Why does water bead up like this on certain surfaces, like a waxed car or glass? Why doesn't it wet the entire surface evenly? The answer is that the forces of attraction between water molecules — intermolecular forces — are stronger than the attractive forces between the water molecules and the surface.
Intermolecular forces are the glue that hold many materials together. They give many substances their properties, such as melting and boiling temperatures.
These bonds are on the order of 100 times stronger than the bonds that result from the attractive forces we'll discuss in this section. Nevertheless, differences in the forces between atoms and molecules can lead to profound differences even between compounds that would seem similar.
All of the intermolecular forces we'll talk about in this section arise from the fundamental arrangement of the electrons in atoms and molecules.
The type and strength of an intermolecular attraction or repulsion depends on how many electrons are present, how they are arranged in bonds, what kinds of bonds are present, and so on.
In this section, we will discuss five fundamental kinds of intermolecular forces. As you work through these, try to keep in mind their similarities more than their differences.
The electrical force between charged particles (atomic or molecular ions, protons or electrons) is one of the four fundamental kinds of forces in the universe (the others are gravity and the strong and weak nuclear forces). We call this the electrostatic force.
We describe these forces using Coulomb's law (→). The Coulomb force is the strongest of the intermolecular forces; it accounts for the ionic bonding of salts, such as NaCl.
For oppositely-charged atoms, the electrostatic force is attractive, but that's an oversimplification of what really happens between atoms.
All atoms are surrounded by negatively-charged electrons, so at very close range, when the electrons come into close contact, they actually begin to repel one-another. This repulsive force is what causes objects not to fall right through objects upon which they sit, eg. a lamp on a table. It's the repulsion between electrons that keep objects from actually "touching".
In the graph below, the black line is the electrostatic force for two oppositely-charged ions. It shows (unrealistic) infinite attraction at small distances, r. The red curve incorporates the added repulsion at short distances, for a more realistic view of the force at work between charges.
Any chemical or physical bond (bond due to IMFs) occurs because its formation reduces the potential energy of the particles involved, and is a balance between attractive forces and electrostatic repulsion between electron clouds.
Think for a moment about an atom like an inert gas atom, say Argon (Ar). Although we know that the electrons aren't arranged in a strictly spherical cloud around the nucleus (p-orbitals and all that), because the atom is free to rotate, the distribution of electrons is more-or-less spherical.
Because an atom is spherically symmetric, we can't identify any one "side" or portion of its electron cloud that is any more or less positive or negative than any other.
Now think of a molecule like hydrogen chloride (HCl). It doesn't have that spherical symmetry. It's a sphere with a "bump".
In fact, because the chlorine atom is quite electronegative, it pulls the single H-atom electron mostly away from the hydrogen, leaving a mostly bare proton. HCl has two different "ends", one clearly more positive than the other.
The H-atom side is more positive than the Cl-atom side (see the picture below).
This situation is sometimes described with this language: An atom is isotropic. It looks the same viewed from any direction. The HCl molecule is anisotropic (not isotropic) because as we approach it from different directions, we get a different view. Space is isotropic. It has no special direction. We only define "up" as north on our globe out of long-standing convention, not because space has any special up or down direction.
A thing that is isotropic looks the same from any direction. It has no special oreientation, like a uniformly-colored sphere.
A thing that is anisotropic has one or more distinguishable orientations. A pencil is anisotropic because it has two distinguishable ends, but (aside from the labeling) it is isotropic with respect to rotation about its long axis.
You can see how the symmetry of a molecule, or we often say the "breaking" of some underlying symmetry, can lead to an asymmetric distribution of the electrons around it.
This makes one part of the molecule "less negative" (and therefore more positive) than another. A molecule that clearly has two sides or ends, one more negative than the other, is said to possess a permanent dipole (two poles, like a battery, + and -), and has interesting properties because of it. The linear molecule HCl is a good example (below).
Non-linear molecules can have permanent dipoles, too. In the top row of the table below are some molecules which are highly symmetric, and because of that symmetry, cannot have permanent dipoles.
We need a way to measure the relative strengths of dipoles, because a dipole can exert a force on another atom or molecule, and the amount of force exerted will depend on the strength or magnitude of the dipole.
The strength of a dipole depends on how much charge difference exists between one "end" of a molecule and another, and the separation of those ends. We call the measure of dipole strength the dipole moment. In mathematics, a moment is a (vague) measure of shape.
← This table lists the dipole moments of some commonly encountered substances. The unit of measurement of dipole moment is the Debye (D).
|Water (H2O)||1.85 D|
|Chloromethane (CH3Cl)||1.8 D|
|Ethanol (C2H5OH)||2.69 D|
|Carbon dioxide (CO2)||0|
The figures on the right show how two or more dipoles interact when they are close. First, we introduce a shorthand notation, the crossed-arrow symbol; the plus goes on the positive side of the dipole and the arrow on the negative side.
