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Dipole moment is the degree of polarity, i.e. the seperation of positive and negative charges. But I am not getting the intuition why and how lone pairs affect the polarity and dipole moment. I cannot connect the separation of positive and negative charges definition with lone pair.

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How does lone pair of a central atom affect the dipole moment?

There is no single answer to your question, let me explain. Unlike a typical covalent bond where the electrons are shared between two nuclei and the electron density is spread out over the entire bond, in a lone pair the electrons are not shared and the electron density is more localized around the atom that has the lone pair of electrons. This increased electron density could lead to a more significant contribution from the lone pair electrons to the molecular dipole moment than from electrons spread out more diffusely in a covalent bond. Next we must understand the directionality of the lone pair of electrons. Consider the two molecules pictured below, ammonia and phosphine. The molecules appear to be very similar, they are in the same column in the Periodic Table.

enter image description here

However in ammonia the $\ce{H-N-H}$ angle is around 107 degrees and the molecule is roughly $\ce{sp^3}$ hybridized, the lone pair and the 3 $\ce{N-H}$ bonds roughly pointing towards the corners of a tetrahedron. You can see that in this case (as shown by the arrows, the "arrowhead" end representing the negative end of a dipole), the lone pair on nitrogen will make a contribution to the molecular dipole moment. Next, let's examine phosphine. The $\ce{H-P-H}$ angle is around 90 degrees and the molecule can be viewed as being unhybridized, the lone pair is an $\ce{s}$ orbital and the 3 $\ce{P-H}$ bonds are constructed from phosphorous $\ce{p}$ orbitals. You can see that in this case, the lone pair on phosphorous, due to its spherical symmetry will not make a contribution to the overall molecular dipole moment.

So in summary, a lone pair of electrons can make a significant contribution to the magnitude of a molecular dipole moment due to the fact that they are more localized than bonding electrons and consequently there is a high electron density. But, the directionality (or lack thereof) of the lone pair must also be assessed, since a lack of directionality may preclude it from making a significant contribution to the overall molecular dipole moment.

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  • $\begingroup$ Could you elaborate a bit more on the directionality part? I don't understand why the dipole moment of the lone pair in ammonia points away from the nitrogen atom. Shouldn't more electron density be closer to the nitrogen atom (where nuclear attraction is greater)? This would imply that the dipole moment's direction (due to the lone pair) would be opposite to the one actually observed. $\endgroup$ – Gerard May 24 '15 at 8:34
  • $\begingroup$ @Gerard The nitrogen nucleus is positively charged. The lone pair electrons are contained in an orbital directed away from the nucleus. The vector describing the nitrogen nucleus - lone pair dipole would be a straight line having it's positive end around the nucleus and its negative end around the center of the lone pair orbital - as I've drawn in the figure above. $\endgroup$ – ron May 24 '15 at 13:49
  • $\begingroup$ @ron , JD lee says, “ the lone pair has no contribution to he dipole moment of the molecule , if the lone pair is present in the pure s or p orbital as s is spherically symmetrical and p is objected equally in opposite directions” Does this statement agree with your answer? $\endgroup$ – Aaryan Dewan Jul 8 '16 at 4:28
  • $\begingroup$ @AaryanDewan Yes, read what I say about lone pair "directionality". $\endgroup$ – ron Jul 8 '16 at 12:46
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    $\begingroup$ @YUSUFHASAN The key word is nucleus. The nucleus contains protons and neutrons and is always positively charged. Add the electrons and you then have a neutral atom. $\endgroup$ – ron Oct 25 '18 at 23:26
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Dipole moment can be defined as the product of magnitude of charges and the distance of separation between the charges. Dipole moment may refer to:

  • Electric dipole moment, the measure of the electrical polarity of a system of charges
  • Transition dipole moment, the electrical dipole moment in quantum mechanics
  • Molecular dipole moment, the electric dipole moment of a molecule.
  • Bond dipole moment, the measure of polarity of a chemical bond
  • Electron electric dipole moment, the measure of the charge distribution within an electron
  • Magnetic dipole moment, the measure of the magnetic polarity of a system of charges
  • Electron magnetic dipole moment
  • Nuclear magnetic moment, the magnetic moment of an atomic nucleus
  • Topological dipole moment, the measure of the topological defect charge distribution
  • The first order term (or the second term) of the multipole expansion of a function

(from wikipedia)

The dipole moment is actually affected by the presence of a lone pair of electrons because the electrons on the central atom can cause shielding effect as the inner orbital electrons does for the outer-orbitals. It can cause hindrances by charge imbalances which could result in imbalance of charges and which could destabilise the atom due to decrease in dipole moment.

Consider a molecule of carbon dioxide where two oxygens are present at each side of carbon:

enter image description here

Here the formal charge on carbon is zero and the same formal charge is there on both the oxygens. Coming to the topic of dipole moments you should consider the lone pair of electrons on oxygen. The lone pair is generally resided on the other side of the bond [as shown in figure]. These electrons are a part of oxygen and these makes oxygen more electronegative because of which the dipole moment increases ans hence the bond becomes stronger.
In most of the cases the central atom in any compound of oxygen is more electropositive atom and electropositive atoms do not contain lone pairs. You can just imagine there are lone pairs on the central atom carbon which could generate some partial negative character in them and hence the bond between them must not be as strong as the bond between an electropositive and electronegative atom.

This all is governed by dipole moment only.

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  • $\begingroup$ the pure shielding effect is not seen in the MOLECULES they are just seen in ATOMS in order to make you understand just used the term but it does'nt mean that it is as same as in atoms here it just hinders the central atom [make's a sort of covering ] from the attached atoms so that the complete atom is'nt exposed to the central atom and there will surely be a difference in the lenght of bond but this is not what you asked keep this topic aside $\endgroup$ – agha rehan abbas Jun 25 '14 at 17:11
  • $\begingroup$ i will try to make the same same thing with ammonia $\endgroup$ – agha rehan abbas Jun 25 '14 at 17:12
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    $\begingroup$ This was hard to read, edit, and I believe most of it is wrong. For example the central atom in $\ce{NH3}$ is the most electronegative element. The picture of $\ce{CO2}$ is just utter garbage and has no source given. The lone pairs of oxygen do not make oxygen more electronegative. $\endgroup$ – Martin - マーチン Oct 10 '17 at 10:56

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