Molecular Structure (Cont.)

Electronegativity

Readings for this section.

Petrucci: Section 10.3 and 10-7

Experimental evidence leads us to recognize that not all covalent bonds involve the equal sharing of the electron pair in the bonds.  Certain atoms attract the electrons in the bond pair more than other atoms.  If a molecule contains two different types of atoms that are covalently bonded to each other, it is almost certain that the electron pair will not be equally shared.  In the extreme, the bond is not a covalent bond but an ionic one.

Introduction

Linus Pauling ,(L. Pauling, The Nature of the Chemical Bond and the Structure of Molecules and Crystals, Cornell University Press:  Ithaca, N.Y., 1948, pp. 58 - 75) used bond energies in his determination of a scale that could be used to measure the degree of unequal sharing or degree of ionic character in a bond.  He used symbol χA and χB (Greek letter chi), to represent the electronegativities of atoms A and B, respectively.

There is no real way to calculate an atoms ability to pull electrons to itself in a covalent bond (a.k.a. electronegativity) without actually taking into account the other atom involved in the bond.  We can calculate the difference in electronegativities  using bond energies as the difference between the bond energies of the A-B molecule less the geometric average of the bond energies of the A-A and B-B molecules.

The difference in the electronegativities can now easily be calculated using

We rarely actually use these equations, relying mostly on tables of data compiled from large amounts of statistical data.  Pauling originally set the value of the electronegativity for F to 4.00 and calculated other atoms' electronegativities from there.  Today, we use 3.98 for F, giving a scale with all atoms having electronegativities less than four and greater than zero.

Differences in electronegativities of two bonded atoms can be used as a predictor of the bond's character.  Differences greater than 2.0 are considered indicative of an ionic bond while differences less than about .4 indicate a pure covalent bond.  Intermediate values for the electronegativity differences indicate a polar covalent bond where the electrons are not equally shared between the two atoms that are bonded.

Table of average Electronegativities calculated according to the method outlined above.

H
2.20  
Li Be   B C N O F
0.98 1.57 2.04 2.55 3.04 3.44 3.98
Na Mg Al Si P S Cl
0.93 1.31 1.61 1.90 2.19 2.58 3.16
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br
0.82 1.00 1.36 1.54 1.63 1.66 1.55 1.90 1.88 1.91 1.95 1.65 1.81 2.01 2.18 2.55 2.96
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I
0.82 0.95 1.22 1.33 1.60 2.24 1.90 2.20 2.28 2.20 1.93 1.69 1.78 1.88 2.05 2.10 2.66
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At
0.79 0.89 1.27 1.30 1.50 2.36 1.90 2.20 2.20 2.28 2.54 2.00 1.83 2.10 2.02 2.00 2.20

 

The Table of electronegativities as presented in Petrucci Fig 10.6, p. 382.

H
2.1  
Li Be   B C N O F
1.0 1.5 2.0 2.5 3.0 3.5 4.0
Na Mg Al Si P S Cl
0.9 1.2 1.5 1.8 2.1 2.5 3.0
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br
0.8 1.0 1.3 1.5 1.6 1.6 1.5 1.8 1.8 1.8 1.9 1.6 1.8 1.8 2.0 2.4 2.8
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I
0.8 1.0 1.2 1.4 1.6 1.8 1.9 2.2 2.2 2.2 1.9 1.7 1.7 1.8 1.9 2.1 2.5
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At
0.8 0.9 1.3 1.3 1.5 2.4 1.9 2.2 2.2 2.2 2.4 1.9 1.8 1.8 1.8 2.0 2.2

Dipole Moments

One last item we need to discuss while we're on the topic of the three-dimensional structure of the molecules is their dipole moments.  Dipole moments (μ) are caused by two opposite charges of magnitude δ  in Coulombs separated by distance R in meters.

|μ|= δ R.

Dipole moments are most often expressed in units of debye where 1 debye = 3.33610-30 coulomb meters.  

Here, we can see that the larger the charge, the larger the dipole moment and the smaller the distance the smaller the moment.  Thus molecules that have large charge separations can have large dipole moments.  The absolute value sign around the symbol |m| is because this equation only represents the magnitude of the dipole moment.  There is also a direction.  We usually denote the direction as being from the positive to the negative charge.  Thus, if two charges are separated by a distance as indicated in the diagram below, the dipole moment can be represented by a vector starting at the positively charged atom and going along the bond to the negatively charged one.

In theory, any bond that has charge separation of any amount will be polar although we may not be able to measure the polarity of some such bonds with very small charge separations.  The only truly non polar bonds are those that are 100% covalent like the bond in a homonuclear diatomic molecule, e.g., H2, O2 or N2, etc.

HF dipole momentHere we can see an example of the HF molecule where two charges (both of magnitude d less than the magnitude of one electron charge,  |d| < |e|) are separated by the bond length d of the molecule.  Thus, we draw the dipole moment as a vector starting at the positive H (note the + sign) and pointing to the F atom.  Unfortunately, the centre of the charge around each atom may not be exactly the centre of the nucleus so R is not necessarily the same as d.

We say that HF is polar since the molecule has a dipole moment.


In the case of more complicated molecules, we must add the vector dipole moments of each bond to get an overall dipole moment of the molecule.  In the case of water, there are two bonds, each of which are polar.

Here, we see that the two bond dipoles (red) lie along the H-O bond directions.  These two add vectorially to produce a molecular dipole (red) that lies half-way between the two bonds.  Additionally, the lone pairs increase the molecular dipole since they consist of significant negative concentration.

We call water a polar molecule because it has a molecular dipole.

In the case of BeCl2, there are two polar bonds but they are at 180? from each other and of the same magnitude.  Thus, they cancel each other out.  There is no molecular dipole.  The molecule is non-polar even though the bonds are polar.

In this case, BF3, there are three bond dipoles, each at 120 degrees from each other.  These add vectorially in such a way as to cancel each other out.  There is no molecular dipole and so the molecule is termed non-polar, even though the bonds themselves are polar.

The methane molecule has four polar bonds, as indicated in the picture.  These are oriented at the tetrahedral angle of 109.5? from each other.  again, these four bond dipoles cancel out so that the molecule becomes non-polar even though the bonds themselves are polar.


In Phosphorous pentachloride, it may not be quite so obvious but think of the equatorial (trigonal) ligands like that of the BF3.  The dipole moments cancel out.  The axial ligands are at 180? to each other and also cancel out.  Thus, there is no molecular dipole even though the bonds are polar.


Similarly, the octahedral molecule SF6 would be non-polar despite the fact that all the bonds are polar. 

In general, if any of these molecules' symmetry is broken by reaction, they would become polar.  For example, there is a species called the methyl radical.  It is short-lived, highly reactive with a single electron where one hydrogen would have been.  It is also polar as the picture shows. Similarly ammonia, NH3 with three hydrogens and a lone pair would be a polar molecule.

 

Thus, the polarity of a molecule depends not only on the polarity of the bonds in the molecule but also on it's symmetry.  Molecules with certain types of symmetry are not polar even if they have polar bonds.

Can we use these ideas to explain why ammonia (NH3) has a larger molecular dipole than NF3?

the polarity of the individual bonds would be similar (but in opposite directions) for NH3 as for NF3, however, in the case of the NH3, the lone pair's negative concentration will augment the polarity contribution from the polar bonds while for the NF3, the lone pair's dipole would subtract from the dipole contributed by the NF bonds.

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Prof. Michael J. Mombourquette.
Copyright © 1997
Revised:August 29, 2013.