Water's Chemical Properties.
Water's Chemical Properties
You probably know water's chemical description is H2O. As the diagram to the left shows, that is one atom of oxygen bound to two atoms of hydrogen. The hydrogen atoms are "attached" to one side of the oxygen atom, resulting in a water molecule having a positive charge on the side where the hydrogen atoms are and a negative charge on the other side, where the oxygen atom is. Since opposite electrical charges attract, water molecules tend to attract each other, making water kind of "sticky." As the right-side diagram shows, the side with the hydrogen atoms (positive charge) attracts the oxygen side (negative charge) of a different water molecule. (If the water molecule here looks familiar, remember that everyone's favorite mouse is mostly water, too).
All these water molecules attracting each other mean they tend to clump together. This is why water drops are, in fact, drops! If is wasn't for some of Earth's forces, such as gravity, a drop of water would be ball shaped -- a perfect sphere. Even if it doesn't form a perfect sphere on Earth, we should be happy water is sticky.
Water is called the "universal solvent" because it dissolves more substances than any other liquid. This means that wherever water goes, either through the ground or through our bodies, it takes along valuable chemicals, minerals, and nutrients.
Pure water has a neutral pH. Pure water has a pH, of about 7, which is neither acidic nor basic.
Diagram about pH
Water's Physical Properties
* Water is unique in that it is the only natural substance that is found in all three states -- liquid, solid (ice), and gas (steam) -- at the temperatures normally found on Earth. Earth's water is constantly interacting, changing, and in movement.
* Water freezes at 32o Fahrenheit (F) and boils at 212o F (at sea level, but 186.4° at 14,000 feet). In fact, water's freezing and boiling points are the baseline with which temperature is measured: 0o on the Celsius scale is water's freezing point, and 100o is water's boiling point. Water is unusual in that the solid form, ice, is less dense than the liquid form, which is why ice floats.
* Water has a high specific heat index. This means that water can absorb a lot of heat before it begins to get hot. This is why water is valuable to industries and in your car's radiator as a coolant. The high specific heat index of water also helps regulate the rate at which air changes temperature, which is why the temperature change between seasons is gradual rather than sudden, especially near the oceans.
* Water has a very high surface tension. In other words, water is sticky and elastic, and tends to clump together in drops rather than spread out in a thin film. Surface tension is responsible for capillary action, which allows water (and its dissolved substances) to move through the roots of plants and through the tiny blood vessels in our bodies.
* Here's a quick rundown of some of water's properties:
o Weight: 62.416 pounds per cubic foot at 32°F
o Weight: 61.998 pounds per cubic foot at 100°F
o Weight: 8.33 pounds/gallon, 0.036 pounds/cubic inch
o Density: 1 gram per cubic centimeter (cc) at 39.2°F, 0.95865 gram per cc at 212°F
By the way:
1 gallon = 4 quarts = 8 pints = 128 ounces = 231 cubic inches
1 liter = 0.2642 gallons = 1.0568 quart = 61.02 cubic inches
1 million gallons = 3.069 acre-feet = 133,685.64 cubic feet
water
water, odorless, tasteless, transparent liquid that is colorless in small amounts but exhibits a bluish tinge in large quantities. It is the most familiar and abundant liquid on earth. In solid form (ice) and liquid form it covers about 70% of the earth's surface. It is present in varying amounts in the atmosphere. Most of the living tissue of a human being is made up of water; it constitutes about 92% of blood plasma, about 80% of muscle tissue, about 60% of red blood cells, and over half of most other tissues. It is also an important component of the tissues of most other living things.
Chemically, water is a compound of hydrogen and oxygen, having the formula H2O. It is chemically active, reacting with certain metals and metal oxides to form bases, and with certain oxides of nonmetals to form acids. It reacts with certain organic compounds to form a variety of products, e.g., alcohols from alkenes. Because water is a polar compound, it is a good solvent. Although completely pure water is a poor conductor of electricity, it is a much better conductor than most other pure liquids because of its self-ionization, i.e., the ability of two water molecules to react to form a hydroxide ion, OH-, and a hydronium ion, H3O+. Its polarity and ionization are both due to the high dielectric constant of water.
Water has interesting thermal properties. When heated from 0°C, its melting point, to 4°C, it contracts and becomes more dense; most other substances expand and become less dense when heated. Conversely, when water is cooled in this temperature range, it expands. It expands greatly as it freezes; as a consequence, ice is less dense than water and floats on it. Because of hydrogen bonding between water molecules, the latent heats of fusion and of evaporation and the heat capacity of water are all unusually high. For these reasons, water serves both as a heat-transfer medium (e.g., ice for ...
