The water molecule has two hydrogen atoms held at an angle of 104º 27' and at distances of nearly 0.1 nm from the oxygen atom. The hydrogen atoms are held at about 0.1 nm from each other. The molecule is covalently bonded, electrons being shared between the hydrogens and the oxygen, giving overall electrical neutrality, but the oxygen nucleus has greater affinity for electrons than that of the hydrogen. There is thus, on average, a slight displacement of negative charge towards the oxygen, which leaves a slight positive charge on the hydrogen atoms. Consequently there is an attraction between the slightly positive hydrogen on one molecule, and the slightly negative oxygen on the other, which links them together with a 'hydrogen bond'. The angles at which oxygen and hydrogen are held in the water molecule, coupled with this hydrogen bonding, result in a crystal structure for ice that is based on tetrahedra.
Oxygen is at the centre of each tetrahedron, surrounded by four hydrogen atoms, two covalently and two hydrogen-bonded. Such a crystal is quite open, compared with those of other substances, which often have twelve neighbours (as opposed to four in ice) packed around each molecule, Ice consequently has a low density. It floats on liquid water and, by forming an insulating layer at the surface of water bodies, often prevents them from freezing solid and thus killing fish and other organisms.
The hydrogen bonding in ice is quite strong, because the displacement of negative charge towards the oxygen atom is powerful. The temperature at which melting of ice takes place - a measure of the energy needed to begin breaking down the hydrogen-bonded structure - is thus relatively high, compared with H2S and H2Se, where the charge displacement in the molecules is small.
Water (relative molecular mass = 18) is a liquid at room temperature, in contrast with other substances with small molecules such as methane (r.m.m. = 16), ammonia (17), hydrogen sulphide (34) and carbon dioxide (44) which are gases. In gases the molecules are widely spaced and free to move about independently of each other, whereas in liquids the molecules are close together. In water, the molecules are held close together by hydrogen bonding. It is the hydrogen bonds that cause water to exist as a liquid at the temperatures and pressures that normally prevail on the Earth's surface. Life as we know it depends on this property.
Water has a high heat capacity; that is, a great deal of heat energy is required to raise the temperature of water. This is because much of the energy is used to break the hydrogen bonds which restrict the mobility of the molecules. As a result, water is relatively slow to heat up or cool down. In fact, the specific heat capacity of liquid water is the highest of any known substance.
The latent heat (enthalpy) of fusion of water (the heat energy needed to melt ice) is unusually high. By the same token, relatively large amounts of heat energy must be extracted from liquid water before it freezes. The latent heat of vaporisation of water (the heat energy required to vaporise liquid water) is also unusually high. Thus the evaporation of water requires a great deal of energy and has a remarkable cooling effect.
Most liquids contract on cooling, reaching their maximum density at their freezing point. Water is unusual in reaching its maximum density above its freezing point - at 4¢XC. So when water freezes the ice formed is less dense than water and floats on top. Ice on the surface effectively insulates the water below, thereby making it less likely that the bulk of water (sea, pond or lake) will freeze up even if the air above is very cold.
At the surface of water the molecules are orientated so that most hydrogen bonds point inwards towards other water molecules. This give water a very high surface tension, higher than any other liquid except mercury. Despite this, water molecules slide past each other relatively easily, and water has a remarkably low viscosity.
Compared with other liquids, water has extremely strong adhesive and cohesive properties that prevent it breaking under tension. Water adheres strongly to most surfaces, and can be drawn up through narrow tubes without the water column breaking.
Cohesion and surface tension
Cohesion is the tendency of molecules of one substance to hold together by mutual attraction. The hydrogen bonding of water results in strong cohesive forces. One effect of this is that the surface of a drop of water will assume the smallest possible area, and the drop therefore forms a sphere. The water molecules at the surface are drawn in towards the body of the drop forming a skin-like layer of molecules at the surface. This force is called surface tension. Insects walking on the surface of water and the movement of water up plants are two biological processes that can occur as a result of the cohesive properties of water molecules.
Adhesion and capillarity
Adhesion is the attraction of molecules of different compounds to one another. The ability of water to cling readily to other molecules is responsible for the upward movement of water when a small-bore tube is dipped into it. This phenomenon is called capillarity. Xylem vessels of a diameter 0.02mm can, in theory, support a column of water of height 1.5m by capillarity forces. One of its main biological effects is the upward movement of water in the soil.
