and
K.
Van Krevelen recommends to choose
K, then
and
101.6 K.
Using such universal parameters allows one to guess the temperature dependence of a polymer by knowingthe viscosity at a single temperature.
In reality the universal parameters are not that universal, and it is much better to fit the WLF parameters fromthe experimental data.
Viscosity of water equation accurate to within 2.5% from 0 °C to 370 °C:[citation needed]
where T has units of Kelvin, and μ has units of N·s/m².
Models for kinematic viscosity
The effect of temperature on the kinematic viscosity (ν) has also been described by a number of empiricalequations.
Walther formula
The Walther formula[1] is typically written in the form
where λ is a shift constant, and A, B are empirical parameters.
Wright model
The Wright model[1] has the form
where an addition function f(ν), often a polynomial fit to experimental data, has been added to the Waltherformula.
Seeton model
The Seeton model[1]is based on curve fitting the viscosity dependence of many liquids (refrigerants,hydrocarbons and lubricants) versus temperature and applies over a large temperature and viscosity range:
where T is absolute temperature in kelvins, is the kinematic viscosity in centistokes, is the zero order
modified Bessel function of the second kind, and A and B are liquid specific values. This form should not beapplied to ammonia or water viscosity over a large temperature range.
For liquid metal viscosity as a function of temperature, Seeton proposed:
Notes
1. ̂a b c d Seeton, Christopher J. (2006), "Viscosity-temperature correlation for liquids"
Properties of waterFrom Wikipedia, the free encyclopedia
Water (H2O) is the most abundant compound onEarth's surface, covering about 70 percent of theplanet. In nature, water exists in liquid, solid, andgaseous states. It is in dynamic equilibrium betweenthe liquid and gas states at standard temperatureand pressure. At room temperature, it is a tastelessand odorless liquid, nearly colorless with a hint ofblue. Many substances dissolve in water and it iscommonly referred to as the universal solvent.Because of this, water in nature and in use is rarelypure and some of its properties may vary slightlyfrom those of the pure substance. However, thereare also many compounds that are essentially, if notcompletely, insoluble in water. Water is the onlycommon substance found naturally in all threecommon states of matter and it is essential for all
life on Earth.[4] Water usually makes up 55% to
78% of the human body.[5]
In keeping with the basic rules of chemicalnomenclature, water would have a systematic name
of dihydrogen monoxide,[6] but this is not amongthe names published by the International Union of
Pure and Applied Chemistry[7] and, rather thanbeing used in a chemical context, the name isalmost exclusively used as a humorous way to referto water.
Contents
1 Forms of water
2 Physics and chemistry
2.1 Water, ice and vapor
2.1.1 Heat capacity and heatsof vaporization and fusion
Main hazards Drowning (see also Dihydrogen monoxide
hoax)
Water intoxication
NFPA 704
Related compounds
Other cations Hydrogen sulfide
Hydrogen selenide
Hydrogen telluride
Hydrogen polonide
Hydrogen peroxide
Related solvents acetone
methanol
Related
compounds
water vapor
ice
heavy water
(verify) (what is: / ?)
Except where noted otherwise, data are given for materials
2.2.1 Electrical conductivity2.2.2 Electrolysis
2.3 Static dielectric constant
2.4 Polarity and hydrogen bonding
2.4.1 Cohesion and adhesion
2.4.2 Surface tension
2.4.3 Capillary action
2.4.4 Water as a solvent
2.5 Water in acid-base reactions
2.5.1 Ligand chemistry
2.5.2 Organic chemistry
2.5.3 Acidity in nature2.6 Water in redox reactions
2.7 Geochemistry
2.8 Transparency
2.9 Heavy water and isotopologues3 History4 Systematic naming
5 See also6 Notes
7 References8 External links
Forms of water
Like many substances, water can take numerousforms that are broadly categorized by phase ofmatter. The liquid phase is the most commonamong water's phases (within the Earth'satmosphere and surface) and is the form that isgenerally denoted by the word "water." The solidphase of water is known as ice and commonlytakes the structure of hard, amalgamated crystals,such as ice cubes, or loosely accumulated granularcrystals, like snow. For a list of the many differentcrystalline and amorphous forms of solid H2O, see
the article ice. The gaseous phase of water isknown as water vapor (or steam), and ischaracterized by water assuming the configurationof a transparent cloud. (Note that the visible steamand clouds are, in fact, water in the liquid form asminute droplets suspended in the air.) The fourthstate of water, that of a supercritical fluid, is muchless common than the other three and only rarelyoccurs in nature, in extremely uninhabitableconditions. When water achieves a specific criticaltemperature and a specific critical pressure (647 Kand 22.064 MPa), liquid and gas phase merge to
one homogeneous fluid phase, with properties ofboth gas and liquid. One example of naturallyoccurring supercritical water is found in the hottestparts of deep water hydrothermal vents, in which water is heated to the critical temperature by scalding volcanicplumes and achieves the critical pressure because of the crushing weight of the ocean at the extreme depths atwhich the vents are located. Additionally, anywhere there is volcanic activity below a depth of 2.25 km
(1.40 mi) can be expected to have water in the supercritical phase.[8]
Vienna Standard Mean Ocean Water is the current international standard for water isotopes. Naturallyoccurring water is almost completely composed of the neutron-less hydrogen isotope protium. Only 155 ppm
include deuterium (2H or D), a hydrogen isotope with one neutron, and fewer than 20 parts per quintillion
include tritium (3H or T), which has two.
