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Image source: NASA (modified)
Water is important in many other ways in which we are utterly dependent. Our food sources, plant an animal, need it to grow. We exploit the unique properties of water to heat and cool, and to generate power. We ship heavy loads over water and we recreate in it. The properties of water have profound consequences on all kinds of chemistry (aqueous chemistry). In this section, I'll try to build a foundation so you can have some insight into why water behaves as it does.
The basic geometry of the water molecule gives rise to big differences in its properties when compared with other triatomic molecules.
The average structure of the water molecule is shown on the right.
It's important to remember that any such structure is an average over a constantly translating, rotating and vibrating molecule. Molecules never stop (cannot stop) moving. The two most important features of water are (1) that the molecule is bent, with the angle between the Hydrogen atoms approximately 105˚ (as you will see below, the consequences of this fact are immense), and (2) it is composed of a highly electronegative (electron-withdrawing) atom and two H atoms, which are not electronegative and bring only one electron to the bond.
Oxygen has six outer-shell electrons, two of which it shares in molecular orbitals with the H atoms so that it has a full octet. The other two reside in p-like orbitals in a roughly tetrahedral configuration, as VSEPR theory would predict.
These lone pairs, the fact that water is a bent molecule, and the fact that oxygen withdraws some of the electron density from the hydrogens, make water quite polar. Water has a negative end and a positive end. We refer to these as "δ +" and "δ-" ("delta-plus" & "delta-minus") ends.
The importance of the polarity of water cannot be overstated. It leads to all kinds of effects that determine the outcome of reactions and the structure of large molecules and assemblies of large molecules.
We ought to pause here to note that the lone-pair p-like orbital picture above is a bit idealized. When calculations are performed to map the electron density of the H2O electrons not involved in bonding, we really get a picture more like the one on the right.
This contour map is meant to show that there are two regions of high electron density (the two small closed curves), and that the density decreases as we move away. Think of this map just like a contour map of mountains. The outer rings represent the lowest electron density, the inner ones the "mountain tops".
The region of negative charge is really a little closer to the oxygen atom than the first lone-pair illustration above would lead one to believe.
Here's another view of the electron density of water. It shows that most of the electrons are drawn close to the electronegative oxygen atom, leaving very little to the hydrogens. The hydrogens are nearly "bare" in water.
This is partly what leads water to be so polar: "unshielded" protons on one side, and a region of negative charge on the other. The molecule is electrically neutral, but the distribution of its charges is uneven. Think about a battery. It's charge-neutral because it has a negative end and a positive end, but it has two poles, + and -.
The charge distribution of water lends to the easy formation of hydrogen bonds. Hydrogen bonding among water molecules and between liquid water and solutes has profound implications on all kinds of chemistry. The near tetrahedral geometry of that charge distribution also gives solid water (ice) a beautiful hexagonal crystal structure.
In certain experiments, two water molecules, bound together by a hydrogen bond, have been observed unperturbed by any surrounding molecules. The structure of the water dimer looks something like this →
The distance between oxygen atoms is about 3.5Å (1Å = 10-10 m). A hydrogen bond is weak compared to a covalent bond: The H-bond energy is less than 1/10 of the energy of the O-H bond in water.
The water molecule that contributes its proton to the H-bond is called the H-bond donor, the other the acceptor.
In the liquid phase, water molecules possess enough kinetic energy (the energy of movement) to easily break H-bonds just by their constant motion. But the closeness of the molecules in the liquid, and the fact that each water can form four H-bonds at once (two as donor, two as acceptor), makes re-formation of H-bonds very likely.
In the liquid, H-bonds between water molecules are constantly forming and re-forming — they are transient. On average, each water has at least one H-bond to another at any given instant.
As we remove energy from water, the molecules slow down. They translate and rotate much more slowly. At temperatures close to room temp., they are already vibrating about as slowly as they can.
Most substances can exist in gas, liquid and solid forms. When most substances fuse into their solid forms, they become more dense than the liquid. Solid iron sinks in liquid iron; solid CO2 (dry ice) sinks in liquid CO2, and so on, because the solid is more dense than the liquid. But when water freezes it becomes less dense than its liquid. As you are aware, ice floats on liquid water. That is so because formation of the extensive H-bonding network creates a lot of open space in the ordered crystal lattice of ice. In fact, the maximum density of water occurs at about 4˚C.
To understand water in this way, we need first to recall that temperature is essentially the translational movement of atoms and molecules. Translation is movement of the molecule as a whole along one or a combination of the three-dimensional directions, x, y or z. Let's look at the ways that kinetic energy (the energy of movement) can be expressed in atoms and diatomic molecules before returning to water.
