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Anomalous And Unique Properties of Water

By H.S. Jeon, Ph.D
Mount Kisco, NY

Water is so simple, so beautiful, and so unique!

It is such a simple and small molecule, but it has very unique as well as anomalous chemical and physical properties.  The chemical structure of a single water molecule is shown in Fig. 1.  Two hydrogen atoms bound to one oxygen atom to form a ‘V’ shape with the hydrogen atoms at an angle of 104.5°.  The hydrogen atoms have positive charges, and the oxygen atom on the opposite side has two negative charges. The net interaction between the covalent bond and the attracting and repulsion charges produces the ‘V’ shape of the molecule.

Fig. 1. Shape of A Single Water Molecule (Ball-and-Stick Model).

Think of how much water is important in our daily lives. Different people have different percentages of their bodies made up of water. Babies have the most, being born at about 78%. In adult men, about 60% of their bodies are composed of water whereas old men have less than 50% due to the loose of water as getting older. However, fat tissue does not have as much water as lean tissue. In adult women, fat makes up more of the body than men, so they have about 55% of their bodies made of water. [1] About 90% of our blood is water, which helps blood circulate, transport waste, and control body temperature due to the low viscosity and high heat capacity of water. Each day we need about 2-3 liters of water in order to sustain our health.

Looking at water, you might think that it’s the simplest thing around.  Water is a simple chemical substance with a molecular formula H2O: a water molecule has two hydrogen atoms covalently bonded to a single oxygen atom. Water appears in nature in all three common states of matter, such as gas, solid and liquid.  Pure water is colorless, odorless, and tasteless. But it’s not at all simple—in particular, when water molecules are together. Also it is vital for all life on Earth. Where there is water there is life, and where water is scarce, life is scared!  There is a saying, “You spend money like water!” It can have two different meanings to the people who live in different environments.  It means “waste of money” in Korea due to plenty of water whereas “conserve money” in Israel due to lack of water.

What are the physical and chemical properties of water that make it so unique and necessary for living things? There are many anomalous and unique properties that water should not have according to what we presently know about chemistry and physics. These characteristics strongly point to water as a simple but mysterious molecule.  Therefore, finding and understanding the hidden properties of water give us very interesting subjects.

1. Maximum Density

Density change is not constant as temperature changes. With decreasing temperature water density increases and reaches its maximum at 40C, and then begins to decrease after crossing the critical point of 40C. Moreover the density of crystallized ice water is decreased, meaning that volume of water is expanded as water freezes to ice. Due to the higher density of liquid water than that of ice, ice floats on water surface although ice is solid. This causes an inversion and mixing of water bodies on Earth.  It takes surface oxygen down to the bottom and raises bottom toxic gases to the surface to be neutralized and exhausted. Water is not supposed to be most dense as a liquid at 4o C. All other liquids are most dense when they reach the freezing or solid state. Because of this unusual property, we have circulation of water in lakes or oceans in temperate climates.

Floating ice on water acts as a shielding layer that protects the heat loss of water and the decrease in water temperature. If ice density was larger than that of water, ice would sink to the bottom of lakes, and the lakes in the temperate and arctic climates would be frozen from the bottom up. Thus, living creatures under water can be survived from cold as well as food shortage.

Without water circulation oxygenated surface water would not go to the bottom of lakes to enable life to exist at the bottom so that organic sediments could be biodegraded, bottom toxic gases brought up to the surface and removed, and fish to spawn and feed on bottom-feeding insects.  Without this circulation, there would be no life in our lakes. This circulation is replaced by hurricanes, typhoons, monsoons, and torrential rain in the sub-tropic and tropic zones, where the temperature changes are small.

2. Ultrahigh Melting and Boiling Points

Water has an unusually high melting temperature of 00C instead of -1000C although its molecular weight (MW = 18 g/mole) is smallest among solvents.  Its boiling temperature is 1000C, instead of about -800C.  Graphs of adjacent molecules in the Periodic Table of Elements show a straight-line relationship of melting and boiling points far below 00C.  With using this method, we can thus estimate the theoretical melting and boiling points of water.  According to water’s neighboring molecules in the Periodic Table of Elements, ice should melt somewhere around

-1000C instead of 00C, and should boil at about -800C instead of 1000C.  If melting and boiling points were far below 00C all water would be in the gaseous state and there would be no life on earth!