When dipoles are free to rotate, they tend to class, as shown. This happens because of the twisting force shown. That twisting force is simply due to the electrostatic force: A region of high negative charge will repel the more negative end of a dipole and (relatively speaking) attract the more positive end.
When many dipoles collect, e.g. in a solid, they tend to class as shown in the bottom panel. Note the each + and - charge is surrounded by opposite charges.
When designating the + and - ends of a molecular dipole, we generally use the symbols δ+ and δ- rather than + and -. We use the Greek letter δ (delta) to indicate that we're not talking about a full positive or negative charge, just a partial charge. δ+ means "a little more positive."
It turns out that the dipole moment is not the only kind of charge moment a molecule can have. CO2 has no dipole moment, but it does have a quadrupole moment. While these higher moments must be considered in some weak interactions, they are of less importance than the dipole.
At a meeting of the Faraday Society in England years ago, a Harvard Professor, responding to a point made in a discussion of weak bonding between molecules said, "Aha! The quadrupole moment ... last refuge of a scoundrel!" Nerdy that I remember that.
We have seen that charges and regions of relative charge difference can interact to attract and/or repel one another. How does a non-charged, non-polar atom or molecule interact with a dipole? It turns out that a dipole can induce a dipole moment in, say, an atom. Play the animation below to see how it works.
The animation simulates the distortion of an electron cloud as a permanent dipole (HCl, in this case)is brought near a neutral atom, the anisotropy of its charge distribution redistributes the electron density of the atom. Electrons are attracted away from the atomic nucleus and toward the positive end of the dipole. An atom is spherically symmetric, so it can't have a permanent dipole, but this distortion disrupts that symmetry and induces a small, temporary dipole moment in it.
This effect is stronger with the strength of the permanent dipole, and with the polarizability of the atom.
The induced dipole interacts with any permanent dipole just as a permanent one would, induced dipoles are just weaker, so any attraction or repulsion is weaker.
Any polar molecule can similarly rearrange the electron cloud of any molecule, including molecules that already have permanent dipole moments. In fact, the induction force accounts for about 10% of the total cohesive binding force of liquid water - the force that accounts for the beading of water on a surface, the first image of this section.
An atom or molecule is polarizable if its electron cloud can be distorted in the presence of an electric field, like the one that might be experienced by being near a permanent dipole.
Polarizability increases with the number of electron shells, so it increases down groups (columns) on the periodic table.
The dispersion force, also known as London dispersion, is a purely quantum-mechanical force, and can only be truly understood by applying the laws (and mathematics) of quantum mechanics to the outer electrons of atoms.
It goes something like this: Think about two inert-gas atoms sitting side by side. From time to time in an atom the electron density is momentarily unbalanced, momentarily spoiling the spherical symmetry and forming a small dipole. This moment can induce a similar fluctuation in the other atom, and so on. Such fluctuations can become correlated (they occur more or less in sync.), producing an overall attraction between atoms that we might otherwise think would have no mechanism of attraction. That's roughly what the dispersion force is.
This animation illustrates how such a correlation of electron-cloud asymmetry might look. It's highly stylized, just a cartoon. A little funny really, atoms doing a jiggle dance.
The dispersion force scales roughly with the number of electrons for atoms, and similarly for molecules—the bigger the atoms the larger the polarizability, and the larger the attractive force due to the dispersion effect.
We know that the dispersion force is real by looking at the properties of inert gases, where the only possible intermolecular force is dispersion.
This graph shows that helium (He), with only two electrons, has a much lower boiling point (the temperature at which atoms are liberated from the liquid) than larger inert gas atoms like Krypton (Kr) and Xenon (Xe). As the dispersion force becomes larger, it takes more energy to liberate an inert gas atom from the liquid because it is stuck to its neighbors more strongly.
Note that the dispersion attraction doesn't scale exactly with the atomic number - the graph is not linear. That means that dispersion must scale with more than just the number of electrons.
Hydrogen bonding is a very important kind of intermolecular interaction in biological chemistry. It is a prevalent force in water and any reaction that takes place in it; it is the force that holds the two strands of the DNA helix together; and it is the most important force determining the structure of protein molecules.
Hydrogen is the most abundant element in the universe. It is ubiquitous in the molecules of life: most are hydrocarbons, and of course water is H2O.
When hydrogen is bonded to an electron-withdrawing element, especially oxygen, it is a good candidate for H-bonding.
The essence of H-bonding is the attraction between an essentially bare proton (because it is attached to an electron-withdrawing atom) and an electron-rich atom, such as oxygen, which usually bonds in such a way as to expose two lone electron pairs.
The diagram below shows a H-bond between two water molecules.
The H-bonds between the bases of DNA are shown in the figure below. AT pairs are held together by two H-bonds and GC pairs by three.
Similar H-bonds hold the bases of RNA together to give RNA molecules their unique structures.
The alpha-helices and beta sheets, as well as other structural elements of proteins are held in place by H-bonds.
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