For an atom to possess more or less kinetic energy simply means that it moves faster. Its movement (translation) is described by its x, y and z coordinates. Any motion at all in any direction can always be described as a combination of various amounts of movement in the x, y and z directions. Here we consider the x, y and z directions to be basis vectors that can be added in the necessary proportions to produce any translation vector.
Now think about a diatomic molecule like chlorine, Cl2. The picture is complicated because the bond (which wasn't there in an atom) can stretch and compress — the molecule can vibrate, and the molecule can rotate (atoms don't rotate in any classical sense of the word, and in any case wouldn't have much kinetic energy of rotation because of the smallness of the nucleus and the lightness of the electrons). These motions are shown below.
Now let's look at water. It has the same three translation motions as any other body. The animation shows translations along the x (left-right), y (up-down) and z (in-out) axes.
These three motions are called degrees of freedom of motion. Our diatomic molecule above had a total of five degrees of freedom: three translations (true for any molecule), two kinds of rotation and one kind of vibration. Atoms only have three degrees of freedom, just the 3 translations.
The figure below shows that water has three different kinds of rotational motion around its center of mass. Any arbitrary rotation can be expressed as a little (or none) of the first, a little (or none) of the second and a little (or none) of the third. So far that's a total of six degrees of freedom for water. That's six places to put kinetic energy.
Just as any arbitrary translation can be represented as a sum of different portions of translations along the x, y and z axes, any arbitrary vibrational motion of water can be represented as a sum of the three fundamental kinds of vibrations below. They are called the symmetric stretch, the asymmetric stretch and the symmetric bend. By the way, calculating the location of the three rotational axes or the normal modes of vibration is mathematically a little beyond what one typically encounters in high school ... but it's fun so maybe I'll include it one day.
Play the animation a few times to get a feel for what the symmetric stretching vibration looks like. It is essentially a stretching of the O-H bonds, but notice that the oxygen atom moves a little, too. The center of mass of the molecule stays where it is.
Notice that in water, the dipole moment changes as the molecule vibrates because the regions of opposite charge separate and close a bit. The analogous stretch of CO2 , O=C=O, where the carbon sits still and the oxygens move in and out in unison, does not change its zero dipole moment because the symmetry of the molecule is preserved.
This animation shows the mode of vibration called the asymmetric stretch. The H-atoms alternately compress and stretch their bonds while the oxygen moves to compensate and keep the center of mass stationary.
Linear triatomic molecules like CO2, O=C=O also have an asymmetric stretch. In this vibration CO2 develops an alternating dipole moment where it had none before because of the asymmetry that the motion produces in the electron distribution.
This animation shows the mode of vibration called the symmetric bend. The bond angle H-O-H oscillates about its equilibrium value of about 105˚.
Linear triatomic molecules like CO2, O=C=O also have an symmetric bending vibration. In this vibration CO2 also develops an alternating dipole moment where it had none before.
So for water we now have nine degrees of freedom of motion of the atoms: 3 translations of the entire molecule as a unit, 3 rotations about 3 different axes and 3 vibrational normal modes. These motions provide the basis from which to describe any movement that water can undergo. Movement, of course, is the expression of kinetic energy. All of these motions generally occur simultaneously.
Water has six places to "put" kinetic energy that aren't direct deposits of that energy into molecular translation, which we measure as temperature. That's why it takes a lot of energy input to raise the temperature of water by some number of degrees.
Imagine holding onto a 50 g block of aluminum (about palm sized) over a flame. The heat of the flame would very quickly heat the block to a temperature too hot to hold. Now heat 50g of water in a beaker with the same flame, under the same conditions. Do you think you could hold your finger in the water longer than you could hold on to the Al block? (You could). It takes much more energy to raise the temperature of water than aluminum. Aluminum, a good conductor of heat; it has a low heat capacity. Water, a relatively poor conductor of heat, has a high heat capacity.
Here are a few heat capacities →
Diamond is a perfect crystal of carbon. Its atoms are locked in place and each is bonded to four others. Any vibration of one is very quickly transmitted through the whole crystal - there is no "storage" of heat. Diamond has a small heat capacity.
The heat capacity of liquid water is substantially greater than that of many comparable substances.
Ammonia (NH3)has a higher heat capacity because it has four atoms, thus one more vibrational mode in which to store kinetic energy. What's interesting is how close the heat capacity of water is to that of ammonia.
← Here are some representative heats of vaporization.
