Pauling defined electronegativity as the power of an atom in a molecule to attract electrons to itself. Previously, we discussed covalent bonding, in which electrons are shared between nuclear centers. The concept of electronegativity indicates that sharing may in fact not be equal. Furthermore, since electronegativity is defined as the ability of an atom when in a molecule to attract electrons to itself, it is a relative measure of electron attracting ability. This aspect can be seen most clearly in the case of a heteronuclear diatomic molecule like H-Cl. Measurements show that the Cl end of the molecule has a negative charge relative to the H end of the molecule. Thus, Cl appears to be more electronegative than H. A quantitative measure of the difference in electronegativity depends on assigning a reference point for an lectronegativity scale. Pauling developed a method for calculating electronegativities and defined an electronegativity scale. The following table gives Pauling values for the electronegativity of some common elements.
In general, elements in the upper right hand corner of the periodic table have high electronegativities and those on the left hand side of the table have lower electronegativities. The most electronegative element is F (3.98) and the least electronegative is Fr (0.7). Going down a group we see a decrease in the electronegativities of the elements; as we move across a period we see an increase. Note that electronegativity is dimensionless since it measures electron attracting ability on a relative scale.
Later it will useful to remember that C-F bonds are highly polar with the F atom attracting electrons more strongly than the C atom.Similarly, C is just slightly more electronegative than H and is weakly polar. Compounds like H-Cl are highly polar.
The existence of the solid and liquid phases of matter, including for example the liquid phase of atomic helium, indicates that there are forces of attraction between both atomic and molecular systems. For example, helium, which exists as a monoatomic gas at room temperature, forms from gaseous helium (P=1 atm) when cooled to -269 C. We know that as we cool He, the average thermal kinetic energy of He atoms decreases. At -269 C, weak forces of attraction between He atoms are sufficient to cause the formation of a condensed liquid phase. As water liquefies at 100 C and forms a solid at 0 C; we commonly find water as a gas, liquid and solid in our home kitchens. He exists as a monoatomic gas while water consists of two atoms of H and one of oxygen bound together. The O-H bond distance is a little less than 100 pm
Figure 1. A water molecule-oxygen red, hydrogen white
A study of the chemistry of water indicates that freezing and boiling have no effect on the O-H intramolecular bond; thus, it is reasonable to assume that the forces which cause the condensation and freezing of water occur between water molecules and do not disrupt intramolecular bonds. As a consequence, we can say that intermolecular forces of attraction appear weaker than normal covalent bonds. The fact that water condenses at 100 C while He requires cooling to -269 C indicate that water molecules experience much stronger forces of attraction than He atoms.
Understanding the nature of these forces and their relative strength is important in materials science. Given that atoms and molecules consist of positively charged nuclear centers and electron clouds, it seems reasonable to assume that quantum mechanics and the interaction of charged particles play a major role in determining the nature of intermolecular forces. We will use the quantum mechanical treatment of the problem in a qualitative sense and will spend most of our time exploring intermolecular forces from the standpoint of electrostatic interactions. The three main types of intermolecular forces are dipole-dipole, induced dipole-induced dipole (dispersion) and hydrogen bonding.
In electrostatics, a dipole is defined as a situation in which a positive charge and a negative charge of equal magnitude are separated by a small distance. We can represent this situation schematically as in the following diagram
Here d is a small distance and the equality of the magnitude of the charges can be represented by the relation |σ(+)|=|σ(-)|. The total energy of a two dipole system is reduced when the dipoles interact with each other as pictured below. Note that in this drawing. the positive end of one dipole is attracted to the negative end of a neighboring dipole.
One could envision larger numbers of dipoles interacting with each other in condensed phases with opposite charged atoms attracting one another.
Can molecules exhibit dipoles? Reviewing the major ideas behind electronegativity suggests an affirmative answer to this question. Let�s consider the simple diatomic molecule HCl. The H atom in this molecule has an electronegativity of 2.1 while the electronegativity of Cl is 3.0. The bond length in H-Cl is ~0.13 nm. Thus, the H-Cl molecule forms a dipole with the Cl atom having a negative charge relative to the H atom. This can be represented schematically as:
The interaction of a collection of molecular dipoles should lead to condensation of H-Cl. Experimentally, it is found that H-Cl liquefies at -86 C and forms a solid at -85 C and forms a solid at -114 C. It seems reasonable to assume that this solid consists of a lattice of dipoles as schematically illustrated below on Figure 5.
At low enough temperatures, dipole-dipole forces could hold molecules together in a molecular crystal.