Heavy water is water with a higher-than-average deuterium content, up to 100%. Chemically, it is similar butnot identical to normal water. This is because the nucleus of deuterium is twice as heavy as protium, and thiscauses noticeable differences in bonding energies. Because water molecules exchange hydrogen atoms with oneanother, hydrogen deuterium oxide (DOH) is much more common in low-purity heavy water than pure
dideuterium monoxide (D2O). Humans are generally unaware of taste differences,[9] but sometimes report a
burning sensation[10] or sweet flavor.[11] Rats, however, are able to avoid heavy water by smell.[12] Toxic to
many animals,[12] heavy water is used in the nuclear reactor industry to moderate (slow down) neutrons. Lightwater reactors are also common, where "light" simply designates normal water.
Light water more specifically refers to deuterium-depleted water (DDW), water in which the deuterium contenthas been reduced below the standard 155 ppm level.
Physics and chemistry
See also: Water chemistry analysis
Water is the chemical substance with chemical formula H2O: one molecule of water has two hydrogen atoms
covalently bonded to a single oxygen atom.[13] Water is a tasteless, odorless liquid at ambient temperature andpressure, 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.[2]
Water is primarily a liquid under standard conditions, which is not predicted from its relationship to otheranalogous hydrides of the oxygen family in the periodic table, which are gases such as hydrogen sulfide. Theelements surrounding oxygen in the periodic table, nitrogen, fluorine, phosphorus, sulfur and chlorine, allcombine with hydrogen to produce gases under standard conditions. The reason that water forms a liquid is thatoxygen is more electronegative than all of these elements with the exception of fluorine. Oxygen attractselectrons much more strongly than hydrogen, resulting in a net positive charge on the hydrogen atoms, and a netnegative charge on the oxygen atom. The presence of a charge on each of these atoms gives each watermolecule a net dipole moment. Electrical attraction between water molecules due to this dipole pulls individualmolecules closer together, making it more difficult to separate the molecules and therefore raising the boilingpoint. This attraction is known as hydrogen bonding. The molecules of water are constantly moving in relation toeach other, and the hydrogen bonds are continually breaking and reforming at timescales faster than 200
femtoseconds.[14] However, this bond is sufficiently strong to create many of the peculiar properties of water,such as those that make it integral to life. Water can be described as a polar liquid that slightly dissociates
disproportionately into the hydronium ion (H3O+(aq)) and an associated hydroxide ion (OH
The dissociation constant for this dissociation is commonly symbolized as Kw and has a value of about 10−14 at
25 °C; see "Water (data page)" and "Self-ionization of water" for more information.
Percentage of elements in water by mass: 11.1% hydrogen, 88.9% oxygen.[15]
The self-diffusion coefficient of water is 2.299·10−9 m²·s−1 [16]
Water, ice and vapor
Heat capacity and heats of vaporization and fusion
Main article: Enthalpy of vaporization
Water has a very high specificheat capacity – the secondhighest among all theheteroatomic species (afterammonia), as well as a highheat of vaporization(40.65 kJ/mol or 2257 kJ/kgat the normal boiling point),both of which are a result ofthe extensive hydrogenbonding between itsmolecules. These two unusualproperties allow water tomoderate Earth's climate bybuffering large fluctuations intemperature. According to
Josh Willis, of NASA's Jet Propulsion Laboratory, the oceans absorbone thousand times more heat than the atmosphere (air) and are holding
80 to 90% of the heat of global warming.[18]
The specific enthalpy of fusion of water is 333.55 kJ/kg at 0 °C. Ofcommon substances, only that of ammonia is higher. This propertyconfers resistance to melting on the ice of glaciers and drift ice. Beforeand since the advent of mechanical refrigeration, ice was and still is incommon use for retarding food spoilage.
Note that the specific heat capacity of ice at −10 °C is about 2.05 J/(g·K) and that the heat capacity of steam at100 °C is about 2.080 J/(g·K).
Density of water and ice
The density of water is approximately one gram per cubic centimeter.It is dependent on its temperature, but the relation is not linear and isunimodal rather than monotonic (see table at left). When cooled fromroom temperature liquid water becomes increasingly dense, as withother substances, but at approximately 4 °C (39 °F), pure waterreaches its maximum density. As it is cooled further, it expands tobecome less dense. This unusual negative thermal expansion isattributed to strong, orientation-dependent, intermolecular interactions
and is also observed in molten silica.[22]
The solid form of most substances is denser than the liquid phase;thus, a block of most solids will sink in the liquid. However, a blockof ice floats in liquid water because ice is less dense. Upon freezing, the density of water decreases by about
9%.[23] This is due to the 'cooling' of intermolecular vibrations allowing the molecules to form steady hydrogenbonds with their neighbors and thereby gradually locking into positions reminiscent of the hexagonal packingachieved upon freezing to ice Ih. Whereas the hydrogen bonds are shorter in the crystal than in the liquid, this
locking effect reduces the average coordination number of molecules as the liquid approaches nucleation. Othersubstances that expand on freezing are silicon, gallium, germanium, antimony, bismuth, plutonium and alsochemical compounds that form spacious crystal lattices with tetrahedral coordination.
Only ordinary hexagonal ice is less dense than the liquid. Under increasing pressure, ice undergoes a number oftransitions to other allotropic forms with higher density than liquid water, such as ice II, ice III, high-densityamorphous ice (HDA), and very-high-density amorphous ice (VHDA).
Water also expandssignificantly as thetemperature increases.Water near the boilingpoint is about 96% asdense as water at 4°C.
