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nPhysical properties-the result of hydrogen bonding
Studied in isolation, the water molecule is deceptively simple. Its tow hydrogen atoms are joined to the oxygen atom by single covalent bonds. Because oxygen is more electronegative than hydrogen, the electrons of the polar bonds spend more time closer to the oxygen atom. In other words, the bonds that hold together the atoms in a water molecule are polar covalent bonds, with the oxygen region of the molecule having a partial negative charge and the hydrogen having a partial positive charge. The water molecule, shaped something like a wide V, is a polar molecule, meaning that opposite ends of the molecule have opposite charges.
The anomalous properties of water arise from attractions between these polar molecules. The attraction is electrical; slightly positive hydrogen of one molecule is attracted to the slightly negative oxygen of a nearby molecule. The two molecules are thus held together by a hydrogen bond (FIGURE 1). Each water molecule can form hydrogen bonds to a maximum f four neighbors, and at any given moment, many of the molecules in a sample of liquid water are linked in this way. The extraordinary qualities of water are emergent properties resulting from the hydrogen bonding that orders molecules into a higher level of structural organization.
I will examine four of water¡¦s properties that contribute to the fitness of Earth as an environment for life water¡¦s cohesive behavior, its ability to stabilize temperature, its expansion upon freezing, and its versatility as a solvent.
1.Cohesion
Water molecules stick to each other as result of hydrogen bonding. When water is in its liquid form, its hydrogen bonds are very fragile, about 1/1 as strong as covalent bond. They form, break, and re-form with great frequency. Each hydrogen bond lasts only a few trillionths of a second, but the molecules are constantly forming new bonds with a succession of partners. Thus, at any instant, a substantial percentage of all the water molecules are bonded to their neighbors, making water more structured than most other liquids. Collectively, the hydrogen bonds holds the substance together, a phenomenon called cohesion.
Cohesion due to hydrogen bonding contributes to the transport of water against gravity in plants. Water reaches the leaves through microscopic vessel that extend upward from the roots (FIGURE ). Water that evaporates from a leaf is replaced by water from the vessels in the veins of the leaf. Hydrogen bonds cause water molecules leaving the veins to tug on molecules father down in the vessel, and the upward pull is transmitted along the vessel all the way down to the roots. Adhesion, the clinging of one substance to another, also plays a role. Adhesion of water to the walls of the vessel helps counter the downward pull of gravity.
Related to cohesion is surface tension (FIGURE ), a measure of how difficult it is to stretch or break the surface of a liquid. Water has a greater surface tension than most other liquids. At the interface between water and air is an ordered arrangement of water molecules, hydrogen-bonded to one another and to the water below. This makes the water heaves as though coated with an invisible film. We can observe the surface tension of water by slightly overfilling a drinking glass¡¦ the water will stand above the rim. Water¡¦s surface tension also allows us to skip rocks on a pond. In a more biological example, some animals can sand, walk, or run on water without breaking the surface.
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(1)High heat capacity
Different substances have different capacities for storing energy absorbed as heat. The heat capacity of a substance is the quantity of heat required to raise the temperature of the substance 1C. Specific heat is the heat capacity of a 1-g sample; that is, it is the quantity of that required to change the temperature of 1 g of the substance 1C. Table 1 gives the specific heats of several familiar substances.
Table 1
Specific Heats of Some Familiar Substances
SubstanceSpecific Heat (cal/g.C)
Aluminum0.16
Copper000
Ethanol0.588
Iron0.107
Lead0.006
Silver0.056
Sulfur0.16
Water1.00
Note that it takes almost ten times as much heat to raise the temperature of 1 g of water I C as to raise the temperature of 1 g of iron the same amount. Conversely, much heat is given off by water for even a small drop in temperature.
We can trace water¡¦s high specific heat, like many of its other properties, to hydrogen bonding. Heat must be absorbed in order to break hydrogen bonds, and heat is released when hydrogen bonds form. A calorie of heat causes a relatively small change in the temperature of water because much of the heat energy is used to disrupt hydrogen bonds before the water molecules can begin moving faster. And when the temperature of water drops slightly, many additional hydrogen bonds form, releasing a considerable amount of energy in the form of heat.
The vast amounts of water on the surface of Earth thus act as a giant heat reservoir to moderate daily temperature variations. You need just consider the extreme temperature changes on the surface of the waterless moon to appreciate this important property of water. The high specific heat of water also tends to stabilize ocean temperatures, creating a favorable environment for marine life. Thus, because of its high specific heat, the water keeps temperature fluctuations on land and in water within limits that permit life. Also, because organisms are made primarily of water, they are more able to resist changes in their own temperatures than if they were made of a liquid with a lower specific heat.
() High heat of vaporization
Heat of vaporization is the quantity of heat a liquid must absorb for 1 g of it to be converted from the liquid to the gaseous state. Compared with most other liquids, water has high heat evaporation. To evaporate each gram of water at room temperature, about 580 cal of heat is needed-nearly double the amount needed to vaporize a gram of alcohol or ammonia. Water¡¦s high heat of vaporization is another emergent property caused by hydrogen bonds, which must be broken before the molecules can make their exodus form the liquid.
