Molecular Structure and Polarity

All matter is held together by force. The force between atoms within a molecule is a chemical or intramolecular force. The force between molecules is a physical or intermolecular force. We learned about intramolecular forces and the energy it took to overcome these forces, earlier in our chemical studies. We will now breifly delve into a discussion of intermolecular forces.


The Nature of Intermolecular Forces:

The Intermolecular Forces (forces between molecules) are weaker than Intramolecular Forces (The Chemical Bonds within an Individual Molecule). This distinction is the reason we define the molecule in the first place. The properties of matter result from the properties of the individual molecule (resulting from chemical bonding) and how the molecules act collectively (resulting from intermolecular forces).

Intermolecular Forces are longest-ranged (act strongly over a large distance) when they are electrostatic. Interaction of Charge Monopoles (simple charges) is the longest-ranged electrostatic force.
Charge-Charge forces (found in ionic crystals)

For like charges (+,+) or (-,-), this force is always repulsive. For unlike charges (+,-), this force is always attractive.

Charge-Dipole Forces:
An uncharged molecule can still have an electric dipole moment. Electric Dipoles arise from opposite but equal charges separated by a distance. Molecules that possess a dipole moment are called Polar molecules (remember the polar covalent bond?). Water is polar and has a dipole moment of 1.85 Debye. The Debye is a unit of dipole moment and has a value of 3.336 x 10-30 Coulomb meter.

When salt is dissolved in water, the ions of the salt dissociate from each other and associate with the dipole of the water molecules. This results in a solution called an Electrolyte

The force may be understood by decomposing each of the dipole into two equal but opposite charges and adding up the resulting charge-charge forces. Notice that the Charge-Dipole Forces depend on relative molecular orientation. This means that the forces can be attractive or repulsive depending on whether like or unlike charges are closer together. On average, dipoles in a liquid orient themselves to form attractive interactions with their neighbors, but thermal motion makes some instantaneous configurations exist fleetingly that are, in fact, repulsive.

Dipole-Dipole forces exist between neutral polar molecules. Again, this force may be understood by decomposing each of the dipole into two equal but opposite charges and adding up the resulting charge-charge forces.



The following table demonstrates the effect of the dipole moment on the boiling point of several substances:
 
 

Substance
Molecular Mass
[g/mol]
Dipole moment
[Debye]
Normal Boiling Point
[ K ]
Propane
44
0.1
231
Dimethyl ether
46
1.3
248
Chloromethane
50
2.0
249
Acetaldehyde
44
2.7
294
Acetonitrile
41
3.9
355

Electrostatic forces are defined (categorized) by the symmetry of the partners involved in the interaction. This symmetry is labelled by the first non-zero moment of the charge distribution, i.e. Monopole, Dipole, Quadrupole, etc. Electrostatic forces only exist between molecules with permanent moments of their charge distribution; Molecules do not have to distort or fluctuate in order to exhibit electrostatic intermolecular forces.


Identifying Polar Molecules

A polar molecule is one with a permanent Dipole Moment. A polar molecule must have a slightly positive end opposite a slightly negative one. This cannot happen if the molecule is too symmetric! If a molecule is 'spherical' enough, then each end of the molecule will have the same properties and in must be non-polar.

Examples of molecules that have exactly zero dipole moment and therefore be referred to as non-polar by symmetry are:
Homonuclear Diatomics, and molecules with a center of inversion, eg. H2 and CO2
Symmetric 'Disk" shaped molecules (Trigonal Planar, Square Planar, Hexagonal Planar, etc), eg. Benzene (C6H6) and BCl3
Symmetric 'Ball' shaped molecules (Tetrahedral and Octahedral), eg. Methane (CH4) and SF6
Note: These arguments only hold for symmetrically substituted molecules; Asymmetric substitution giverise to a net dipole.
Molecules that have 'low' symmetry will always have at least a small dipole moment and therefore be referred to as polar. Examples of such low symmetry molecular shapes include:
Bent molecules, eg. water
T-shaped molecules, eg. IF3
See-Saw molecules, eg. SF4
Pyramidal molecules (trigonal pyramidal, square pyramidal, etc) [NOT bi-pyramidal], eg. NF3, BrF5

When molecules are less than 'perfectly' symmetric, a dipole moment results from the unequal sharing of the electrons between bonded atoms. Remember what indicated how unequal the sharing of electrons was? Yup, the electronegativity:

We normally refer to the charge separation resulting from the unequal sharing of the electrons in a chemical bond as a Bond Dipole. The greater the difference in Electronegativities between the bonded atoms, the greater the Bond Dipole.

The molecule's overall dipole moment is the result of the vector sum of all the bond dipoles within it. In the symmetric molecules above, no matter what the bond dipoles are, the net dipole moment of the molecule is zero because the bond dipoles cancel.

Lone pairs contribute to the molecule's dipole moment even though the do not constitute a 'bond'. Clearly the nucleus 'end' of the lone pair is positive and the electron 'end' is negative so one might think of a 'lone pair dipole' contributing to the polarity of the molecule in analogy to a bond dipole. This behavior is demonstrated in the relative magnitudes ad directions of the dipole moments in the molecules PH3 (m=0.58 D), NH3 (m=1.47 D), and PF3 (m=1.03 D)


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