The melting point ofice is 0 °C (32 °F,273.15 K) at standard pressure, however, pure liquid water canbe supercooled well below that temperature without freezing ifthe liquid is not mechanically disturbed. It can remain in a fluidstate down to its homogeneous nucleation point of approximately
231 K (−42 °C).[24] The melting point of ordinary hexagonal icefalls slightly under moderately high pressures, but as icetransforms into its allotropes (see crystalline states of ice) above209.9 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 of Ice VII[25]).
A significant increase of pressure is required to lower the meltingpoint of ordinary ice—the pressure exerted by an ice skater onthe ice only reduces the melting point by approximately 0.09 °C
(0.16 °F).[citation needed]
These properties of water have important consequences in itsrole in Earth's ecosystem. Water at a temperature of 4 °C willalways accumulate at the bottom of freshwater lakes,
irrespective of the temperature in the atmosphere. Since water and ice are poor conductors of heat[26] (goodinsulators) it is unlikely that sufficiently deep lakes will freeze completely, unless stirred by strong currents thatmix cooler and warmer water and accelerate the cooling. In warming weather, chunks of ice float, rather thansink to the bottom where they might melt extremely slowly. These properties therefore allow aquatic life in thelake to survive during the winter.
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 stillfloats in the oceans, otherwise they would freeze from the bottom up. However, the salt content of oceanslowers the freezing point by about 2 °C (see here for explanation) and lowers the temperature of the densitymaximum of water to the freezing point. This is why, in ocean water, the downward convection of colder wateris not blocked by an expansion of water as it becomes colder near the freezing point. The oceans' cold waternear the freezing point continues to sink. For this reason, any creature attempting to survive at the bottom ofsuch cold water as the Arctic Ocean generally lives in water that is 4 °C colder than the temperature at thebottom of frozen-over fresh water lakes and rivers in the winter.
In cold countries, when the temperature of fresh water reaches 4 °C, the layers of water near the top in contactwith cold air continue to lose heat energy and their temperature falls below 4 °C. On cooling below 4 °C, theselayers do not sink but may rise up as fresh water has a maximum density at 4 °C. (Refer: Polarity and hydrogenbonding) Due to this, the layer of water at 4 °C remains at the bottom and above this layers of water 3 °C, 2°C, 1 °C and 0 °C are formed. Because ice is a poor conductor of heat, it does not absorb heat energy from
the water beneath the layer of ice whichprevents the water freezing. Thus, aquatic
creatures survive in such places.[citation needed]
As the surface of salt water begins to freeze (at−1.9 °C for normal salinity seawater, 3.5%) theice that forms is essentially salt free with adensity approximately equal to that offreshwater ice. This ice floats on the surface andthe salt that is "frozen out" adds to the salinityand density of the seawater just below it, in aprocess known as brine rejection. This densersaltwater sinks by convection and the replacingseawater is subject to the same process. Thisprovides essentially freshwater ice at −1.9 °Con the surface. The increased density of theseawater beneath the forming ice causes it to sink towards the bottom. On a large scale, the process of brinerejection and sinking cold salty water results in ocean currents forming to transport such water away from thePoles, leading to a global system of currents called the thermohaline circulation. One potential consequence ofglobal warming is that the loss of Arctic and Antarctic ice could result in the loss of these currents as well, whichcould have unforeseeable consequences on near and distant climates.
Miscibility and condensation
Main article: Humidity
Water is miscible with many liquids, forexample ethanol in all proportions,forming a single homogeneous liquid.On the other hand, water and most oilsare immiscible usually forming layersaccording to increasing density from thetop.
As a gas, water vapor is completelymiscible with air. On the other hand themaximum water vapor pressure that isthermodynamically stable with the liquid(or solid) at a given temperature isrelatively low compared with totalatmospheric pressure. For example, if
the vapor partial pressure[27] is 2% ofatmospheric pressure and the air iscooled from 25 °C, starting at about 22°C water will start to condense, definingthe dew point, and creating fog or dew.The reverse process accounts for thefog burning off in the morning. If thehumidity is increased at room temperature, for example, by running a hot shower or a bath, and the temperaturestays about the same, the vapor soon reaches the pressure for phase change, and then condenses out as minutewater droplets, commonly referred to as steam.
A gas in this context is referred to as saturated or 100% relative humidity, when the vapor pressure of water inthe air is at the equilibrium with vapor pressure due to (liquid) water; water (or ice, if cool enough) will fail tolose 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 vaporpressure, is much more useful. Water vapor pressure above 100% relative humidity is called super-saturated
and can occur if air is rapidly cooled, for example, by rising suddenly in an updraft.[28]
liquid water, ice Ih, and water vapor 611.73 Pa 273.16 K (0.01 °C)
liquid water, ice Ih, and ice III 209.9 MPa 251 K (−22 °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
Water phase diagram: Y-axis = Pressure in pascals (10n); X-axis =
temperature in kelvins; S = solid; L = liquid; V = vapor; CP = critical
point; TP = triple point of water
The compressibility of water is a function of pressure and temperature. At 0 °C, at the limit of zero pressure, the
compressibility is 5.1 × 10−10 Pa−1.[30] At the zero-pressure limit, the compressibility reaches a minimum of
4.4 × 10−10 Pa−1 around 45 °C before increasing again with increasing temperature. As the pressure is
increased, the compressibility decreases, being 3.9 × 10−10 Pa−1 at 0 °C and 100 MPa.
The bulk modulus of water is 2.2 GPa.[31] The low compressibility of non-gases, 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 km depth, where pressures are 40 MPa, there is only a 1.8% decrease in volume.[31]
Triple point
The temperature and pressure atwhich solid, liquid, and gaseous watercoexist in equilibrium is called thetriple point of water. This point isused to define the units oftemperature (the kelvin, the SI unit ofthermodynamic temperature and,indirectly, the degree Celsius andeven the degree Fahrenheit).