Thermal capacity (specific heat)
Another consequence of hydrogen bonding in water is that much heat is needed to cause increased molecular movement and hence gas (steam) formation. The heat energy must first be used to break the hydrogen bonds. For this reason the temperature of water rises only very slowly for a given amount of heat added, when compared with other substances. Similarly it cools more slowly. In all it is thermally stable and so biochemical reactions in a water medium are not subjected to large temperature fluctuations and can take place at a more constant rate. Were it not for hydrogen bonding, water would be a gas at normal environmental temperatures and life as we know it could not exist. For the same reasons much heat is needed to evaporate water and therefore even the evaporation of a small amount of water from the surface of an organism has a large cooling effect, e.g. sweating.
Water has its maximum density at 4ºC. Unlike most other substances it is less dense as a solid than as a liquid. It freezes from the top downwards and the ice that forms at the surface can insulate the warmer water below this layer from the colder temperatures above it. This prevents large bodies of water from freezing solid and has contributed to the survival of aquatic organisms
Dissociation (ionisation), pH and buffers
There is a slight tendency for water molecules to dissociate into ions according to the equation: 2H2O = H3O+ + OH- water oxonium hydroxide molecule ion (hydroxyl) ion It is simpler, however, to consider the dissociation as: H2O = H+ + OH- water oxonium hydroxide molecule ion (hydroxyl) ion In a litre of water this dissociation produces I /10 000 000 (10-7) mole of hydrogen ions. This is equivalent to a pH of 7, which is neutral. If the concentration of hydrogen ions was greater, say 1/1000 (10-3) mole hydrogen ions per litre, the pH would be 3 and the solution would be acidic. Any pH below 7 is acidic, any above is basic. An acid is therefore a substance that donates hydrogen ions and a base is a hydrogen ion acceptor. Note that the pH scale is not linear but logarithmic.
Water has very powerful solvent properties and it dissolves more substances than any other common liquid.
Water as a solvent
Many substances dissolve in water. A substance that dissolves in a liquid is said to be soluble. Sugar, for example, is soluble in water. Water is the solvent and sugar is known as the solute. The solvent and the solute together form a solution. When water is the solvent, the solution is said to be an aqueous solution. Since there is so much water in living organisms, water is a very important solvent. Substances such as glucose dissolve in the water of the blood, which allows it to be carried around the human body. In plants, the most commonly transported substance is sucrose, which is also soluble in water.
There is a limit to the amount of solute that can be dissolved in a solvent. For example, if increasing amounts of sugar are added to water there will come a point at which no more sugar will dissolve in the water. The solution is then said to be saturated. The point at which saturation is reached depends on the temperature of the solvent, because warm water will dissolve more sugar than cold water. There is an important exception to this general rule. Gases can dissolve in liquids. However, gases actually become less soluble as the temperature rises. Water will hold more dissolved oxygen at a low temperature than it will at a high temperature. This has an important effect on the living organisms found in water. Fish such as salmon and trout need a lot of oxygen and are very sensitive to the amount of oxygen dissolve in the water. One particular type of water pollution, thermal pollution, can have a dramatic effect on these fish. Power stations use water from rivers and seas as a coolant and return it to the source as warmer water. However, the warmer water contains less oxygen, so salmon and trout often e when hot water is pumped into their river.
Sometimes it is possible to make more of a substance dissolve in water without raising the temperature. One method of achieving this is to increase the pressure being applied to the liquid. If a force is applied to a liquid, more gas molecules will dissolve in the liquid. Fizzy drinks are an example of this technique in action. The bubbles in drinks are made by the gas carbon dioxide, which is dissolved into the drink under pressure. When the can or bottle is opened, the pressure is released and the excess gas is able to escape in the form of bubbles. As the gas bubbles to the surface, the drink loses its fizz. It will eventually become flat. In some places of the world water comes out of the ground already fizzy! The water in the ground is under pressure and contains dissolved gases, such as carbon dioxide. As the water leaves the ground the pressure is reduced, allowing the bubbles of gas to escape into the atmosphere.
When an active metal such as sodium is placed in contact with liquid water, a violent exothermic (heat-producing) reaction occurs that releases flaming hydrogen gas.
2Na(s) + 2H20(l) = 2Na+(aq) + 2OH-(aq) + H2(g)
This is an example of an oxidation-reduction reaction, which a reaction in which electrons are transferred from one atom to another. In this case, electrons are transferred from sodium atoms (forming Na+ ions) to water molecules to produce hydrogen gas and OH- ions. The other alkali metals give similar reactions with water. Less active metals react slowly with water. For example, iron reacts at a negligible rate with liquid water but reacts much more rapidly with superheated steam to form iron oxide and hydrogen gas.