The amount of energy required to convert liquid water at its boiling temperature into gaseous water (steam) at its boiling temperature is very high compared to other substances.
Ammonia forms H-bonds, but not the kind of extensive H-bonding network that water does.
The heats of vaporization of most molten metals are quite high. Atoms in many pure metals form an extensive network of bonds to neighbor atoms, thus it takes more energy to pry these atoms away from the bulk and release them into the gas phase.
Now let's look at a schematic diagram of continuous heating of water from a temperature well below the freezing point to one well above the vaporization point. Temperature of a sample of water is plotted vs. heat added at a constant rate. This is an amazing curve and one worth staring at for a while.
As we add heat to ice, its temperature rises. That's no big surprise. It rises at a linear rate that is governed by the heat capacity of ice, something that can be looked up in a table.
What happens at 0˚C is really remarkable. We add heat but the temperature does not increase, even after a substantial amount of heat is added. That's a very strange result. Imagine if you were the first to discover it. People might not believe you!
What is happening, of course, is that ice is undergoing a phase transition: ice → liquid. It turns out that this phase transition alone, without rise in temperature, requires an extra amount of heat that we call the latent heat of fusion, ΔHf. Mathematically, the heat capacity of water at this temperature (and at 100˚C) is infinite.
After the solid melts to liquid, the liquid reflects added heat as a rise in its temperature. The slope of the rise isn't quite the same as that of the solid because the heat capacity of water is a little different that that of ice (see table under Heat).
At 100˚C we observe another phase transition: Liquid → gas (steam). Look at the (relatively speaking) immense amount of heat it takes to convert liquid water to steam. It's huge. That's why sweating is such an efficient cooling mechanism. Sweat requires a tremendous amount of heat energy (taken from the body) in order to evaporate into the gas phase. We call this extra energy the latent heat of vaporization, ΔHv.
Finally, we can heat steam almost arbitrarily. Again, the rise in temperature has a slope that is determined by the heat capacity of steam, which is quite different from those of liquid water and ice. At some very high temperature, of course, the O—H bonds of water will break and there won't be any more water, just H and O.
Here are a couple of pictures (below) of the first solvation shell that tends to form around ions
in solution. Charged species are like water in that they bear a charge imbalance (in the case of ions, they are a charge imbalance). The positive "ends" (the hydrogens) of water molecules tend to orient toward a negative ion, and the lone-pair electron density tends to orient toward a positive ion. Remember that liquid water is a dynamic substance, with each water molecule in a constant tumbling motion. Still, at any one time a picture more or like this is likely to exist:
Oils, waxes and other nonpolar substances are called hydrophobic, which means "water fearing". Hydrophobic substances dissolve in other hydrophobic solvents like CCl4, hexane or benzene. But when a hydrophobic substance is placed in water it tends to associate with the most hydrophobic thing present, usually itself. Thus oil aggregates into droplets, then drops, and finally into a layer completely separate from water, minimizing the unfavorable interaction between hydrophobic molecules and water.
It is instructive to think about detergents. As shown in the diagram below, detergents consist of a charged or polar "head group", usually a sulfate or a phosphate, attached to a long hydrophobic tail, usually just composed of repeating CH3 units: -CH2-CH2-CH2-CH2-CH2-CH2- ... The detergent shown below is one of the most common, sodium dodecyl sulfate, a sulfate head group with 12 CH2 groups in the tail.
Detergents tend to form micelles in water. Micelles are little globs of detergent molecules with the polar heads forming the interface with water molecules, and the hydrophobic tails buried inside the spherical micelle. Everything's happy: The head groups are surrounded by polar water molecules, and water is excluded from the core of the micelle, where hydrophobic tails solvate one another.
Formation of micelles like this is driven by the hydrophobic effect.
Lipid molecules (fats) are similar to detergents in that they consist of a charged or polar head and a hydrophobic tail. These molecules tend to form membranes - like cell membranes - and vesicles, small, spherical closed membrane sacs that are used by cells to transport molecules.
Look at the pictures of the vesicle and membrane section on the right. Notice that the head groups are on the outside where they can be solvated by water molecules and the hydrophobic tails are inside of a region that excludes water.
Again, the hydrophobic effect drives the formation of these structures. They form spontaneously when lipids are placed in aqueous solution.
Note: These membrane drawings are highly idealized. Real cell membranes can contain a variety of different lipid types, and usually membranes are pierced by many kinds of proteins that exist to help transport something into or out of the cell, or to signal something that is going on inside the cell to the "world" outside.
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