As a consequence, water's triplepoint temperature is a prescribedvalue rather than a measured quantity.
The triple point is at a temperature of273.16 K (0.01 °C) by convention,and at a pressure of 611.73 Pa. This
pressure is quite low, about 1⁄166 of the
normal sea level barometric pressure of101,325 Pa. The atmospheric surfacepressure on planet Mars is 610.5 Pa,which is remarkably close to the triplepoint pressure. The altitude of thissurface pressure was used to definezero-elevation or "sea level" on that
planet.[33]
Although it is commonly named as "thetriple point of water", the stablecombination of liquid water, ice I, andwater vapor is but one of several triplepoints on the phase diagram of water.Gustav Heinrich Johann ApollonTammann in Göttingen produced dataon several other triple points in the early20th century. Kamb and others
documented further triple points in the 1960s.[32][34][35]
Pure water containing no exogenous ions is an excellent insulator, but not even "deionized" water is completelyfree of ions. Water undergoes auto-ionization in the liquid state, when two water molecules form one hydroxide
anion (OH−) and one hydronium cation (H3O+).
Because water is such a good solvent, it almost always has some solute dissolved in it, often a salt. If water has
even a tiny amount of such an impurity, then it can conduct electricity far more readily.[citation needed]
It is known that the theoretical maximum electrical resistivity for water is approximately 182 kΩ·m at 25 °C.This figure agrees well with what is typically seen on reverse osmosis, ultra-filtered and deionized ultra-purewater systems used, for instance, in semiconductor manufacturing plants. A salt or acid contaminant levelexceeding 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 slight electrical conductivity of 0.055 µS/cm at 25 °C.Water can also be electrolyzed into oxygen and hydrogen gases but in the absence of dissolved ions this is avery slow process, as very little current is conducted. In ice, the primary charge carriers are protons (see proton
conductor).[36]
Electrolysis
Main article: Electrolysis of water
Water can be split into its constituent elements, hydrogen and oxygen, by passing an electric current through it.
This process is called electrolysis. Water molecules naturally dissociate into H+ and OH
− ions, which are
attracted toward the cathode and anode, respectively. At the cathode, two H+ ions pick up electrons and form
H2 gas. At the anode, four OH− ions combine and release O2 gas, molecular water, and four electrons. The
gases produced bubble to the surface, where they can be collected. The standard potential of the waterelectrolysis cell (when heat is added to the reaction) is a minimum of 1.23 V at 25 °C. The operating potential isactually 1.48 V (or above) in practical electrolysis when heat input is negligible.
One of the important properties of water is that it has a high dielectric constant. This constant shows its ability tomake electrostatic bonds with other molecules, meaning it can eliminate the attraction of the opposite charges ofthe surrounding ions.
Polarity and hydrogen bonding
See also: Chemical polarity
An important feature of water is its polar nature. The water molecule forms an angle, with hydrogen atoms at thetips and oxygen at the vertex. This angle formed is 104.3 degrees as opposed to the typical tetrahedral angle of109 degrees. Because oxygen has a higher electronegativity than hydrogen, the side of the molecule with the
oxygen atom has a partial negative charge. Also the presence of the lone pairs tend to push the oxygen away.An object with such a charge difference is called a dipole meaning two poles. The oxygen end is partiallynegative and the hydrogen end is partially positive, because of this the direction of the dipole moment pointsfrom the oxygen towards the center of the hydrogens. The chargedifferences cause water molecules to be attracted to each other (therelatively positive areas being attracted to the relatively negativeareas) and to other polar molecules. This attraction contributes tohydrogen bonding, and explains many of the properties of water, such
as solvent action.[37]
A water molecule can forma maximum of fourhydrogen bonds because itcan accept two and donatetwo hydrogen atoms. Othermolecules like hydrogenfluoride, ammonia,methanol form hydrogen bonds but they do not show anomalousbehavior of thermodynamic, kinetic or structural properties like thoseobserved in water. The answer to the apparent difference betweenwater and other hydrogen bonding liquids lies in the fact that apartfrom water none of the hydrogen bonding molecules can form fourhydrogen bonds, either due to an inability to donate/accept hydrogensor due to steric effects in bulky residues. In water, local tetrahedralorder due to the four hydrogen bonds gives rise to an open structureand a 3-dimensional bonding network, resulting in the anomalous
decrease of density when cooled below 4 °C.
Although hydrogen bonding is a relatively weak attraction compared to the covalent bonds within the watermolecule itself, it is responsible for a number of water's physical properties. One such property is its relativelyhigh melting and boiling point temperatures; more energy is required to break the hydrogen bonds between
molecules. The similar compound hydrogen sulfide (H2S), which has much weaker hydrogen bonding, is a gas atroom temperature even though it has twice the molecular mass of water. The extra bonding between watermolecules also gives liquid water a large specific heat capacity. This high heat capacity makes water a good heatstorage medium (coolant) and heat shield.