3. Freeze Point for Water Under Pressure

3.1 Freeze Point Increase

Normally, water becomes ice at 00C and 1atm. It is, however, possible for water to become ice at 200C or for ice to become water at -200C at different pressure conditions. W hen you squeeze ice you lower its freezing temperature slightly so that it melts more easily. This behavior in which pressure can induce melting is extremely unusual. Almost every other material in nature becomes more difficult to melt as you squeeze it. That’s because most materials expand during melting and have to do work against the pressure. Added pressure makes it harder for them to melt and you can even make them freeze by squeezing them. But ice shrinks during melting because ice is less dense than water. Rather than keeping ice from melting, pressure can actually liquefy ice, even below its normal melting temperature.

A related question is “Would it be possible to pass a thin wire through an ice cube without breaking the ice cube?” The answer is “possible”.  Here is an example. A piece of thin wire with a heavy weight on each side secured to the wire is put on the ice cube. Over a period of time the wire would cut through the ice and the ice would refreeze again where it had passed through. Finally the wire would completely pass through the ice. The freeze point of ice under the wire would be reduced as below zero due to the pressure of the wire causing the wire to cut through the ice and then refreeze. It means that the wire is actually passing through the thin water layer due to the melting of ice with a condition that is below the freeze point. Unlike nearly every other substance we know, water takes up more volume in its solid state than in its liquid state. When you put ice under pressure, you decrease the volume available to it, and it cannot remain as a solid. This application of pressure subsequently lowers the freeze point since the crystalline structure of the water is distorted or crushed.

Another example is how ice skates work, although the exact mechanism is somewhat disputed. If you apply pressure on the surface of ice, it melts.  This is how ice skates work – the pressure of the blade causes a tiny bit of ice to be melted, allowing the blade to slide smoothly over it. While pressure-induced melting has long been used to explain the lubrication with a thin layer of water film between ice and skate blades in many texts, that explanation appears to be not good enough or incorrect.

We estimate an upper bound to the temperature change of ~0.1°C by using the Clausius-Clapeyron equation [2], with a skater’s weight of 100kg and a surface area of two skate blades (2LW = 300 mm2),

ΔTm = T0 (ΔP*Δv)/(Hm),

where ΔTm = the change in the melting temperature to the absolute melting, T0

Δv = change in specific volume from ice to water (~10-4 m3/kg)

Hm = Heat of melting of ice (= 3.34×103 J/kg) [3]

ΔP = pressure change (N/m2)

So, if this were the only cause of the slipperiness, ice-skating would be possible only at temperatures just a few degrees below freezing.  From observation, it is possible to ice skate on ice at much lower temperatures than this. It is, thus, necessary to have another chemical or physical phenomenon to explain this melting phenomenon in skating. We can think of frictional heat from interface between ice and blades. Frictional heating due to the movement of the skate contributes to the melting of ice and generates a thin liquid layer of water.

Ice’s huge melting heat (334,000 J/kg) is what keeps an interface or mixture of water and ice at 0 °C. As long as both water and ice are contacting together at the interface, they are in the process of either melting or freezing. Any heat you add to the interface goes into melting more ice, not into raising its temperature. Any heat you remove from the interface comes from freezing more water, not from lowering its temperature. With ice floating in your drink, it will remain at 0 °C, even in the hottest or coldest weather. After the blade passes the thin layer of liquid refreezes as pressure and temperature return to normal states.

On the other hand, it is easy to explain that why the surface of ice is at least a little slippery. At the surface of ice, water molecules are only hydrogen bonded to their neighbors on one side. Consequently, their energy is not as low as in bulk ice. At equilibrium they must have higher entropy in the air-ice interface than bulk ice. So ice must have a thin water layer on the surface, whose thickness would be expected to increase at temperatures close to the melting of ice. Such melting is particularly easy at the interface between ice and air, where the crystalline structure is incomplete and disordered. Because they lack a full complement of neighbors, the outermost water molecules are relatively mobile and already have a liquid-like character.

When heated by sliding friction, this layer melts entirely and acts as a lubricant to make ice extremely smooth. Since the layer is so thin, very little heat is needed to melt it and a tiny bit of frictional heating is all it takes to get something sliding along the ice. The thickness of a thin water layer between ice and blade can be controlled by given temperature and frictional heating.   At warmer temperatures, the water layer is thicker and melts more easily when heated by friction. That’s why ice is most slippery when the temperature is close to freezing and the surface is “wet”. When the temperature is extremely cold, ice has a “dry” surface and is very difficult to slide solid objects on that surface. For example, if you hold an ice cube in your hand on a very cold winter day, it will stick quickly and strongly to your hand.