Dipole-Induced dipole forces
If one thinks of a dipole approaching a molecule or atom it should be clear that the dipole will induce or form a dipolar arrangement of charge in the molecule or atom it approaches. The induced dipole arises from the attraction of oppositely charged particles and the repulsion experienced by like charged molecules. A schematic representation of this situation is as follows in Figure 6
This interaction is relatively weak (compared to dipole-dipole interactions). Note also that this type of interaction could occur between dissimilar molecules; for example, between a dipolar molecule and one which does not possess a dipole. It also could occur between a dipolar molecule and neutral atom.
Induced dipole-Induced dipole forces
A more complicated situation can occur for both atoms and molecules. Let�s first consider the case of atoms, such as the atoms of the noble gases. Over extended time intervals, the center of negative and positive charge in a neutral atom will coincide and the atom will have no net dipole. However, quantum mechanics does allow for transitory states in which there is a difference in the center of positive and negative charge. In other words, transitory states can lead to the formation of an instantaneous dipole. These instantaneous dipoles can induce dipoles in neighboring atoms. A schematic illustration of this situation is given below.
These transitory states occur throughout a sample, and cycle in such a manner as to produce a weak intermolecular force. The net effect can be significant. For example, in the case of the noble gases it is interactions like these that lead to the formation of the liquid phase for the noble gases. The normal boiling point for the noble gases is given in the following table. We see that all of the noble gases can be liquefied and that the temperature at which liquid forms increases as the number of electrons in the atom increases. It seems logical that the probability of a transitory state in which the center of positive and negative charge do not coincide would increase as the number of electrons increases.
In our introduction to intermolecular forces, we described the physical behavior of water H2O. The following table compares the physical properties of water to the properties of H2S and H2Se. Clearly the physical properties of water appear somewhat anomalous compared to the other hydrides of group VI .
Electronegativity differences in the above table indicate that the H-O bond in water is substantially more polar than the H-S and H-Se bonds with the H end of the bond being negative relative to the O end of the bond. Furthermore, the positively charged H in the O-H bond can interact with the negative portion of the dipole on neighboring H2O molecules. This is, in a sense, a dipole-dipole interaction with the negative part of the dipole being concentrated in the electrons surrounding the oxygen. To the extent that these electrons have an orientation, the hydrogen has directional character. Since the oxygen atom attracts electrons from two O-H bonds there is an opportunity for two hydrogens from neighboring water molecules to interact with the oxygen on a given water molecule.
Hydrogen bonds are weaker than covalent bonds but much stronger than induced dipole-induced dipole interactions. A complete theory of hydrogen bonding requires the application of quantum mechanics. Typical bond strengths for hydrogen bonds in water are on the order of 5kJ/mole. It should be noted that in specialized cases the bond strengths can be much larger than 5 KJ/mole. Searching the web will allow you to observe the range of reported energies for hydrogen bonds.
In general it is found that hydrogen bonds can occur when H is bound to O, N, and F. Acceptor atoms (F, N, O) do not need to be bound to H. Simple examples of hydrogen bonding are given in the following illustrations. Figure 7 (a) shows the bonding in 1,2-dihyroxypropane and Figure 7 (b) shows bonding in 1-hydroxypropyl-methyl ether. In both figures, the hydrogen bond is demoted with a dotted line. In Figure 7(b) we see an example of a situation where the acceptor atom (O) is not bound to H. In both of these examples the length of the hydrogen bond is about 180 pm; substantially larger than the length of the covalent bonds in the molecule.
Figure 7(a). A molecule of 1,3-dihydroxypropane-oxygen red, hydrogen white, carbon black, hydrogen bond dotted line.
Figure 7(b). A molecule of 1-hydroxypropyl-methyl ether-oxygen red, hydrogen white, carbon black, hydrogen bond dotted line.
It is also possible to have hydrogen bonding occurring in situations where carbon is attached to hydrogen and an electronegative group. It is important to note that hydrogen bonding, which is important in determining the properties of water, also has great significance because of its role in determining the structure and properties of the proteins DNA and RNA. Many aspects of this role are described in biochemistry textbooks. Figure 8 illustrates the complex formed between propamidine and DNA. The helix structure of DNA is clear in the figure along with the large number of hydrogen bonds involving the bases in the DNA strand and also involving the propamidine with DNA. The large number of hydrogen bonds in this structure are responsible, collectively, for the stability of the structure.
Figure 8 Complex formed by propamidine and DNA. Hydrogen bonds dotted lines.