This is an updated version of the module Water.
Before we start, get yourself a glass of water. By the time you’ve reached the end, you’ll have a much greater appreciation for this miracle liquid.
Got your glass? Now take a sip and think about all the roles water plays in your life. For one thing, your body can’t function more than a few days without it. You use water to wash yourself, your clothes, and your car. Water puts out fires, cooks our food, makes our soap get sudsy, and hundreds of other things. Water is absolutely essential to our lives on Earth.
Water is so central to our existence that you might be surprised to learn that it’s a rare and unusual substance in the universe. Water is at once so vital and so scarce that exobiologists (scientists looking for life beyond Earth) set their sights on planets where water might exist. Life, it seems, can tough it out in acid, lye, extreme salt, extreme heat, and other conditions that would kill us humans. But it can’t exist without water.
What’s so special about water?
Despite its scarcity across the universe, water is so abundant on Earth that we aren’t always aware of how special it is. For starters, water is the only substance that exists naturally on our planet as a solid (ice and snow), liquid (rivers, lakes, and oceans), and a gas (water in the atmosphere as humidity). As you might recall (or can read about in our module on States of Matter), water molecules are in a different energy state in each phase. The amount of energy required to go from solid to liquid and liquid to gas is related to how water molecules interact with each other. Those interactions are, in turn, related to how the atoms within a water molecule interact with each other.
Our Chemical Bonding: The Nature of the Chemical Bond module discussed how a dipole forms across a water molecule; in the bond between oxygen and hydrogen, the electrons are shared unequally, drawn a bit more to the oxygen. As a result, a partial negative charge (ð-) forms at the oxygen end of the molecule, and a partial positive charge (ð+) forms at each of the hydrogen atom ends (Figure 1).
Since the hydrogen and oxygen atoms in the molecule carry opposite (though partial) charges, nearby water molecules are attracted to each other like tiny little magnets. The electrostatic attraction between the ð+ hydrogen (ð stands for partial charge, a value less than the charge of an electron) and the ð- oxygen in adjacent molecules is called hydrogen bonding (Figure 2).
Hydrogen bonds make water molecules "stick" together. These bonds are relatively weak compared to other types of covalent or ionic bonds. In fact, they are often referred to as an attractive force as opposed to a true bond. Yet, they have a big effect on how water behaves. There are many other compounds that form hydrogen bonds, but the ones between water molecules are particularly strong. Figure 2 shows why. If you look at the central molecule in this figure you see that the oxygen end of the molecule forms hydrogen bonds with two other water molecules; in addition, each hydrogen on the central molecule is attracted to a separate water molecule. As the illustration shows, each water molecule forms attractions with four other water molecules, a network of connections that makes the hydrogen bonding in water particularly strong and lends the substance its many unique properties.
Properties of water that arise from hydrogen bonding
Now it’s time to make use of that glass of water. If you have some ice cubes, drop one in your glass. You’ll notice that it floats. Its ability to bob to the top of the water line means that the ice (water in its solid state) is less dense than liquid water. (To review density and buoyancy, see our Density module) This isn’t a common state of affairs; if you put a chunk of solid wax into a vat of molten wax, it will sink toward the bottom (and possibly melt before it gets there).
To understand what causes ice to float but solid wax to sink, let’s think first about what happens when a liquid turns to a solid (again, the States of Matter module can be a handy review here). In a liquid, the molecules have enough kinetic energy to keep moving around. As molecules come near to each other, they are drawn together by intermolecular forces. At the same time, molecules have enough kinetic energy to break free of those forces and be drawn to other nearby molecules. Thus the liquid flows because intermolecular attractions can be broken and reformed.
A liquid freezes when the kinetic energy is reduced (i.e. the temperature is reduced) enough that the attractive forces between molecules can no longer be broken, and the molecules become locked in a static lattice. For nearly all compounds, the lower energy and lack of movement between molecules means the molecules in a solid are packed together more tightly than the liquid state. This is the case with wax and so solid wax is denser than the liquid and sinks.
In the case of water, though, the shape of the molecule and the strength of the hydrogen bonds affect the arrangement of the molecules. In liquid water, hydrogen bonding pulls molecules closely together. As water freezes, the dipole ends with like charges repel each other, forcing the molecules into a fixed lattice in which they are farther from each other than they are in liquid water (Figure 3). More space between molecules makes the ice less dense than liquid water, and thus it floats.
The universal solvent
Water is sometimes referred to as the “universal solvent,” because it dissolves more compounds than any other liquid known. The polarity of the water molecule allows it to readily dissolve other polar molecules, as well as ions. (See our Solutions, solubility, and colligative properties module for a deeper discussion of dissolution.)
This ability to dissolve substances is one of the properties that makes water vital for life. Most biological molecules, such as DNA, proteins, and vitamins are polar, and important ions such as sodium and potassium are also charged. In order for any of these compounds to carry out functions in the body, they have to be able to circulate in the blood and the fluid within and between cells, all of which are mostly water. Because of its polarity, water is able to dissolve these and other substances, allowing their free movement around the body. A few biomolecules, such as fats and cholesterol, aren’t polar, and don’t dissolve in water – however, the body has developed unique ways to circulate and store these substances.
Water is also able to dissolve gasses such as oxygen, allowing fish, plants, and other aquatic life to access this dissolved oxygen (Figure 4). O2 isn’t a polar molecule; it dissolves because the polar charges in the water molecule induce a dipole in the oxygen, making it soluble and so available to aquatic life. (Learn more about induced-dipole interactions in our Properties of Liquids module.)