Cohesion and adhesion
Water molecules stay close to each other (cohesion), due to thecollective action of hydrogen bonds between water molecules. Thesehydrogen bonds are constantly breaking, with new bonds beingformed with different water molecules; but at any given time in asample of liquid water, a large portion of the molecules are held
together by such bonds.[38]
Water also has high adhesion properties because of its polar nature.On extremely clean/smooth glass the water may form a thin filmbecause the molecular forces between glass and water molecules(adhesive forces) are stronger than the cohesive forces. In biologicalcells and organelles, water is in contact with membrane and proteinsurfaces 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 stronglyheld 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. They are important in biology,
particularly when cells are dehydrated by exposure to dry atmospheres or to extracellular freezing.[39]
Surface tension
Main article: Surface tension
Water has a high surface tension of72.8 mN/m at room temperature, causedby the strong cohesion between watermolecules, the highest of the common non-ionic, non-metallic liquids. This can be seenwhen small quantities of water are placedonto a sorption-free (non-adsorbent andnon-absorbent) surface, such aspolyethylene or Teflon, and the water staystogether as drops. Just as significantly, airtrapped in surface disturbances formsbubbles, which sometimes last long enough
to transfer gas molecules to the water.[citation needed]
Another surface tension effect is capillary waves, which are the surface ripples thatform around the impacts of drops on water surfaces, and sometimes occur withstrong subsurface currents flowing to the water surface. The apparent elasticitycaused by surface tension drives the waves.
Capillary action
Main article: Capillary action
Due to an interplay of the forces of adhesion and surface tension, water exhibits capillary action whereby waterrises into a narrow tube against the force of gravity. Water adheres to the inside wall of the tube and surfacetension tends to straighten the surface causing a surface rise and more water is pulled up through cohesion. Theprocess continues as the water flows up the tube until there is enough water such that gravity balances theadhesive force.
Surface tension and capillary action are important in biology. For example, when water is carried through xylemup stems in plants, the strong intermolecular attractions (cohesion) hold the water column together and adhesiveproperties maintain the water attachment to the xylem and prevent tension rupture caused by transpiration pull.
Water as a solvent
Main article: Aqueous solution
Water is also a good solvent, due to its polarity. Substances that will mix well and dissolve in water (e.g. salts)are known as hydrophilic ("water-loving") substances, while those that do not mix well with water (e.g. fats andoils), are known as hydrophobic ("water-fearing") substances. The ability of a substance to dissolve in water isdetermined by whether or not the substance can match or better the strong attractive forces that watermolecules generate between other water molecules. If a substance has properties that do not allow it to
overcome these strong intermolecular forces, the molecules are "pushed out" from the water, and do notdissolve. Contrary to the common misconception, water and hydrophobic substances do not "repel", and thehydration 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 relativelysmall size of water molecules typically allows many water molecules tosurround one molecule of solute. The partially negative dipole ends of thewater are attracted to positively charged components of the solute, and viceversa for the positive dipole ends.
In general, ionic and polar substances such as acids, alcohols, and salts arerelatively soluble in water, and non-polar substances such as fats and oils arenot. Non-polar molecules stay together in water because it is energeticallymore favorable for the water molecules to hydrogen bond to each other thanto engage in van der Waals interactions with non-polar molecules.
An example of an ionic solute is table salt; the sodium chloride, NaCl,
separates into Na+ cations and Cl
− anions, each being surrounded by water
molecules. The ions are then easily transported away from their crystallinelattice into solution. An example of a nonionic solute is table sugar. The waterdipoles make hydrogen bonds with the polar regions of the sugar molecule(OH groups) and allow it to be carried away into solution.
Water in acid-base reactions
Chemically, water is amphoteric: it can act as either an acid or a base inchemical reactions. According to the Brønsted-Lowry definition, an acid is defined as a species which donates a
proton (a 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, water receives an H+
ion from HCl when hydrochloric acid is formed:
HCl (acid) + H2O (base) H3O+ + Cl
−
In the reaction with ammonia, NH3, water donates a H+ ion, and is thus acting as an acid:
NH3 (base) + H2O (acid) NH+4 + OH
−
Because the oxygen atom in water has two lone pairs, water often acts as a Lewis base, or electron pair donor,in reactions with Lewis acids, although it can also react with Lewis bases, forming hydrogen bonds between theelectron pair donors and the hydrogen atoms of water. HSAB theory describes water as both a weak hard acidand a weak hard base, meaning that it reacts preferentially with other hard species:
H+ (Lewis acid) + H2O (Lewis base) → H3O
+
Fe3+
(Lewis acid) + H2O (Lewis base) → Fe(H2O)3+6
Cl− (Lewis base) + H2O (Lewis acid) → Cl(H2O)−
6
When a salt of a weak acid or of a weak base is dissolved in water, water can partially hydrolyze the salt,producing the corresponding base or acid, which gives aqueous solutions of soap and baking soda their basicpH:
Water's Lewis base character makes it a common ligand in transition metal complexes, examples of which range
from solvated ions, such as Fe(H2O)3+6 , to perrhenic acid, which contains two water molecules coordinated to a
rhenium atom, to various solid hydrates, such as CoCl2·6H2O. Water is typically a monodentate ligand, it formsonly one bond with the central atom.
Organic chemistry
As a hard base, water reacts readily with organic carbocations, for example in hydration reaction, in which 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 addition of water to an organic molecule cleaves themolecule in two, hydrolysis is said to occur. Notable examples of hydrolysis are saponification of fats anddigestion of proteins and polysaccharides. Water can also be a leaving group in SN2 substitution and E2
elimination reactions, the latter is then known as dehydration reaction.
Acidity in nature
Pure water has the concentration of hydroxide ions (OH−) equal to that of the hydronium (H3O
+) or hydrogen
(H+) ions, which gives pH of 7 at 298 K. In practice, pure water is very difficult to produce. Water left exposed
to air for any length of time will dissolve carbon dioxide, forming a dilute solution of carbonic acid, with a limitingpH of about 5.7. As cloud droplets form in the atmosphere and as raindrops fall through the air minor amounts
of CO2 are absorbed, and thus most rain is slightly acidic. If high amounts of nitrogen and sulfur oxides arepresent in the air, they too will dissolve into the cloud and rain drops, producing acid rain.