This is of enormous importance to us because large amounts of body heat can be dissipated by the evaporation of small amounts of water (perspiration) from the skin. This effect also accounts for the climate-modifying property of lakes and oceans. A large portion of the heat that would otherwise heat up the land is used to vaporize water from the surface of lakes or seas. Thus, in summer it is cooler near a large body of water then in interior land areas.
.Expansion upon freezing
In the liquid and solid states, water molecules are strongly associated through hydrogen bonds. In liquid water, the molecules are associated but close together. When water freezes, its molecules take on a more ordered arrangement with large hexagonal holes (FIGURE 4). This three-dimensional structure extends out for billions and billions of molecules. The holes account for the fact that ice is less dense than liquid water.
The ability of ice to float because of the expansion of water as a result of lower density is an important factor in the fitness of the environment. If ice sank, then eventually all ponds, lakes, and even oceans would freeze solid, making life as know it impossible on Earth. During summer, only the upper few inches of the ocean would thaw. Instead, when a deep body of water cools, the floating ice insulated the liquid water below, preventing it from freezing and allowing life to exist under the frozen surface. This enables fish and other aquatic organisms to survive winter in the temperate zones. If ice were denser than liquid water, it would sink to the bottom as it formed., and even the deeper lakes would freeze solid in winter.
4.Solvent
Water is a great solvent. Wherever water flows through the soil or over tree surfaces, it dissolves and takes along valuable materials. Because of its small size and polar nature, water dissolves many materials, more than any other liquid. Water can fit into small surface faults and between molecules which helps dissolve materials. Materials that are ionic or polar can be pulled into water and surrounded by a shell of many water molecules hiding or covering any charge. Many acids, bases and salts ionize easily in a water solution and are immediately surrounded by a hydration layer or shell.
A hydration shell of water surrounding polar or charged materials makes these materials behave as if they were a larger compound. Some relatively large (at the molecular scale), but highly charged materials like clay colloids can be suspended in water. Large molecules with many atoms can be surrounded by water, which minimizes any electrostatic charges and negates any cohesion forces, helping these large molecules dissolve in water. Water is a soft means of dissolving many materials, especially when these materials already have a surface film of water adhering and surrounding them.
However, water is not a universal solvent that can dissolve anything. If it were, it could not be stored in any container, including our cells. According to the theory, substances that are not polar can hardly dissolve in water, oxygen, for example. The fact that oxygen molecules are less abundant and diffuse more slowly in water than in air has been a strong selective force in the evolution of respiratory systems in both terrestrial and aquatic animals.
nChemical properties-the result of dissociation of water molecules
Occasionally, a hydrogen atom shared by two water molecules in a hydrogen bond shifts from one molecule to the other. When this happens, the hydrogen atom leaves its electron behind, and what is actually transferred is a hydrogen ion, a single proton with a charge of +1. The water molecule that lost a proton is now a hydroxide ion, which has a charge of ¡V1. the proton binds to the other water molecule, making that molecule a hydronium ion. We can picture the chemical reaction this way
Although the dissociation of water is reversible and statistically rare, it is exceedingly important in the chemistry of life. Hydrogen and hydroxide ions are very reactive. Changes in their concentrations can drastically affect a cell¡¦s proteins and other complex molecules. The concentration of the two ions are equal in pure water, but adding certain kinds of solutes, called acids and bases, disrupts this balance. Biologists use something called the pH scale to describe how acidic or basic a solution is, which is of great importance to our organisms
1.pH
In any solution, the product of the H+ and OH- concentrations is constant at 10-14.The pH of a solution is defined as the negative logarithm of the hydrogen ion concentration. For a neutral solution, pH is 7. A pH value less than 7 denotes an acidic solution; a pH value above 7 denotes an basic solution. Most biological fluids are within the range from 6to 8.There are few exceptions, however, including the strongly acidic digestive juice of the human stomach, which has a pH of about .
A stable pH is very important to living things. Even a slight change in Ph can harmful because the chemical processes of the cell are very sensitive to the concentrations of hydrogen and hydroxide ions, which is partly because the enzymes can only function well in a fixed pH environment.
.Buffers
Buffers are substances that minimize changes in the concentrations of H+ and OH- in a solution. For example, buffers normally maintain the pH of human blood every close to 7.4. A person cannot survive for more than a few minutes if the blood pH drops to 7 or rises to 7.8.Under normal circumstances, the buffering capacity of the blood prevents such swings in pH.
A buffer works by accepting hydrogen ions from the solution when they are in excess and donating hydrogen ions to the solution when they have been depleted. Most buffer solutions contain a weak acid and its corresponding base, which combine reversibly with hydrogen ions. One of the buffers that contribute to pH stability in human blood and many other biological solutions is carbonic acid. Look at the equation below
The chemical equilibrium between carbonic acid and bicarbonate acts as a pH regulator, the reaction shifting left to right as other processes in the solution add or remove hydrogen ions. If the H+ concentration in blood begins to fall, more carbonic acid dissociated, replenishing hydrogen ions. But when H+ concentration in blood begins to rise, the bicarbonate ion acts as a base an removes the excess hydrogen ions from the solution. Thus, the carbonic acid-bicarbonate buffering system consists of an acid and a base in equilibrium with each other.. Most other buffers are also acid-base pairs..
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