Cohesion and surface tension
Let’s return to your water glass. Fill the glass just to the rim and stop. Then, slowly, add a little bit more. You’ll see that you can actually fill the glass a bit past its rim, and the edges of the water will round out against the glass, holding the water in.
Once again, hydrogen bonding is behind this act, resulting in cohesion. Cohesion occurs when molecules of the same kind are attracted to each other. In the case of water, the molecules form strong hydrogen bonds, which hold the substance together. As a result, water is highly cohesive, in fact, it is the most cohesive of all non-metallic liquids.
Cohesion occurs throughout your glass of water, but it’s especially strong at the surface. Molecules there have fewer neighbors (because they have none at the very surface), and so create stronger bonds with the molecules that are near them. The result is called surface tension, or the ability of a substance to resist disruption to its surface. Dip your finger into your water glass and then pull it out. The drop that forms at the end of your fingertip is held together by surface tension.
Surface tension was the misunderstood central player in a raucous debate between Galileo Galilei and his chief rival, Ludovico delle Colombe in 1611. Delle Colombe, a philosopher, was at odds with some of Galileo’s ideas, including his explanation that ice floats on water because it is less dense. So the philosopher challenged Galileo to a debate, which delle Colombe believed would prove his own intellectual superiority.
Delle Colombe championed the (incorrect) idea that ice floats not because of density, but because of its shape, which he saw as broad and flat, as is ice on a lake. To prove the “truth” of his theory, he used ebony wood, which is slightly denser than water, in a demonstration before an audience of curious spectators. He dropped a sphere of the wood into water, and it sank. He then placed a thin wafer of the wood flat on the water’s surface, and it floated. Delle Colombe pronounced himself the winner.
Galileo left frustrated. His observations of the world gave him evidence that his explanation, not delle Colombe’s, was right, but he couldn’t explain the outcome of delle Colombe’s experiment.
Had he known about molecules and dipoles and hydrogen bonds at the time, Galileo certainly would have offered this explanation: When delle Colombe floated the thin ebony disc, he was taking advantage of the cohesive nature of water and the surface tension that arises from it (Figure 5). As the ebony wafer appeared to float on the water, the force exerted by its mass was distributed throughout the surface of the water beneath it. In other words, a single pinpoint-sized area of surface water only had to support the pinpoint-sized piece of ebony just above it. The hydrogen bonds between the water molecules were strong enough to support the weight of the disc. When delle Colombe placed the sphere in the water, however, the pinpoint-sized area that first touched the water bore the weight of the entire sphere, which was more than the water’s surface tension could support. Had Galileo known this at the time, he could have disproved delle Colombe easily – had he simply pushed the wafer through the surface to break the surface tension, the wafer would have sunk.
This same surface tension is what allows leaves to stay at the surface of a lake and dewdrops to adhere to a spider’s web. Even some animals take advantage of this phenomenon – the Basilisk lizard (Figure 6), water striders, and a few other small animals and bugs appear to "walk" on water by taking advantage of the surface tension of water.
Adhesion and capillary action
For your next observation, take another sip of water, and notice the side of the glass. Chances are you’ll see a few drops stuck to it. Gravity is pulling down on these drops, so something else must be keeping them stuck there. That something else is adhesion, the attraction of water to other kinds of molecules; in this case, the molecules that make up the glass. Because of the polarity of the molecule, water exhibits stronger adhesion to those surfaces that have some net electrical charge, and glass is one such surface. But place a drop of water on a non-polar surface, such as a piece of wax paper and you will see it take a different shape than one to which it adheres. On the wax paper, the water droplets take the shape of a true droplet because there is little adhesion and the cohesive forces pull the drop into a sphere. But on glass you will see the droplets flatten and deform a bit as the adhesive forces draw it more to the surface of the glass.
Both cohesion and adhesion (Figure 7) occur with many compounds besides water. Pressure sensitive tapes, for example, stick to surfaces because they are coated with a high viscosity fluid that adheres to the surface to which they are pressed. Generally, you can overcome this adhesive force by pulling, for example – you can easily lift a Post-it® Note from a page. But sometimes the adhesive forces are stronger than the forces holding the surface together – pull tape off of a piece of paper and you remove pieces of the paper with the tape.
Let’s return to our glass of water, and look inside to where the water surface meets the glass. The very edge of the water surface curves upward slightly on the glass. That’s also adhesion – the water is drawn up the surface by adhesion with the glass. If you have a clear plastic straw, you can put one end of it into the water and see that the liquid climbs up the straw a bit, above the surface of the remaining glass of water. It’s actually moving upward against gravity!
What’s happening in your straw is a phenomenon called capillary action (Figure 7). Capillary action occurs in small tubes, where the surface area of the water is small, and the force of adhesion—water’s attraction to the polar glass or other material—overcomes the force of cohesion between those surface molecules.
Another way to see the effects of adhesion and cohesion is to compare the behavior of polar and nonpolar liquids. When you put water in a test tube, adhesion makes the water along the edges move slightly upward and creates a concave meniscus. Liquid mercury, on the other hand, is not polar and therefore not attracted to glass. In a test tube, cohesion at the surface of the mercury is much stronger than adhesion to the glass. The surface tension in the mercury forms a convex meniscus, much the same as the way water forms a slight bulge over the top of your very full glass (Figure 8).
Adhesion and capillary action are among the forces at play that help plants take up water (and dissolved nutrients) in their roots. Capillary action also keeps your eyes from drying out, as saline water flows from tiny ducts in the outer corners of your eyes. With each blink, you spread the water away from the duct, and capillary action brings more fluid to the surface.
If you want to see capillary action at work, put a few drops of red food coloring in your glass of water, and then drop a stalk or two of leafy celery into it. After a day or two, your green celery will be streaked with red.
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