0.7 Molecular structure and physical properties  (Page 4/5)

Focusing again on the water molecule, how can we account for the existence of a dipole moment in a neutralmolecule? The existence of the dipole moment reveals that a water molecule must have an internal separation of positive partialcharge $\delta$ and negative partial charge $-\delta$ . Thus, it must be true that the electrons in the covalent bondbetween hydrogen and oxygen are not equally shared. Rather, the shared electrons must spend more time in the vicinity of one nucleus thanthe other. The molecule thus has one region where, on average, there is a net surplus of negative charge and one region where, onaverage, there is a compensating surplus of positive charge, thus producing a molecular dipole. Additional observations reveal thatthe oxygen "end" of the molecule holds the partial negative charge. Hence, the covalently shared electrons spend more time near theoxygen atom than near the hydrogen atoms. We conclude that oxygen atoms have a greater ability to attract the shared electrons in thebond than do hydrogen atoms.

We should not be surprised by the fact that individual atoms of different elements have differing abilities toattract electrons to themselves. We have previously seen that different atoms have greatly varying ionization energies,representing great variation in the extent to which atoms cling to their electrons. We have also seen great variation in the electronaffinities of atoms, representing variation in the extent to which atoms attract an added electron. We now define the electronegativity of an atom as the ability of the atom to attract electrons in a chemical bond. Thisis different than either ionization energy or electron affinity, because electronegativity is the attraction of electrons in a chemical bond , whereas ionization energy and electron affinity refer to removal andattachment of electrons in free atoms. However, we can expect electronegativity to be correlated with electron affinity andionization energy. In particular, the electronegativity of an atom arises from a combination of properties of the atom, including thesize of the atom, the charge on the nucleus, the number of electrons about the nuclei, and the number of electrons in thevalence shell.

Because electronegativity is an abstractly defined property, it cannot be directly measured. In fact, thereare many definitions of electronegativity, resulting in many different scales of electronegativities. However, relativeelectronegativities can be observed indirectly by measuring molecular dipole moments: in general, the greater the dipolemoment, the greater the separation of charges must be, and therefore, the less equal the sharing of the bonding electrons mustbe.

With this in mind, we refer back to the dipoles given in . There are several important trends in these data. Note that each hydrogenhalide ( $HF$ , $H\mathrm{Cl}$ , $H\mathrm{Br}$ , and $HI$ ) has a significant dipole moment. Moreover, the dipole momentsincrease as we move up the periodic table in the halogen group. We can conclude that fluorineatoms have a greater electronegativity than do chlorine atoms, etc. Note also that $HF$ has a greater dipole moment than $H_{2}O$ , which is in turn greater than that of $N{H}_3$ . We can conclude that electronegativity increases as we move across the periodic table from left to right in a single period. These trends hold generally incomparisons of the electronegativities of the individual elements. One set of relative electronegativities of atoms in the first threerows of the periodic table is given in .