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Water has interesting thermal properties. When heated from 0°C, its melting point, to 4°C, it contracts and becomes more dense; most other substances expand and become less dense when heated. Conversely, when water is cooled in this temperature range, it expands. It expands greatly as it freezes; as a consequence, ice is less dense than water and floats on it. Because of hydrogen bonding between water molecules, the latent heats of fusion and of evaporation and the heat capacity of water are all unusually high. For these reasons, water serves both as a heat-transfer medium (e.g., ice for cooling and steam for heating) and as a temperature regulator (the water in lakes and oceans helps regulate the climate).
Structure of the Water Molecule
Many of the physical and chemical properties of water are due to its structure. The atoms in the water molecule are arranged with the two H-O bonds at an angle of about 105° rather than on directly opposite sides of the oxygen atom. The asymmetrical shape of the molecule arises from a tendency of the four electron pairs in the valence shell of oxygen to arrange themselves symmetrically at the vertices of a tetrahedron around the oxygen nucleus. The two pairs associated with covalent bonds (see chemical bond) holding the hydrogen atoms are drawn together slightly, resulting in the angle of 105° between these bonds. This arrangement results in a polar molecule, since there is a net negative charge toward the oxygen end (the apex) of the V-shaped molecule and a net positive charge at the hydrogen end. The electric dipole gives rise to attractions between neighboring opposite ends of water molecules, with each oxygen being able to attract two nearby hydrogen atoms of two other water molecules. Such hydrogen bonding, as it is called, has also been observed in other hydrogen compounds. Although considerably weaker than the covalent bonds holding the water molecule together, hydrogen bonding is strong enough to keep water liquid at ordinary temperatures; its low molecular weight would normally tend to make it a gas at such temperatures.
Various other properties of water, such as its high specific heat, are due to these hydrogen bonds. As the temperature of water is lowered, clusters of molecules form through hydrogen bonding, with each molecule being linked to others by up to four hydrogen bonds, each oxygen atom tending to surround itself with four hydrogen atoms in a tetrahedral arrangement. Hexagonal rings of oxygen atoms are formed in this way, with alternate atoms in either a higher or lower plane than their neighbors to create a kinked three-dimensional structure.
Liquid Water
According to present theories, water in the liquid form contains three different molecule populations. At the highest temperatures single molecules are the rule, with little hydrogen bonding because of the high thermal energy of the molecules. In the middle range of temperatures there is more hydrogen bonding, and clusters of molecules are formed. At lower temperatures aggregates of clusters also form, these aggregates being the most common arrangement below about 15°C. On the basis of these three population types and the transitions between them, many aspects of the anomalous behavior of water can be explained. For example, the tendency of water to freeze faster if it has been cooled rapidly from a relatively warm temperature than if it has been cooled at the same rate from a lower temperature is explained in terms of the greater number of irregularly shaped cluster aggregates in the cooler water that must find a suitable means of fitting together with a neighboring aggregate.
The discovery in the late 1960s of "superwater," or "polywater," helped to shed light on some aspects of the structure of water. This substance was thought by some to be a giant polymer of water molecules, 40 times denser and 15 times more viscous than ordinary water. Studies showed, however, that these new and unexplained properties were connected with the presence of contaminants in the water. Even so, the interaction of the water molecules with these other substances may be helpful in understanding the way in which water molecules interact with each other.
Ice
In ice, each molecule forms the maximum number of hydrogen bonds, resulting in crystals composed of open, hexagonal columns. Because these crystals have a number of open regions and pockets, normal ice is less dense than water. However, other forms of ice also exist at conditions of higher pressure, each of these different forms (designated ice II, ice III, etc.) having greater density and other distinct physical properties that differ from those of normal ice, or ice I. As many as eight different forms of ice have been distinguished in this manner. The higher pressures creating such forms cause rearrangements of the hexagonal columns in ice, although the basic kinked hexagonal ring is common to all forms.
When ice melts, it is thought that the fragments of these structures fill many of the gaps that existed in the crystal lattice, making water denser than ice. This tendency is the dominant one between 0°C and 4°C, at which temperature water reaches its maximum density. Above this temperature, expansion due to the increased thermal energy of the molecules is the dominant factor, with a consequent decrease in density.
Carbon Dioxide is present in water in the form of a dissolved gas. Surface waters normally contain less than 10 ppm free carbon dioxide, while some ground waters may easily exceed that concentration. Carbon dioxide is readily soluble in water. Over the ordinary temperature range (0-30 C) the solubility is about 200 times that of oxygen. Calcium and magnesium combine with carbon dioxide to form carbonates and bicarbonates.
Aquatic plant life depends upon carbon dioxide and bicarbonates in water for growth. Microscopic plant life suspended in the water, phytoplankton, as well as large rooted plants, utilize carbon dioxide in the photosynthesis of plant materials; starches, sugars, oils, proteins. The carbon in all these materials comes from the carbon dioxide in water.