3.2 Freeze Point Decrease

In 1956, Buswell and Rodebush pointed out two very interesting natural phenomena about water freezing at above a normal freezing point [6]. One of them is that freezing water is sometimes found in a natural gas pipe at 200 C. Natural gas (mostly methane) is a non-electrolyte, which has, thus, a very limited solubility in water. Water is an excellent solvent due to its polarity, high dielectric constant and small size, particularly for polar and ionic compounds and salts.

They explained that some solubility is even possible not due to the attraction force between water and natural gas but due to the lack of attraction force. We can thus understand this unusual freezing by studying the solubility of methane in water. When methane is mixed with water, the reaction generates unexpectedly high heat despite the fact that methane has nearly no interactions with water and thus has a little solubility in it. Mixing of methane in water exerts ten times the heat than that of methane in hexane, which is a very good solvent for methane. It was discovered that the unusual heat generated from the water molecules forms a “cage” around a methane molecule. When methane molecules, which are relatively larger than water molecules, are dissolved in water, they push out many water molecules that break out the hydrogen bonding and reduce the other interaction forces between water molecules. This causes the reduction of strong internal forces of water molecules and an increase of freezing point of water. Water thus freezes at the interface between methane and water.

The other unique phenomenon is corn, which sometimes freezes at 40C. This can also be explained using the same physical chemistry principles as explained in the freezing water in methane gas as previously noted. The same phenomenon can occur between water and protein molecules. All protein molecules are much larger than water molecules and majority of them have non-polar or non-ionic groups of atoms like in methane. Therefore, water molecules can also form aggregates on the surface of protein molecules, and tend to be crystallized on the surface of protein molecules. The corn is thus damaged due to the expansion of water when it freezes to ice at 40C.

4. Very High Heat Capacity

Water has a very high heat capacity (or specific heat) compared to other materials. This means that it is more difficult to raise the temperature of water compared to other substances. Heat capacity is the amount of heat in calories required to raise the temperature of 1g of material 1 0C to raise its temperature.

For example, the heat capacity of water is 1.0cal/g0C while the heat capacity for ice is only 0.5cal/g 0C.

If water is frozen, its specific heat reduced by half, so ice tends to warm easily. If it is liquid, it tends to be more difficult to raise the temperature. Due to the very high differences in the heat capacity of water, ice, and steam, water tends to remain near the most desirable temperature for life on Earth regardless of drastic changes in atmospheric temperatures. The anomalously high heat capacity of water and the right quantity of water stabilize the Earth’s temperature.

4.1 Strange Climates in the Deserts

 

In a desert region, air temperature drops quickly after sunset and remains much lower than that of daytime.  The temperature differences between day and night in desert regions are much bigger than those at a typical area with common humidity. It is all due to the very high heat capacity of water than that of dry sand. At night, water releases the heat absorbed during daytime as the air temperature falls after sunset, whereas dry sand doesn’t hold enough heat to release at night because of its much lower hear capacity.  Deserts without humid air, plants, and lakes also exhibit very cold temperatures at night.  That is the reason why deserts are much colder at night than during daytime. Due to the same reason, hydraulic power plants with large water dams also experience significant temperature change at wintertime.

4.2 Freezing Speed of Hot or Cold Water

In very cold winter days when the temperature falls below 10 0C, we can observe water freezing on glass on contact, such as when we throw water on a car windshield.  The question here is “Which water will freeze first: hot or cold?” Surprisingly, the answer is hot water! Why?  Because hot water evaporates faster than cold water, thus losing its energy and reaching its freezing point more quickly.  It is again due to the different specific heats of water: latent heat (1cal/g 0C), heat of vaporization (540cal/g), and heat of melting (80cal/g).

For example, we need 75cal for heating one gram of water at 250C to 1000C, 540cal for heating one gram of water at 1000C to 1000C steam, and 80cal for heating one gram of ice at 00C to water at 00C. Therefore, 180cal must be released in order to turn one gram of 1000C water into ice. It indicates that the rate of evaporation is proportional to the decreasing of water temperature.

The same phenomenon is at work when we sweat to reduce skin temperature, by utilizing the properties of water as sweat evaporates on our skin. Heat capacities of various materials are given in Table 1 [5,7].