Water in redox reactions
Water contains hydrogen in oxidation state +1 and oxygen in oxidation state −2. Because of that, water oxidizes
chemicals with reduction potential below the potential of H+/H2, such as hydrides, alkali and alkaline earth
metals (except for beryllium), etc. Some other reactive metals, such as aluminum, are oxidized by water as well,but their oxides are not soluble, and the reaction stops because of passivation. Note, however, that rusting ofiron is a reaction between iron and oxygen, dissolved in water, not between iron and water.
2 Na + 2 H2O → 2 NaOH + H2
Water itself can be oxidized, emitting oxygen gas, but very few oxidants react with water even if their reduction
potential is greater than the potential of O2/O2−
. Almost all such reactions require a catalyst.[40]
4 AgF2 + 2 H2O → 4 AgF + 4 HF + O2
Geochemistry
Action of water on rock over long periods of time typically leads to weathering and water erosion, physicalprocesses that convert solid rocks and minerals into soil and sediment, but under some conditions chemicalreactions with water occur as well, resulting in metasomatism or mineral hydration, a type of chemical alterationof a rock which produces clay minerals in nature and also occurs when Portland cement hardens.
Water ice can form clathrate compounds, known as clathrate hydrates, with a variety of small molecules that can
be embedded in its spacious crystal lattice. The most notable of these is methane clathrate, 4CH4·23H2O,naturally found in large quantities on the ocean floor.
Transparency
Main article: Water absorption
Water is relatively transparent to visible light, near ultraviolet light, and far-red light, but it absorbs mostultraviolet light, infrared light, and microwaves. Most photoreceptors and photosynthetic pigments utilize theportion of the light spectrum that is transmitted well through water. Microwave ovens take advantage of water'sopacity to microwave radiation to heat the water inside of foods. The very weak onset of absorption in the redend of the visible spectrum lends water its intrinsic blue hue (see Color of water).
Heavy water and isotopologues
Several isotopes of both hydrogen and oxygen exist, giving rise to several known isotopologues of water.
Hydrogen occurs naturally in three isotopes. The most common (1H) accounting for more than 99.98% ofhydrogen in water, consists of only a single proton in its nucleus. A second, stable isotope, deuterium (chemical
symbol D or 2H), has an additional neutron. Deuterium oxide, D2O, is also known as heavy water because of itshigher density. It is used in nuclear reactors as a neutron moderator. The third isotope, tritium, has 1 proton and
2 neutrons, and is radioactive, decaying with a half-life of 4500 days. T2O exists in nature only in minutequantities, being produced primarily via cosmic ray-induced nuclear reactions in the atmosphere. Water with one
deuterium atom HDO occurs naturally in ordinary water in low concentrations (~0.03%) and D2O in far loweramounts (0.000003%).
The most notable physical differences between H2O and D2O, 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.
The difference in boiling points allows the isotopologues to be separated. The self-diffusion coefficient of H2O at
25°C is 23% higher than the value of D2O.[41]
Consumption of pure isolated D2O may affect biochemical processes – ingestion of large amounts impairskidney and central nervous system function. Small quantities can be consumed without any ill-effects, and evenvery large amounts of heavy water must be consumed for any toxicity to become apparent.
Oxygen also has three stable isotopes, with 16
O present in 99.76%, 17
O in 0.04%, and 18
O in 0.2% of water
molecules.[42]
History
The first decomposition of water into hydrogen and oxygen, by electrolysis, was done in 1800 by an Englishchemist William Nicholson. In 1805, Joseph Louis Gay-Lussac and Alexander von Humboldt showed thatwater is composed of two parts hydrogen and one part oxygen.
Gilbert Newton Lewis isolated the first sample of pure heavy water in 1933.
The properties of water have historically been used to define various temperature scales. Notably, the Kelvin,Celsius, Rankine, and Fahrenheit scales were, or currently are, defined by the freezing and boiling points ofwater. The less common scales of Delisle, Newton, Réaumur and Rømer were defined similarly. The triple point
of water is a more commonly used standard point today.[43]
Systematic naming
The accepted IUPAC name of water is oxidane[44] or simply water, or its equivalent in different languages,
although there are other systematic names which can be used to describe the molecule.[45]
The simplest systematic name of water is hydrogen oxide. This is analogous to related compounds such ashydrogen peroxide, hydrogen sulfide, and deuterium oxide (heavy water). Another systematic name, oxidane, is
accepted by IUPAC as a parent name for the systematic naming of oxygen-based substituent groups,[46]
although even these commonly have other recommended names. For example, the name hydroxyl isrecommended over oxidanyl for the –OH group. The name oxane is explicitly mentioned by the IUPAC asbeing unsuitable for this purpose, since it is already the name of a cyclic ether also known as tetrahydropyran.
The polarized form of the water molecule, H+OH−, is also called hydron hydroxide by IUPAC
nomenclature.[47]
Dihydrogen monoxide (DHMO) is a rarely used name of water. This term has been used in various hoaxesthat call for this "lethal chemical" to be banned, such as in the dihydrogen monoxide hoax. Other systematicnames for water include hydroxic acid, hydroxylic acid, and hydrogen hydroxide. Both acid and alkali namesexist for water because it is amphoteric (able to react both as an acid or an alkali). None of these exotic namesare used widely.
