On several occasions up to this point I have used the term 'intermolecular attractive forces'. I first used it when we looked at the animation of the container of gas as it was cooled. (See the phase transition animation at the particulate level.) Recall in this animation a container of gas particles was cooled. As the temperature of a collection of particles was lower we observed the particles slowing down. At the lower velocities colliding particles appeared to stick together forming groups of particles. As the temperature continued to drop the number of particles in these groups increased. Eventually the groups of particles are of sufficient size that they fall to the bottom of the container as a result of force of gravity, forming a liquid. As the tmperature continues to drop the particles become more ordered, and their translational energy drops to a very small value and a solid forms. Condensation occurs when the intermolecular attraction between a pair of particles exceeds the kinetic energy of the collision.

The 'stickyness' exhibited by particles at the lower temperatures, which result in the formation of liquids and eventually solids is due to intermolecular attractive forces. Intermolecular means between molecules. Intramolecular means between atoms. Intramolecular forces are what we call covalent bonds and are very strong (100 - 1000 kJ/mol) (see also Table 9.2 in Silberberg). Intermolecular forces are between molecules and are weak (0.1 - 40 kJ/mol). Intermolecular forces are less directional compared to covalent bonds and operate over a longer range compared to covalent bonds. It is intermolecular forces which explain the formation of liquids and solids in covalent compounds. Intermolecular attractive forces are electrostatic in nature.

Intermolecular forces are classified into the following categories;

* ion-dipole

* dipole-dipole

* induced dipole-induced dipole (London dispersion forces)

* hydrogen-bonding

To determine which type(s) of intermolecular attractive forces occur for a particular substance the most important characteristic to determine for the substance is its polarity. A polar molecule possess a permanent dipole moment (see pages 384 - 386 in Silberberg) as a result of its molecular shape and from the unequal sharing of electrons in chemical bonds which produces the separation of charge which produces the dipole.

To recognize a polar substance (you have to do this on the review problem set) you must draw the Lewis structure (see pages 361 - 371 in Silberberg, CHEM 1314 lecture notes). After drawing a Lewis structure look at the central atom.

Rules for predicting whether a molecule is polar (has a permanent dipole) or is nonpolar

A simple example is HCl. The Lewis structure for HCl is;

The pair of electrons in the covalent bond between hydrogen and chlorine is unequally shared due to the difference in electronegativity between hydrogen and chlorine. Chlorine has a greater electronegativity compared to hydrogen and as a result the electrons in the covalent bond spend a greater proportion of the time closer to the chlorine nucleus. So for chlorine the electrons in the covalent bond spend more time nearer its nucleus producing a small partial negative charge in the molecule. Since the hydrogen atom 'sees' the electron in the covalent bond less frequently it has a partial positive charge. This permanent separation of some small amount of charge on the HCl molecule produces a permanent dipole. Here is a picture of the charge on a space-filling model of the HCl molecule;

A collections of HCl molecules will align themselves such that the negative end of one HCl molecule is attracted to the positive end of an adjacent HCl molecule.

The ion-dipole force results from the attraction of an ion of positive or negative charge and the oppositely charged end of the dipole on the polar molecule. The strength of the attraction depends on the magnitude of the charge on the ion, the magnitude of the dipole moment and the distance between center of the ion and the midpoint of the polar molecule. Ion-dipole forces are important for ionic compounds which dissolve in water.

In class we viewed an animation of NaCl dissolving in water. Here are two figures from the animation.

In this figure there are two examples of ion-dipoe attractive forces. Notice the water molecules surrounding the chloride (large green sphere) and the water molecules surrounding the sodium ion (smaller gray sphere). In particular notice the orientation of the water molecules. Chloride ion has a neqative charge in NaCl, and water is a polar molecule with a permanent dipole. Since water has a dipole there is a region of partial negative charge, on the oxygen atom, and regions of partial positive charge, on the hydrogen atoms. The water molecule is oriented so the hydrogen atoms, with the partial positive charge, are close to the negatively charged chloride ion. In the case of the positively charged sodium ion, the oxygen atoms are aligned towards ion. The interaction between the chloride ion, and the sodium ion and the water molecules is an example of ion-dipole interaction.