5. Very High Surface Tension

Surface tension is the property of the surface of a liquid that allows it to resist external force. This force gives liquid droplets their perfect spherical shape to minimize surface area if its surface tension is bigger.  For example, a water droplet has a more perfect spherical shape than that of oil in the air because the surface tension of water (73 dyn/cm) is higher than that of acetone (24 dyn/cm) at room temperature [4].  Another example is the ability of some insects (e.g. water striders) to sit and run on the water surface. This property is caused by the cohesion of like molecules, and is responsible for many of the behaviors of liquids. Surface tension has the dimension of force per unit length, or of energy per unit area. Surface tension of water is very large and 2-4 times larger than that of most organic liquids.

The amount of force required to break apart pure water is enormous. The force of the resistance against this force is called tension or tensile strength. The theoretically calculated value of tensile strength is 95,000kg in order to break a pure water cylinder with a diameter of 1 inch and without any structural defects [8]. But there is no such perfect water in the real world. All normal waters in this world have structural defects and impurities, such as ions (H+, OH-), minerals (Mg, K, Na), and also isotopes (H2, H3, O17, O18).  The experimental value of the tensile strength, 454 kg/in2 was obtained in a laboratory setting using the best quality of water [8]. This value can be comparable to some other materials: concrete (~174 kg/in2), glass (~2,130 kg/in2), ASTM A36 steel (~25,800 kg/in2), Diamond (~170,600 kg/in2) [7].

Why Does Water Have Such Anomalous and Unique Properties?

We reviewed some of the unique properties of water, many of which are the result of the “hydrogen bond” between water molecules. The polarity of water allows it to bind with other molecules, including another water molecule. Water molecules form the hydrogen bonds between hydrogen and oxygen molecules, giving shape to water as a liquid at normal conditions. Each single water molecule can form hydrogen bonds with four other water molecules in a tetrahedral arrangement. Although these bonds are weak compared to “covalent bond” they lead to many other unique properties. The V-shape of the water molecule is also important because it allows for other configurations of water to be formed (see Fig. 1).

These anomalous properties also give water the ability to transport minerals and waste products in water bodies, plants and animals.  It gives water the ability to hold oxygen for animal life, and carbon dioxide for plant life.  The unique dipole moment of water establishes the enormous extent of permanent-polarized bonding, and the angle between chemical bonds.  These determine the water’s ability to create the multitude of necessary molecules involved in every life process.  For example [9], intra-molecular hydrogen bonding between hydrogen atoms and oxygen atoms enables molecules to fold into proteins having specific three-dimensional shapes essential for biological activity.  If the angle between hydrogen atoms in the water molecule was not 104.50, there would be no complex life-giving molecules, and no life on Earth.

 

References

  1. Jeffrey Utz, Allegheny University (May 15, 2000), http://www.madsci.org/posts/archives/2000-05/958588306.An.r.html.
  2. Çengel, Yunus A.; Boles, Michael A. (1998). “Thermodynamics – An Engineering Approach”, 3rd ed., Boston, MA.: McGraw-Hill. ISBN 0-07-011927-9.
  3. Lide, D. R., ed (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton (FL): CRC Press. ISBN 0-8493-0486-5.
  4. Yaws, Carl L., ‘”Handbook of Thermodynamic and Physical Properties of Chemical Compounds”, Publisher: Knovel, Electronic ISBN: 978-1-59124-444-8 (Feb 1, 2003).
  5. Laider, Keith, J. (1993). The World of Physical Chemistry. Oxford University Press. ISBN 0-19-855919-4.
  6. A.M. Bushwell and W.H. Rodebush, Scientific American, April 1956.
  7. Wikipedia (Oct. 2011).
  8. Kenneth S. Davis and John A. Day, “Water: The Mirror of Science: Science study series”, Anchor Book, 1st ed. (1961).
  9. Intelligent Design Theory, “Ch.14. Unique Properties of Water (2011)”, www.intelligentdesign.org.

 

Table. 1. Heat Capacities of Various Materials (Unit: cal/g0C) [5,7].

 

Material                         Heat capacity Material                         Heat capacity
Water                            1.00 Lead                               0.04
Ice                                  0.50 Mercury                        0.03
Steam                            0.47 Nitrogen                       0.25
Alcohol                          0.55 Oxygen                          0.22
Aluminum                    0.22 Silver                             0.06
Glass                              0.12 Sand                              0.19
Gold                               0.03 Soil (wet)                     0.35
Hydrogen                     3.41 Turpentine                  0.41
Iron                                0.11 Wood                            0.42

 

Category: Environment

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