Ice and Water

Submitted by ChemPRIME Staff on Thu, 12/16/2010 - 13:03

The next simplest, and by far the most important, example of hydrogen bonding is that which occurs in H2O. Again there is clear evidence of hydrogen bonding in the structure of the solidA state of matter having a specific shape and volume and in which the particles do not readily change their relative positions.. Figure 1 shows two computer-drawn diagrams of the crystal latticeAn orderly, repeating arrangement of points in 3-D space in which each p;oint has surroundings identical to every other point. A crystal's constituent atoms, molecules, and ions are arranged about each lattice point. of ice. In the model we can clearly see that each O atom is surrounded by four H atomsThe smallest particle of an element that can be involved in chemical combination with another element; an atom consists of protons and neutrons in a tiny, very dense nucleus, surrounded by electrons, which occupy most of its volume. arranged tetrahedrally. Two of these are at a distance of 99 pm and are clearly covalently bonded to the O atom. The other two are at a distance of 177 pm.


Figure 1. Two computer images of the structure of ice. The water molecules have been arranged, so that each oxygen atom is surrounded by four hydrogen atoms in tetrahedral geometry. Two of these atoms are covalently bound to oxygen, while the other two are hydrogen bonding with the oxygen.


They are covalently bonded to other O atoms but are hydrogen bonded to the one in question. The situation is thus:

Image:Covalent and H Bonding in Water.jpg

As in the case of HF, the distance between molecules is abnormally short. The sum of the van der Waals radii of H and O is 260 pm, considerably larger than the observed 177 pm.

The tetrahedral orientation of H atoms around O atoms which results from hydrogen-bond formation has a profound effect on the properties of ice and of liquidA state of matter in which the atomic-scale particles remain close together but are able to change their positions so that the matter takes the shape of its container water. In the space-filling diagram of ice, most of the electronA negatively charged, sub-atomic particle with charge of 1.602 x 10-19 coulombs and mass of9.109 x 1023 kilograms; electrons have both wave and particle properties; electrons occupy most of the volume of an atom but represent only a tiny fraction of an atom's mass. densityThe ratio of the mass of a sample of a material to its volume. of each H and O atom is enclosed by a boundary surface. As you can see, hydrogen bonding causes the H2O molecules to adopt a rather open structure with hexagonal channels running through it. These channels contain an almost perfect vacuum-in them there is a little electron density from the surrounding atoms, but nothing else.

When ice melts, some of the hydrogen bonds are broken and the rigid crystal lattice collapses somewhat. The hexagonal channels become partially filled, and the volume of a given amount of H2O decreases. This is the reason that ice is less dense than water and will float on it. As the temperatureA physical property that indicates whether one object can transfer thermal energy to another object. is raised above 0°C, more hydrogen bonds are broken, more empty space becomes occupied, and the volume continues to decrease. By the time 4°C has been reached, increased molecular velocities allow each H2O molecule to push its neighbors farther away. This counteracts the effect of breaking hydrogen bonds, and the volume of a given amount of H2O begins to increase with temperature.

Most solids expand when they melt, and the corresponding liquids expand continually with increasing temperature, so the behavior of water is rather unusual. It is also extremely important in the environment. When water freezes in small cracks in a rock, the greater volume of the ice can split the rock into smaller pieces. These eventually become able to support plant life, and so water contributes to the formation of fertile soil. The same process happens to roadways, and is the reason for new cracks and potholes seen on roads after a cold winter. The ice bomb experiment, seen below, is perhaps the the most dramatic example of water expanding when frozen.

In the video, water is poured into a cast iron container, which is tightly sealed. The container is then placed in a acetone/dry iceSolid carbon dioxide; called dry because of its tendency to sublime (form a gas without first forming a liquid). At standard pressure carbon dioxide sublimes at -78 °C. sludge, which is at a temperature of -77°C. After a short periodThose elements from a single row of the periodic table. of time, the ice freezes, expands, and causes the cast iron container to explode, blasting off the cover of the acetone/dry ice bath, and spraying the bath itself everywhere. Even though the cast iron container had ⅛ inch thick sides, the pressureForce per unit area; in gases arising from the force exerted by collisions of gas molecules with the wall of the container. of the expanding ice was still able to blow it apart.

Since water has maximum density at 4°C, water at that temperature sinks to the bottom of a deep lake, providing a relatively uniform environment all year around. If ice sank to the bottom, as most freezingThe process of forming a solid from a liquid. liquids would, the surface of a lake would not be insulated from cold winter air. The remaining water would crystallize much more rapidly than it actually does. In a world where ice was denser than water, fish and other aquatic organisms would have to be able to withstand freezing for long periods.

Hydrogen bonding also contributes to the abnormally large quantities of heatEnergy transferred as a result of a temperature difference; a form of energy stored in the movement of atomic-sized particles. that are required to melt, boil, or raise the temperature of a given quantity of water. Heat energyA system's capacity to do work. is required to break hydrogen bonds as well as to make water molecules move faster, and so a given quantity of heat raises the temperature of a gram of water less than for almost any other liquid. Even at 100°C there are still a great many unbroken hydrogen bonds, and almost 4 times as much heat is required to vaporize a mole of water than would be expected if there were no hydrogen bonding. This extra-large energy requirement is the reason that water has a higher boiling point than any of the other hydrides.

The fact that it takes a lot of heat to melt, boil, or increase the temperature of water, makes this liquid ideal for transferring heat from one place to another. Water is used by engineers in automobile radiators, hot-water heating systems, and solar-energy collectors. More significantly, circulation (in the bloodstream) and evaporation (from the skin) of water regulate the temperature of the human body. (You are between 55 and 65 percent water if female and between 65 and 75 percent water if male.) Because of this (as well as for many other reasons) water is an important component of living systems. Water’s ability to store heat energy is also a major factor affecting world climate. Persons who live near large lakes or oceans experience smaller fluctuations in temperature between winter and summer than those who inhabit places like Siberia, thousands of kilometers from a sizable body of water. Ocean currents, such as the Gulf Stream, convey heat from the tropics to areas which otherwise would be quite cold. It is interesting to ask, for example, whether European civilization could have developed without the aid of warmth transported by the common, but highly unusual, liquid—water.