Water is arguably the most important solvent. Life on Earth arose in and evolved to its present state with water as its most vital ingredient. While there may be other life forms in the universe that don't depend on water, it's very difficult for us to conceive of any.
Of the planets in our solar system and others we have found, none rival Earth for volume of liquid water. We know that there is a bit of water on Mars, and that likely there was more in the long-ago past, so many of our efforts to find life on other planets have focused on Mars.
We also know that Europa, one of Jupiter's moons, harbors a great deal of liquid water underneath its icy surface.
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.
Water (H2O) is a nonlinear covalently-bonded triatomic molecule with some interesting properties that arise from its shape and electronic structure. 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 in this figure.
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 contribute their sole electrons 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 with the hydrogens, 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 relatively negative end and a relatively positive end. We refer to these as "δ +" and "δ-" ("delta-plus" & "delta-minus") ends, where δ is the lower-case Greek letter delta.
This charge separation leads to a fairly large dipole moment for water, μ = 1.8 D. (μ is the Greek letter mu.) The dipole moment is a measure of the degree of separation of charge in a molecule. The linear molecule CO2, O=C=O, has no dipole moment because its charge distribution is symmetric; while each oxygen carries slightly more charge than the central carbon, they are symmetrically located around the carbon, so there's no "+ end" and no "- end".
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 below.
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 depicted by the contours is really a little closer to the oxygen atom than the first lone-pair illustration above would lead one to believe.
Below is 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 oxygen electron density peak has been cut off to provide contrast with the small hydrogen peaks – it's that much larger.
The charge distribution of water lends to the easy formation of hydrogen bonds. These unique attractions occur between hydrogen atoms bound to electron-withdrawing atoms (like O, S) and smaller atoms with exposed lone pairs, most frequently oxygen).
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 open 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 strength of the O-H bond in water.
For this reason, H-bonds are relatively easily broken and reformed, and that leads to some important phenomena, like the ability of DNA strands to separate and re-form in cells.
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.
Water turns to ice, its solid form, when enough energy has been removed that small vibrations and rotations of the molecules can no longer overcome the (relatively low) strength of the hydrogen bonds. Each water molecule is bonded to four others in a near-tetrahedral arrangement (lower right, below). Water ice has a hexagonal structure. Along many directions, six-membered rings of water molecules (upper right in the figure below) can be identified.
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 as atoms and molecules move closer together. 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.
Compared to other small molecules of its size, water has a very large heat capacity, the ability to "soak up" kinetic energy with little reflection of that energy in its temperature. Another way to say this is that it takes much more energy to raise the temperature of a gram of water than it does many other materials. This property has significant consequences for the behavior and properties of gaseous, liquid and solid water which, in turn, have implications for almost every facet of the lives we lead, from body temperature regulation to weather and more.
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. Non-bonded atoms only have three degrees of freedom, just the 3 translations.
Any translation of a water molecule in any direction can be expressed as a linear combination of some of each of the x-, y- and z-translations – a little bit of x, a little bit of y and a little bit of z, where "a little bit" could also mean none.
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.
Lastly, there are the vibrations. Each bond in the water molecule can stretch/compress and bend. This amounts to motions of the individual atoms, constrained by the bonds they share. In any molecule, we can identify 3N normal modes of vibration, where N is the number of atoms (for linear molecules it's 3N-1). A normal mode is a fundamental kind of vibration, just as moving along the x-axis is a fundamental kind of translation or rotating around the axis of symmetry (left in the figure above) is a fundamental kind of rotation.
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.
The symmetric stretching motion, one of the three vibrational normal modes of water, 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. This is true of any rotation or vibration, although "where it is" can be on some translational path.
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 dipole moment (which is zero by symmetry) because the symmetry of the molecule is preserved throughout the vibration.
The vibrational normal modes of water are labeled with the Greek letter nu ( ν ), as ν1, ν2 and ν3
The symmetric stretch of H2O occurs at a frequency of 1.09 × 1014 vibrations per second (Hz), so each vibrational period lasts about one femtosecond (fs).
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˚. Here again, the oxygen moves just a bit to preserve the location of the center of mass of the molecule.
Linear triatomic molecules like CO2, O=C=O also have an symmetric bending vibration. In such a vibration CO2 also develops an alternating dipole moment where it had none in its normal or equilibrium configuration.
Each period of the symmetric bending vibration takes twice as long as a symmetric stretching vibration, or about 2 fs.
The asymmetric stretching vibration of water is the highest-frequency mode of vibration, with a frequency just higher than the symmetric stretch.
In this asymmetric motion, the O-H bonds alter in length and the oxygen atom moves roughly from side to side to compensate, keeping the center of mass in place.
So for water we now have nine degrees of freedom of motion of the atoms:
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.
|Heat of vap.
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.
Let's trace that curve from left to right across five regions. Note that on the bottom axis we're just adding heat to the sample at a steady rate, but what is happening to the temperature is anything but steady.
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.
Because water has nine degrees of freedom of movement, it has a high heat capacity for a triatomic molecule.
Because water is polar and forms hydrogen bonds, it has large latent heats of fusion and vaporization, and its solid form is less dense than its liquid form.
A very important consequence of the fact that water is polar might be summed up by the phrase "like dissolves like." That mnemonic means that polar or charged substances (like CH3OH or HCl) tend to dissolve in polar solvents (like water or ethanol), while non polar substances (like methane and benzene) tend to dissolve in non polar solvents (like carbon tetrachloride, CCl4). Charged ions remain in solution because of the polarity of water molecules, i.e. that they have, relatively speaking, partially negative and partially positive ends. In solution, the positive sides of water molecules tend to align (on average – there is always random tumbling in solution) toward a negative ion, as illustrated:
When a positive ion is in an aqueous solution the positive "ends" (the hydrogens) of water molecules tend to orient away from it, 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:
These arrangements of water molecules are called solvation shells, and the order that they require can persist beyond just the first shell to create 2 or more shells, each losing more organization as it gets farther from the ion and more screened from it.
When we attempt to dissolve very non polar solutes in water, they often just don't dissolve. For example, non polar oils and waxes, which are mostly composed of long chains of -CH3- units, tend to separate from water when we try to mix them. Generally oil is of lower density than water and when left undisturbed, floats on top of the water.
Oils, waxes and other non polar 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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