When the oxygen concentration in waters containing organic matter is reduced, the carbon dioxide concentration rises. The rise in carbon dioxide makes it more difficult for fish to use the limited amount of oxygen present. To take on fresh oxygen, fish must first discharge the carbon dioxide in their blood streams and this is a much slower process when there are high concentration of carbon dioxide in the water itself.
The Unique Structure of Water
Polarity of water molecules results in hydrogen bonding. The water molecule is relatively simple in structure. Two hydrogen atoms are joined to a single oxygen atom by single covalent bonds.
Oxygen is more electronegative than the hydrogen atoms which allows the electrons of the polar bonds to spend more time closer to the oxygen side of the molecule. The oxygen side becomes more negative in charge, and the hydrogen atoms have a slight positive charge. This forms the polar molecule.
The water molecule is shaped like an isosceles triangle, with a slight bond angle of 104.5 degrees at the oxygen nucleus. The weak Coulombic characteristics of the bonding of hydrogen atoms to the weakly electronegative oxygen atom result in both ionized and covalent states that simultaneously maintain the integrity of water. Water is one of the only compounds that possess these characteristics.
An electrostatic attraction occurs between the polar water molecules. The slight positive charged hydrogen atom is attracted to the slight negative charged oxygen atom of another water molecule. This weak attraction is called a hydrogen bond. Every water molecule is hydrogen bonded to its four nearest neighbors.
Simple exercise to demonstrate the polar nature of water:
. Fill a burette with tap water attached to a ring stand over a 400 ml beaker.
2. Rub an air filled balloon against a wool cloth.
3. Open the valve on the burette to allow a stream of water to flow into the beaker below.
4. Position the balloon near the stream of water.
5. Students should share their observations.
Cohesion of Water Molecules
When water is in liquid form, its weak hydrogen bonds are about one-twentieth as strong as a covalent bond. Hydrogen bonds constantly form and break. Each hydrogen bond lasts for a fraction of a second, but the molecules continuously form new bonds with other water molecules around them. At any time a large percentage of water molecules are bonded to neighboring water molecules which gives water more structure than most other liquids. Collectively, the hydrogen bonds hold water together by the property of cohesion.
Cohesion due to hydrogen bonding contributes to the formation of waves and other water movements that occur in lakes. Water movements are integral components of the lake system and play an important role in the distribution of temperature, dissolved gases, and nutrients. These movements also determine the distribution of microorganisms and plankton.
Related to cohesion is surface tension, a measure of how difficult it is to stretch or break the surface of a liquid. Water has a greater surface tension than all other liquids except mercury. At the interface between water and air is an ordered arrangement of water molecules which are hydrogen bonded to one another and the water below. The result is an interface surface or film under tension. Students can observe the surface tension of water by overfilling a glass of water to the point where water stands above the rim.
The air-water interface forms a special habitat for organisms adapted to living in this surface film. This community is called the neuston. Water's high surface tension serves as a supporting surface for many organisms. Many aquatic organisms have evolved adaptations that allow them to spread their body weight over a large surface area to prevent breaking water's surface tension.
Water's Specific Heat
Water has a high heat capacity. Specific heat a measure of heat capacity, is the heat required to raise the temperature of 1 gram of water 1°C. Water, with its high heat capacity, therefore, changes temperature more slowly than other compounds that gain or lose energy.
The heat capacity of water stems directly from its hydrogen bonded structure. Although hydrogen bonds are weak, their combined effect is enormous. As heat is added to ice or liquid water, the energy first breaks hydrogen bonds, which allows the molecules to move freely. Since temperature is a measure of the average kinetic energy of molecules (the rate at which they move), the temperature of water rises slowly with the addition of heat. When the temperature of water drops slightly, many additional hydrogen bonds form and release a considerable amount of energy in the form of heat.
This resistance to sudden changes in temperature makes water an excellent habitat because organisms adapted to narrow temperature ranges may die if the temperature fluctuates widely. The heat requiring and heat retaining properties of water provide a much more stable environment than is found in terrestrial situations. Fluctuations in water temperature occur very gradually, and seasonal and diurnal extremes are small in comparison to terrestrial environments.
The high specific heat can have profound effects on climatic conditions of adjacent air masses. When it warms only a few degrees, a large lake can absorb and store a huge amount of heat from the sun in the daytime and summer. At night and during winter, the gradually cooling water can warm the air. This is the reason Michigan and areas east of the Great Lakes have more moderate climates than the Great Lakes region. Mild winters with higher precipitation rates and moist, cool summers are common in Michigan and areas east of the Great Lakes.