3Fe(s)+4H2O(g) =w600oC? Fe3O4(s)+4H2(g)
Noble metals such as gold and silver do not react with water at all. ¡@
The combustion of ammonia proceeds with difficulty but yields nitrogen gas and water.
4NH3 + 302 + heat = 2N2 + 6H2O
Ammonia readily dissolves in water with the liberation of heat.
NH3 + H20 NH4 + + OH-
These aqueous solutions of ammonia are basic and are sometimes called solutions of ammonium hydroxide (NH40H). The equilibrium, however, is such that a 1.0 molar solution of NH3 provides only 4.2 millimoles of hydroxide ion. The hydrates NH3¡DH20, 2NH3¡DH20, and NH3¡D2H20 exist and have been shown to consist of ammonia and water molecules linked by intermolecular hydrogen bonds. Liquid ammonia is used extensively as a nonaqueous solvent. The alkali metals as well as the heavier alkaline earth metals and even some inner transition metals dissolve in liquid ammonia, producing blue solutions. Physical measurements, including electrical conductivity studies, provide evidence that this blue colour and electrical current are due to the solvated electron.
Oxyacids and their salts
Nitric acid, HNO3,
was known to the alchemists of the 8th century as "aqua
fortis" (strong water). It is formed by the nitrogen dioxide
(NO2) with water.
When pure, nitic acid is a colourless liquid that boils at 86º C and freezes at -42º C. Upon being exposed to light or heat, it decomposes to produce oxygen, water, and a mixture of nitrogen oxides (primarily NO2).
4HNO3 + light (or heat) = 4NO2 + 2H2O + O2
Consequently, nitric acid is often yellow or brown in colour because of the NO2 that forms as it decomposes. Nitric acid is stable in aqueous solution, and 68% solutions of the acid (i.e, 68 grams of HNO3 per 100 grams of solution) are sold as concentrated HNO3. It is both a strong oxidizing agent and a strong acid. Nonmetallic elements such as carbon (C), iodine (I), phosphorus (P), and sulfur (S) are oxidized by concentrated HNO3 to their oxides or oxyacids with the formation of NO2: e.g.,
S+ 6HNO3 ?H2SO4+ 6HO2 + 2H2O
A water molecule is formed of one H atom and two O atoms linked by covalent bonds. Water is an unusual liquid. It has a very high boiling point and a high heat of vaporization. The maximum density is at 4ºC, and the water expands upon freezing. It has a very high surface tension and is a very good solvent for salts and polar molecules. These properties are a consequence of the dipolar character of H2O. The electron cloud around H2O results from the hybridization of s and p electrons, yielding two bonding orbitals between the O and the two H atoms and two nonbonding sp3 orbitals on the O. The molecule has a high negative charge density near 0 and high positive charge density near H.
In the vapor phase the equivalent size of a water molecule is 3.3 Å, or 0.33 nm (1 Å= 10-10 m). The molecules move at high speed, and their translation energy is so high that during collisions the van der Waals forces are inefficient for creating bonds. The vapor expands when the temperature increases.
In the liquid phase the molecules are close together and occupy a volume of 29.7 Å3, indicating a porosity of 36.7%. Liquid water forms a heterogeneous fluid structure of water molecules, clusters of molecules, and H+ and OH- molecules. The molecular structure influences the density and the viscosity. Liquid water consists of an ice-tridymite structure, a quartzlike structure, and a close-packed ammonia-like structure. In the ice structure each 0 atom is bonded to four others by hydrogen bonds in a tetrahedral configuration. The H atoms in the O-H ... O bond are no longer 0.96 Å from the O atom but may be either 0.99 Å or 1.77 Å away. The volume per H2O molecule in ice is 32.3 Å3. The volume per gram-molecule is 19.56 ml, and the resulting density is 0.92.
