intermolecular+forces

An **//intermolecular force//** (a.k.a. Van der Waals force) is an attractive force between two or more molecules. There are four major types of intermolecular forces (IMFs):
 * London dispersion
 * dipole-dipole
 * hydrogen-bond
 * ionic

Each of these will be introduced below.

There are many other factors involved in understanding the strength of an IMF. Here are a few:
 * 1) Does the molecule have a //dipole moment//?
 * 2) Does the molecule contain //polarizable// atoms?
 * 3) Are the bonds in a molecule //nonpolar// or //polar//?
 * 4) How much //surface area// does the molecule have?
 * 5) How //heavy// is the molecule?

Once the strength of the IMF for a molecule is understood, it can then be used to study some physical properties of that molecule. These IMFs help to explain properties such as:
 * the phase of the substance at a given temperature (i.e. room temperature)
 * the //normal// //melting point// for the substance
 * the //normal// //boiling point// for the substance
 * the //lattice strength// of a crystalline substance
 * the //viscosity// of a liquid
 * the //RMS velocity// and //rate of effusion// of a gas

The four IMFs will now be defined, along with some important terms that were used above that relate to each type of IMF.


 * __The London Dispersion IMF:__**

This force occurs between nonpolar covalent molecules. Its strength relies primarily on the //surface area// of the molecule, sometimes referred to as its //Van der Waals// area. The larger the molecule is, the more surface area it has. Here are three examples: propane, octane, and paraffin:




 * Recall that propane is used in gas grills and in normal conditions (1 atm, 298 K) is a gas. This means its molecules are not tightly held together at all, and it has a weak London dispersion IMF.
 * Octane and its isomers are the ingredients of gasoline, which is a liquid in normal conditions. These molecules are larger and have more surface area than propane molecules. This leads to stronger IMF strength, thus the molecules stick together so much that they form a liquid.
 * Paraffin is a waxy hydrocarbon used in candles. Its large size (and thus surface area) makes it a solid at room temperature, with a boiling point greater than 325 C.


 * __The Dipole-Dipole IMF:__**

When a molecule contains a highly-electronegative element (such as oxygen, nitrogen, or a halogen) bound to a low-electronegative element, such as hydrogen, carbon, or silicon), there exists **polarity**.

In simple diatomic or linear molecules, the most critical IMF function is individual bond polarity. This determines whether the bond is nonpolar covalent, polar covalent, or polar ionic.

This bond must still be **covalent** in order for the compound to be considered dipole-dipole.

Take a look at some small organic molecules, acetone and ethyl acetate:



The oxygen atoms (in red) have a high electronegativity, and will thus have lots of "negative" electrons concentrated around them. The hydrogens (white) are going to be sharing their electron with their adjacent carbons, so their positive nuclei are more exposed. This sets up a positive-negative "field" on the molecule, and it becomes polar (a 'dipole').

Think of a dipole molecule as a magnet. Magnets are dipoles (N and S poles). They have attraction when they align N to S, or S to N. The same thing will occur in polar molecules.

Acetone and ethyl acetate contain three carbons. Acetone boils at 56 C, and ethyl acetate at 77 C. Both are liquids at room temperature. Compare these to propane, which also has three carbons (see above), and boils at -42 C.

__**The Hydrogen-Bond**__ __**IMF:**__

The term hydrogen-bonding is a bit of a misnomer - there isn't a covalent bond being formed between molecules. However, this attractive force is one of the most powerful in terms of small-molecule IMFs.

This IMF applies to a unique structure: there must be a nitrogen, oxygen, or fluorine atom having a free lone pair, covalently bonded to a hydrogen, within the structure of the molecule.

Any molecule containing this N-H, O-H, or F-H structure will have the hydrogen-bonding IMF.

Previously, we looked at three-carbon molecules like propane and acetone. Now look at isopropanol (the main ingredient in rubbing alcohol), another three-carbon molecule:

As you can see, the oxygen atom meets the criteria: it has lone pairs and is bonded to a hydrogen atom.

The boiling point of isopropanol is 82.5 C, higher than the other three-carbon molecules we discussed earlier.

Water is the most obvious of examples for hydrogen bonding, however. It is such a small and light molecule, yet boils at an astounding 100 C!



Let us take a look at three diatomic molecules: nitrogen gas, nitrogen monoxide gas, and crystalline beryllium oxide. Their boiling points are 77 K, 110 K, and 4200 K, respectively. These physical properties can be explained using IMFs and polarity.

First, nitrogen gas contains one N-N bond (actually a triple bond), which is considered nonpolar covalent because the two atoms have the same electronegativity value. This means the molecule has no net //dipole moment// and only attract using London dispersion force. Because of the tiny surface area of the N 2 molecule, its IMF will be very weak. Therefore, N 2 has a very low boiling point.

In nitrogen monoxide, NO, the O atom is more electronegative than the N atom. Electron density will be shifted over O, and N will be slightly more positive. This sets up a dipole moment, and therefore NO molecules will have a slightly stronger force holding them together.

Beryllium oxide is ionic. There is no covalent bond; oxygen has all the valence electrons, and beryllium has been stripped of its. Therefore, there are two distinct ionic charges: +2 for Be, and -2 for O. This sets up a strong electrostatic attractive force that holds these ions together for an astoundingly high boiling point.