Because of water's high specific heat, the water that covers most of the earth's surface keeps temperature fluctuations within limits that allow living organisms to survive. Also, because organisms consist mostly of water, they are more able to resist changes in their own temperatures.
Evaporation and Cooling
Water has a high heat of vaporization - the energy required to convert liquid water to a gas. Because of the energy needed to break the hydrogen bonds holding a water molecule to its neighbors, more energy is required to evaporate liquid water than most other substances. To evaporate each gram of water at room temperature, about 580 calories of heat are needed, which is nearly double the amount needed to vaporize a gram of alcohol or ammonia.
Water's high heat of vaporization helps moderate the earth's climate. A considerable amount of energy from the sun is absorbed by lakes during the evaporation of its surface waters. As water evaporates, the remaining surface water cools. This evaporative cooling occurs because the warmest molecules are those with the greatest kinetic energy and are most likely to leave in the gaseous state. Evaporative cooling of water contributes to the stabilization of temperature in lakes.
Water's Liquid Temperature Range
Water remains liquid over a wide temperature range, from 0 - 100°C. Most other substances remain liquid over a narrower range. Since the chemical reactions of metabolism depend on interactions between molecules moving about in liquid water, the limits of life are set by water's freezing and boiling points. This property of water makes possible a wide variety of aquatic habitats. Some fish species survive in temperatures at or near freezing while some bacteria and algae survive in hot springs where the water temperature is near boiling.
Water as the Universal Solvent
Water is a substance that can almost dissolve anything. Salts such as sodium chloride (NaCl), dissolve in water by dissociating as each ion becomes surrounded by the polar water molecules. Shielded by a shell of water molecules, the ions stay in solution because they are no longer affected by attractive forces from other ions.
Frozen Lake Density
Water is one of the few substances that are less dense as a solid than as a liquid. While most substances contract when they solidify, water expands. This property is due to the hydrogen bonding. When water is above 4 °C it behaves like other liquids; it expands as it warms and contracts when it cools. Water starts to freeze when the temperature approaches 0°C and the molecules no longer move vigorously enough to break their hydrogen bonds. As the temperature reaches 0°C the water molecules become locked into a crystalline lattice, and each water molecule is bonded to the maximum of four partners.
When the surface temperature in a lake reaches 0°C, ice forms and floats on top of the lake. The ice becomes an insulating layer on the surface of the lake; it reduces heat loss from the water below and enables life to continue in the lake. When ice absorbs enough heat for its temperature to increase above 0°C, the hydrogen bonds can be broken and allow the water molecules to slip closer together. If ice sank, lakes would be packed from the bottom with ice, and many of them would not be able to thaw out, since the energy from the air and the sunlight does not penetrate very far.
Density Relationships of Water
A lake's physical, chemical, and metabolism dynamics are governed to a very great extent by differences in density. The density of ice is almost ten times lighter than liquid water. Water's density increases to a maximum at 3.98°C. Therefore, warmer waters are always found on top of cooler water in lakes and produce layers of water called strata. This is typical of a lake that is stratified during the summer. In winter the density differences in water cause a reverse stratification where ice floats on top of warmer waters.
The polarity of water
Water has a simple molecular structure. It is composed of one oxygen atom and two hydrogen atoms. Each hydrogen atom is covalently bonded to the oxygen via a shared pair of electrons. Oxygen also has two unshared pairs of electrons. Thus there are 4 pairs of electrons surrounding the oxygen atom, two pairs involved in covalent bonds with hydrogen, and two unshared pairs on the opposite side of the oxygen atom. Oxygen is an "electronegative" or electron "loving" atom compared with hydrogen.
Water is a "polar" molecule, meaning that there is an uneven distribution of electron density. Water has a partial negative charge () near the oxygen atom due the unshared pairs of electrons, and partial positive charges () near the hydrogen atoms.
An electrostatic attraction between the partial positive charge near the hydrogen atoms and the partial negative charge near the oxygen results in the formation of a hydrogen bond as shown in the illustration.
The ability of ions and other molecules to dissolve in water is due to polarity. For example, in the illustration below sodium chloride is shown in its crystalline form and dissolved in water.
Many other unique properties of water are due to the hydrogen bonds. For example, ice floats because hydrogen bonds hold water molecules further apart in a solid than in a liquid, where there is one less hydrogen bond per molecule. The unique physical properties, including a high heat of vaporization, strong surface tension, high specific heat, and nearly universal solvent properties of water are also due to hydrogen bonding. The hydrophobic effect, or the exclusion of compounds containing carbon and hydrogen (non-polar compounds) is another unique property of water caused by the hydrogen bonds. The hydrophobic effect is particularly important in the formation of cell membranes. The best description is to say that water "squeezes" non-polar molecules together.