The number of broken H bonds increases with temperature up to 50% at 40ºC. Broken hydrogen bonds are responsible to a considerable extent for its high dielectric constant (D = 78.55 at 250ºC), contributing to tire fact that water is the best solvent for polar compounds
The properties of water
Water undergoes various types of chemical reactions. One of the most important chemical properties of water is its ability to behave both as an acid (a proton donor) and a base (a proton acceptor), the characteristic property of amphoteric substances. This behaviour is most clearly seen in the autoionization of water:
H20(l) + H20(l) H30+(aq) + OH-(aq),
where the (l) represents the liquid state, the (aq) indicates that the species are dissolved in water, and the double arrows indicate that the reaction can occur in either direction and an equilibrium condition exists. At 25ºC the concentration of hydrated H+ (i.e., H3O+ , known as the hydronium ion) in water is 1.0 x 10-7 M, where M represents motes per litre. Since one OH- ion is produced for each H30+ ion, the concentration of OH- at 25ºC is also 1.0 x 10-7 M. In water at 25ºX £n£b£` must always be 1.0 x 10-14:
[H+][OH-] = 1.0 x 10-14,
where [H+] represents the concentration of hydrated H+ ions in motes per litre and [OH-] represents the concentration of OH- ions in motes per litre. When an acid (a substance that can produce H+ ions) is dissolved in water, both the acid and the water contribute H+ ions to the solution. This leads to a situation where the H+ concentration is greater than 1.0 x 10-7 M. Since it must always be true that [H+][OH-] = 1.0 x 10-14 at 25ºC, the [OH-] must be lowered to some value below 1.0 x 10-7. The mechanism for reducing the concentration of OH- involves the reaction
H+ + OH-?H20,
which occurs to the extent needed to restore the product of [H+] and [OH-] to 1.0 x 10-14 M. Thus, when an acid is added to water, the resulting solution contains more H+ than OH-; that is, [H+] > [OH-]. Such a solution (in which [H+] > [OH-]) is said to be acidic. The most common method for specifying the acidity of a solution is its pH, which is defined in terms of the hydrogen ion concentration: pH = -log [H+], where the symbol tog stands for a base-10 logarithm. In pure water, in which [H+] = 1.0 x 10-7 M, the pH = 7.0. For an acidic solution, the pH is less than 7. When a base (a substance that behaves as a proton acceptor) is dissolved in water, the H' concentration is decreased so that [OH-] > [H+]. A basic solution is characterized by having a pH > 7. In summary, in aqueous solutions at 25ºC: ¡@
[H+] = [OH-]
pH = 7
[H+] > [OH-]
pH < 7
[OH-] > [H+]
pH > 7
Density and Elasticity
The density of water at 4ºC is I g per ml. Table I gives density values at different temperatures. Water is assumed to be an incompressible fluid. Nevertheless it has a modulus of elasticity of about 300,000 psi, meaning a volumetric decrease of about 0.000048 for each added atmosphere of pressure.
The viscosity of a fluid is the proportionality factor in the expression for the intensity of viscous shear at a point in the moving fluid:
where £n is the shear per
unit area of surface normal to the s-direction, dv /ds is the
maximum velocity gradient at the point, with the s-direction
representing the direction in which the maximum occurs, v is the
kinematic viscosity (=£g/p), and £g is the absolute viscosity
(force ¡D time)/length ².
The unit of viscosity is the poise (dyne = s/cm²).
The viscosity of pure water at atmospheric pressure, as a function of the temperature, is presented in Table 2. The intensity of viscous shear corresponds to the internal energy loss. The velocity gradient and the shear intensity are important in flocculation, settling, and filtration processes.
Vapor Pressure and Relative Humidity
The vapor pressure of a liquid is the pressure of the liquid vapor in contact with the liquid at which vapor molecules condense as fast as they evaporate from it. Vapor pressure is a function of temperature.
Water molecules are held together by attractive forces. Beyond a certain radius, Rcritical, the attractive forces become negligible. Molecules closer than Rcritical, to a free surface are attracted to the interior of the liquid by the resultant force. The potential energy per unit surface area is the surface energy. The numerical value of the surface energy is equal to the surface tension of the liquid. The surface tension decreases with increasing temperature. The interfacial tension between water and another liquid that is immiscible with water is approximately equal to the difference between their surface tensions. Gibbs' rule shows that the addition of a solute to a solvent leads to different behaviors, depending on the surface tension. If the solute at a low concentration has a weak surface tension it will be concentrated at the surface of the solvent and lower the surface tension of the solution. On the contrary, large amounts of a solute of high surface tension will concentrate away from the surface and will not increase the surface tension of the solution. This phenomenon is of great interest in the treatment of surface water and wastewater.
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C0126220(ThinkQuest 2001). All rights reserved.