Example 3.5. Predicting, where or not a molecule is polar.

Problem: is the IF₅ molecule polar?

Solution:the first step is to determine the shape of the molecule using the VSEPR theory.

Iodine is the central atom of the molecule; its valent shell configuration is 5s²5p⁵, additionally I atom has empty 5d orbitals. The coordination number of iodine in IF₅ molecule equals 5, all I-F bonds are single ones. There is one lone pair in the valence shell of the iodine atom, so, the molecule can be presented by a formula AX₅E. According to VSEPR theory, the AX₅E molecules have a tetragonal pyramidal shape.

The I-F bonds are polar, because electronegativities of the atoms are different (2.8 and 4.1 for I and F, respectively). Four bond dipoles located in one plane cancel each other, thus, remaining dipole has a lone pair on the opposite side and is not cancelled. So, the bond dipoles in IF₅ molecule don’t cancel each other entirely. Therefore, the molecule will be polar.

Answer: the IF₅ molecule is expected to be polar.

3.6. Intermolecular Forces

Van der Waals forcesarerelatively weak electric forces that attract neutral molecules to one another. Intermolecular attractions are much weaker than intramolecular attractions. There are three kinds of intermolecular interactions: permanent dipole - permanent dipole, permanent dipole-induced dipole and dispersion forces between instantaneous dipoles. Attractions between atoms within molecules (chemical bonds) determine chemical properties of a substance; attractions between molecules determine its physical properties, such us melting and boiling points.

Dipole-dipole attraction. If molecules of a substance are polar, the individual dipoles tend to orient themselves so that partial positive charge on one is near the partial negative charges on others. This alignment is far from perfect, particularly in liquids. Nevertheless, the attractions between the oppositely charged ends of the dipoles outweigh the repulsion between like-charged ends, and a net overall dipole-dipole attraction exists between them. Dipole-dipole attractions are normally considerably weaker than ionic or covalent bonds, usually they are only about 1% as strong.

Hydrogen bonds. A particularly strong dipole-dipole attraction occurs between molecules in which hydrogen is covalently bonded to a small atom of the most electronegative element such as fluorine, oxygen or nitrogen. In these instances the bond is highly polar and the small hydrogen atom carries substantial partial positive charge. Because the positively charged end of this dipole can closely approach the negatively charged end of the neighboring dipole, the force of attraction between the two is quite large. This special case of the dipole-dipole attraction is called a hydrogen bond. Hydrogen bonds are about 5 to 10% as strong as an ordinary covalent bond. Because of hydrogen bonds hydrogen fluoride, water and ammonia have melting and boiling points, significantly higher than that of HCl, H₂S and PH₃, respectively. Hydrogen bonds are also responsible for orientation of molecules in the crystalline substances. Water molecule has two hydrogen atoms and two lone pairs, therefore it form four hydrogen bonds in ice – each water molecule is surrounded by four others (a tetrahedral arrangements of molecules is observed). This causes ice to be a very “open” structure and makes ice less dense than liquid water.

Dispersion forcesare also known as "London forces" (named after German physicist Fritz London who first suggested how they might arise).Even uncombined atoms and nonpolar molecules experience weak attraction; otherwise substances like helium and hydrogen couldn’t be condensed to liquid. Attractive forces that hold uncombined atoms and nonpolar molecules together in the liquid or solid state are called dispersion forces or London forces. When electrons move around in the atom or molecule, their motion is somewhat random. So, at any instant there is a chance that more electrons will be on one side of the particle than the other.The dipole appeared is called an instantaneous dipole because its existence is only momentary. In a collection of atoms or molecules the electron motions in nearby neighboring particles are not entirely independent. As instantaneous dipole forms in one of atoms or molecules, negatively charged end of the dipole repels electrons away in the particle alongside. Thus an induced dipole in neighboring atom or molecule appears. The instantaneous dipoles attract each other, but these attractions are rather weak in average because of their fleeting existence. Nevertheless, dispersion attraction forces are present between all particles – atoms, molecules (polar or nonpolar) or ions. They play only minor role in the attractions between ions; but their role in attractions between molecules of all kinds, especially the nonpolar ones is significant.

The strengths of the dispersion forces depend on several factors. One of these is the number of atoms in the molecule. In complex molecules with many atoms, the cumulative effect of the dispersion forces may be quite substantial. Another factor that influences the strengths of the dispersion forces is an ability of the atoms composing the moleculeto deform calledpolarizability. As the size increases, the attraction between outer electrons and the nuclei decreases. Because of this, the electron cloud of a large atom is more easily distorted (polarized) and it is easier to create the instantaneous dipoles that are responsible for the dispersion forces. The result is that molecules, composed of large, easily polarized atoms, attract each other more strongly, than molecules composed of small atoms. You may consider molecules with the same general formula, such as halogens (F₂, Cl₂, Br₂ and I₂) as an example.

[1] Except for He, atom of every noble element has eight electrons in its outer shell (filled ns²np⁶ shell).

[2] For example, quadruple bonds are observed in chromium (II) acetate Cr₂(CH₃COO)₄×2H₂O and in [Re₂Cl₈]^2- anion

[3] Chemical bonding in transition metals is more complex than in s- and p- ones and may involve some covalent bonding due to overlap of localized d-orbitals .