See also
Double distilled waterElectromagnetic absorption by water
Flexible SPC water modelFluid dynamicsOptical properties of water and ice
Superheated waterTrioxidane
Viscosity of waterWater cluster
Water dimerWater thread experiment
Notes
1. ^ Definition of Hydrol (http://www.merriam-webster.com/dictionary/hydrol) Merriam-Webster
2. ̂a b Braun, Charles L.; Sergei N. Smirnov (1993). "Why is water blue?"
(http://www.dartmouth.edu/~etrnsfer/water.htm). J. Chem. Educ. 70 (8): 612. Bibcode:1993JChEd..70..612B(http://adsabs.harvard.edu/abs/1993JChEd..70..612B). doi:10.1021/ed070p612(http://dx.doi.org/10.1021%2Fed070p612).
3. ̂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). Other isotopic compositions melt or boil at slightlydifferent temperatures.
4. ^ United Nations (http://www.un.org/waterforlifedecade/background.html). Un.org (2005-03-22). Retrieved on2011-11-22.
5. ^ Utz, Jeffrey. Re: What percentage of the human body is composed of water?(http://www.madsci.org/posts/archives/2000-05/958588306.An.r.html), The MadSci Network
6. ^ Leigh, pp. 27–28.
7. ^ Leigh, p. 34.
8. ^ 22.064 MPa / ((1 kg × gravity on earth) per liter) = 2.25 km
9. ^ Urey, Harold C. et al. (15 Mar 1935). "Concerning the Taste of Heavy Water". Science 81 (2098) (New York:The Science Press). p. 273. doi:10.1126/science.81.2098.273-a(http://dx.doi.org/10.1126%2Fscience.81.2098.273-a).
10. ^ "Experimenter Drinks 'Heavy Water' at $5,000 a Quart" (http://books.google.com/books?
id=MSoDAAAAMBAJ&pg=PA17). Popular Science Monthly 126 (4) (New York: Popular Science Publishing).Apr 1935. p. 17. Retrieved 7 Jan 2011.
11. ^ Müller, Grover C. (June 1937). "Is 'Heavy Water' the Fountain of Youth?" (http://books.google.com/books?
id=eiYDAAAAMBAJ&pg=PA22). Popular Science Monthly 130 (6) (New York: Popular Science Publishing).pp. 22–23. Retrieved 7 Jan 2011.
12. ̂a b Miller, Inglis J., Jr.; Mooser, Gregory (Jul 1979). "Taste Responses to Deuterium Oxide". Physiology &
13. ^ Campbell, Neil A.; Brad Williamson; Robin J. Heyden (2006). Biology: Exploring Life(http://www.phschool.com/el_marketing.html). Boston, Massachusetts: Pearson Prentice Hall. ISBN 0-13-250882-6.
14. ^ Smith, Jared D.; Christopher D. Cappa, Kevin R. Wilson, Ronald C. Cohen, Phillip L. Geissler, Richard J.Saykally (2005). "Unified description of temperature-dependent hydrogen bond rearrangements in liquid water"(http://www.lbl.gov/Science-Articles/Archive/sabl/2005/October/Hydrogen-bonds-in-liquid-water.pdf). Proc.
15. ^ "How much hydrogen and oxygen is in water"(http://wiki.answers.com/Q/How_much_hydrogen_and_oxygen_is_in_water). Retrieved 2012-10-12.
16. ^ M. Holz, S. R. Heil, A. Sacco (2000). "Temperature-dependent self-diffusion coefficients of water and sixselected molecular liquids for calibration in accurate 1H NMR PFG Measurements". Phys. Chem. Chem. Phys.
17. ^ Heat of Vaporization of Water vs. Temperature (http://www.xydatasource.com/xy-showdatasetpage.php?datasetcode=35484&dsid=111&searchtext=water). Xydatasource.com. Retrieved on 2011-11-22.
18. ^ NASA – Oceans of Climate Change (http://www.nasa.gov/multimedia/podcasting/jpl-earth-20090421.html).Nasa.gov (2009-04-22). Retrieved on 2011-11-22.
19. ^ Constant pressure heat capacity of water vs. temperature (http://www.xydatasource.com/xy-showdatasetpage.php?datasetcode=6841&dsid=104&searchtext=water). Xydatasource.com. Retrieved on2011-11-22.
20. ^ Lide, D. R. (Ed.) (1990). CRC Handbook of Chemistry and Physics (70th Edn.). Boca Raton (FL):CRC Press
21. ^ Water – Density and Specific Weight (http://www.engineeringtoolbox.com/water-density-specific-weight-d_595.html). Engineeringtoolbox.com. Retrieved on 2011-11-22
22. ^ Shell, Scott M.; Debenedetti, Pablo G. and Panagiotopoulos, Athanassios Z. (2002). "Molecular structuralorder and anomalies in liquid silica" (http://www.engr.ucsb.edu/~shell/papers/2002_PRE_silica.pdf). Phys. Rev.
E 66: 011202. arXiv:cond-mat/0203383 (//arxiv.org/abs/cond-mat/0203383). Bibcode:2002PhRvE..66a1202S(http://adsabs.harvard.edu/abs/2002PhRvE..66a1202S). doi:10.1103/PhysRevE.66.011202(http://dx.doi.org/10.1103%2FPhysRevE.66.011202).
25. ^ "IAPWS, Release on the pressure along the melting and the sublimation curves of ordinary water substance,2011" (http://www.iapws.org/relguide/MeltSub.htm). Retrieved 2013-02-19.