Dipole-dipole forces of attraction occur between polar molecules. This type of attractive intermolecular force occurs between the polar molecules of a pure substance, such as HCl, or between two different polar molecules. Here is an animation depicting the attraction when two HCl molecule approach each other. The attraction arises because of the permanent dipole in the polar molecule. Dipole-dipole attractions are relatively weak compared to ion-ion attractions, because the charges on polar molecules are generally quite small. In liquids the molecules are free to move relative to each other, but will do so such that sometimes orientations of adjacent molecules are attractive and sometimes they are repulsive. The overall average effect is attractive.

The larger the magnitude of the dipole moment the stronger the attraction given the molecules are of similar mass and size.

London dispersion forces occur between atoms or molecules of nonpolar substances. Monoatomic atoms (noble gases), diatomic molecules (H2, N2, O2, F2, Cl2) and nonpolar compounds (CH4, CCl4, BF3, BeH2, etc.) are all characterized by a symmetric sharing of electrons in the atom or molecule. These compounds do not have permanent dipoles as occur in heteronuclear diatomic molecules (HCl, HBr, HI, etc) and polar compounds (SO2, H2S, NCl3, etc).


In class we saw an animation of a collection of monoatomic neutral atoms. If we take "snapshots" of the electron distribution we would generally see a symmetric distribution of the electron density. However, occasionally we see, but not very often, an unequal sharing of the electrons. A "snapshot" an instant later would reveal a return to an equal distribution of the electrons. Every once and a while we observe that the electrons are unequally distributed around the nuclei, when this occurs there is a very small charge separation created which gives rise to an instantaneous dipole.

Another atom near this instantaneous dipole will also be effected causing a shift of its electron distribution resulting in a small dipole around it. When this occurs, even for an instant there is a small attraction between the two molecules. The strength of the London dispersion forces depends on how easily the electron cloud is distorted or polarized.

The larger the molecule the further the electrons are from the nucleus and the easier the electron cloud can be polarized. So the magnitude of the dispersion forces increases with increasing molecular size. Because increasing molecular size generally means increasing molecular mass it can be re-stated that dispersion forces increase with increasing molecular mass. Note: dispersion forces operate in all molecules whether they are polar or nonpolar.

It has been shown (not in class) that the force of attraction between two nonpolar molecules is inversely proportional to the seventh power of the distance and directly proportional to a property of each molecule called polarizability. Polarizability of an atom or a molecule is a measure of the ease with which the electrons and nuclei can be displaced from their average positions. When the electrons occupy a large volume of space, which occurs in an atom or molecule with many electrons, the polarizability of the substance is large. The units on polarizability are the units of volume, m3.

When the polarizability is large for a particular atom or molecule the magnitude of the instantaneous dipole can be large with the result producing a stronger attraction between particles.

The electrons which are the most easily displaced in an atom or molecule are the valence electrons, these are the furthest from the nucleus. So valence electrons make the greatest contribution to the polarizability. The force acting on the valence electrons depends on their distance from the nucleus and on the core charge. For any group in the periodic table the core charge remains constant, so we expect polarizability to increase as the atomic size increases. So the polarizability of HI is greater than the polarizability of HF.

In molecules with large numbers of atoms the polarizability will be larger compared to smaller molecules. The polarizability of a molecule increases with both increasing size and increasing numbers of atoms in the molecule. So we expect the magnitude of the instantaneous dipoles, and therefore the strength of the London forces, to be greater the greater the number of atoms in a molecule and the larger the atoms.

The polarizability of N2 is greater than H2, and that of CCl4 greater than CH4, and CO2 greater than that of CO. The strength of the intermolecular attractive forces is reflected in the boiling points of the substances.


Two contributions to the intermolecular attractions between covalent molecules; 1) dipole-dipole forces (present only when the molecule is polar) and, 2) London forces (present between all molecules and are particularly important for large molecules.) London forces are often stronger than the dipole-dipole forces between polar molecules. It is rare that dipole-dipole forces will dominate the properties of a molecule unless the dipole-dipole forces are particularly strong. This occurs for hydrogen-bonded systems.

Lets look at boiling point trends.

All of these properties can be understood in terms of the intermolecular attractive forces which exist between water molecules. The particular intermolecular attractive force is called hydrogen-bonding. Hydrogen bonding is another intermolecular force, which is stronger than London and dipole-dipole forces. Hydrogen bonding forces occurs in a particularly special group of polar compounds. These compounds are characterized by the X-H bond, where X can be O, N, or F. Examples of compounds that exhibit hydrogen bonding forces are H2O, NH3 and HF. Oxygen, nitrogen and fluorine are small strongly electronegative atoms. In a covalent bond with hydrogen these atoms attract the pair of electrons giving rise to a partial positive charge on the hydrogen atom. This partial positive charge on the hydrogen atom is very interested in any negative charge in another adjacent molecule that comes close to it. So when another polar molecule which contains an atom such as O, N, F, Cl or S come near an X-H bond a hydrogen bond can form.

So what would a hydrogen-bonding interaction look like when two or three molecules of water were close to each other? To do this lets begin with just a lewis structure for water. Draw the Lewis structure for water. When you are finished look at my drawing.


Bulk Properties of Liquids

As an introduction to the need to define and discuss intermolecular attractive forces we began our discussion of the bulk properties of gases, liquids and solids. Some of these are discussed in Section 12.4 on pages 440 - 447 in Silberberg.

We talked about the following bulk properties;

* molar volume

* compressibility

* thermal expansion

* viscosity

* surface tension

Molar volumes of gases are much larger compared to the molar volume of liquids and solids. For gases, assuming ideal behavior, the molar volume is 22,400 mLs at 0 degrees C and 1 atmosphere. For liquids and solids the molar volume ranges between 10 and 100 mLs at room temperature and pressure. Molar volume is inversely proportional to the density of the substance.

Compressibility is a measure of the extent to which the volume of a given amount of substance will shrink when it is squeezed. Liquids and solids are nearly incompressible, while gases are highly compressible. The high compressibility of gases is a result of the large amount of empty space which exists between gas particles. In liquids and solids very little free space is present between particles. Compressing liquids or solids, even at high pressures, results in very small changes in volume.

Thermal expansion is the increase in volume of a substance with increasing temperature. For ideal gases the coefficient of thermal expansion is 1/273 at 0 degrees C. A temperature increase of 1 degrees C changes the volume of a gas by 0.366 %. At 25 degrees C the volume of 1 g of water is 1.00296 mL, at 26 degrees C the volume is 1.00323 or a change of about 0.026 %.

Viscosity measures the resistance to flow for a liquid. The higher the viscosity the slower a liquid will flow. From a atomic view viscosity can be understood in terms of the ability of particles to slide over and past each other. If a fluid has a high viscosity particles "stick" to each other and do not slide past each other easily. One wonders what causes this stickiness in the liquid phase.

Surface tension measures how much energy is required to move a particle from the middle of a fluid to the surface. The easier this can occur the lower the surface tension. Liquids under the influence of gravity and on a surface for which the adhesive forces are small, e.g. water on wax paper, will cause the liquid to "bead-up". In an environment where the gravity is low or absent fluids will form a spherical shape as this shape minimizes the surface area. Water has a comparatively high surface tension. It is difficult to move a liquid from within the liquid to the surface. One wonders why it is difficult to move a molecule of water from within the middle of the liquid to the surface. Water molecules must "stick" to each other.