26. ^ Thermal Conductivity of some common Materials (http://www.engineeringtoolbox.com/thermal-conductivity-d_429.html). Engineeringtoolbox.com. Retrieved on 2011-11-22
27. ^ The pressure due to water vapor in the air is called the partial pressure (Dalton's law) and it is directlyproportional to the concentration of water molecules in air (Boyle's law).
28. ^ Adiabatic cooling resulting from the ideal gas law.
29. ^ Brown, Theodore L., H. Eugene LeMay, Jr., and Bruce E. Burston (2006). Chemistry: The Central Science.10th ed. Upper Saddle River, NJ: Pearson Education, Inc., ISBN 0131096869.
30. ^ Fine, R.A. and Millero, F.J. (1973). "Compressibility of water as a function of temperature and pressure".
Journal of Chemical Physics 59 (10): 5529. Bibcode:1973JChPh..59.5529F(http://adsabs.harvard.edu/abs/1973JChPh..59.5529F). doi:10.1063/1.1679903(http://dx.doi.org/10.1063%2F1.1679903).
31. ̂a b Nave, R. "Bulk Elastic Properties" (http://hyperphysics.phy-astr.gsu.edu/hbase/hph.html). HyperPhysics.Georgia State University. Retrieved 2007-10-26.
32. ̂a b Schlüter, Oliver (2003-07-28). Impact of High Pressure — Low Temperature Processes on CellularMaterials Related to Foods (http://edocs.tu-berlin.de./diss/2003/schlueter_oliver.pdf) (PDF). TechnischenUniversität Berlin.
33. ^ Zeitler, W.; Ohlhof, T.; Ebner, H. (2000). "Recomputation of the global Mars control-point network"(http://www.asprs.org/a/publications/pers/2000journal/february/2000_feb_155-161.pdf). Photogrammetric
34. ^ Tammann, Gustav Heinrich Johann Apollon (1925). The States Of Aggregation. Constable And CompanyLimited.
35. ^ Lewis, William Cudmore McCullagh and Rice, James (1922). A System of Physical Chemistry. Longmans,Green and co.
36. ^ Crofts, A. (1996). "Lecture 12: Proton Conduction, Stoichiometry"(http://www.life.uiuc.edu/crofts/bioph354/lect12.html). University of Illinois at Urbana-Champaign. Retrieved2009-12-06.
37. ^ Campbell, Mary K. and Farrell, Shawn O. (2007). Biochemistry (http://books.google.com/books?id=NYa45_BxgukC&pg=PA37) (6th ed.). Cengage Learning. pp. 37–38. ISBN 978-0-495-39041-1.
38. ^ Campbell, Neil A. and Reece, Jane B. (2009). Biology (8th ed.). Pearson. p. 47. ISBN 978-0-8053-6844-4.
40. ^ Charlot, G. (2007). Qualitative Inorganic Analysis (http://books.google.com/books?id=Ml-AJ9YbnTIC).Read Books. p. 275. ISBN 1-4067-4789-0.
41. ^ Edme H. Hardy, Astrid Zygar, Manfred D. Zeidler, Manfred Holz, Frank D. Sacher: Isotope effect on thetranslational and rotational motion in liquid water and ammonia. In: J. Chem Phys. 114, 2001, pp. 3174–3181
42. ^ IAPWS (2001). "Guideline on the Use of Fundamental Physical Constants and Basic Constants of Water"(http://www.iapws.org/relguide/fundam.pdf).
43. ^ A Brief History of Temperature Measurement (http://home.comcast.net/~igpl/Temperature.html).Home.comcast.net. Retrieved on 2011-11-22.
44. ^ Mononuclear hydrides (http://www.acdlabs.com/iupac/nomenclature/93/r93_185.htm) in A Guide to IUPACNomenclature of Organic Compounds (Recommendations 1993) online version by ACDLabs
45. ^ Preamble (http://www.acdlabs.com/iupac/nomenclature/93/r93_35.htm) to chemical nomenclature. IUPAC.
46. ^ Leigh, p. 99.
47. ^ "hydron hydroxide compound summary at PubChem"(http://pubchem.ncbi.nlm.nih.gov/summary/summary.cgi?cid=22247451&loc=ec_rcs).
References
Leigh, G. J. et al. (1998). Principles of chemical nomenclature: a guide to IUPAC
recommendations (http://old.iupac.org/publications/books/principles/principles_of_nomenclature.pdf).Blackwell Science Ltd, UK. ISBN 0-86542-685-6.
Release on the IAPWS Industrial Formulation 1997 for the Thermodynamic Properties of Water and
Steam (http://www.iapws.org/relguide/IF97-Rev.pdf) (fast computation speed)Release on the IAPWS Formulation 1995 for the Thermodynamic Properties of Ordinary Water
Substance for General and Scientific Use (http://www.iapws.org/relguide/IAPWS95.pdf) (simplerformulation)Online calculator using the IAPWS Supplementary Release on Properties of Liquid Water at 0.1 MPa,
September 2008 (http://www.staff.uni-bayreuth.de/~bt150361/tools/h2o/h2o_gui.html)Chaplin, Martin. "Water Structure and Science" (http://www.lsbu.ac.uk/water/sitemap.html). London
South Bank University. Retrieved 2009-07-07.Calculation of vapor pressure (http://ddbonline.ddbst.de/AntoineCalculation/AntoineCalculationCGI.exe?
component=Water) of waterWater Density Calculator (http://www.linkingweatherandclimate.com/ocean/waterdensitycalc.php)
Why does ice float in my drink?(http://www.nasa.gov/audience/foreducators/topnav/materials/listbytype/Why_Does_Ice_Float.html),NASA
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