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Lone Pairs of Electrons

For any chemical bonding the presence or absence of lone pair of non-bonded or lone pairs helps in determining the overall shape of the molecule. When lone pairs are present in chemical bonding the existence of several different repulsive forces are observed as well.

These forces exist between bonding pairs, between lone pair and bonded pair as well as in between lone pairs. The VSEPR model is based on different repulsive forces for the three situations.

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Lone Pairs of Electrons Definition

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The lone pairs of electrons are nothing but the non-bonded pairs which remain in many chemical compound. These are usually seen in molecules where the combination of large atom with small atoms where a few of the electrons of the valence cloud remain unpaired.

The presence of these lone pair of electrons in valence cloud is important as these are responsible for all sorts of shapes of the molecules. For every lone pair the angle of 3 degree is reduced from the normal geometrical electron domain angles.

The forces which exist between the pairs are in the order of the greatest repulsion of
  • lone pair and lone pair
  • lone pair and bonding pair
  • bonding pair and bonding pair
The presence of lone pairs are typical of covalent bonding as the ionic bonding is all about transferring of electrons in the valence orbit or cloud.

It’s a specific characteristic of covalent bonding or bonding where the sharing of electrons take place in the valence orbit electron cloud and in spite of sharing between almost all the electrons that are present in valence shell, some of the electron pairs remain unused or unbonded.

These are the electron pairs which finally give rise to the lone pair theory and the theory of repulsion between lone pair and bonded pair and the characteristic shapes of each of the covalently bonded molecules and ions.

How to Find Lone Pairs?

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In order to understand this simple practice of finding the lone pairs in any molecule or polyatomic ions, we need to draw the electron dot structure first. In order to get the electron dot structure we need to follow the basic steps.
  • Provide the symbol of the element
  • Each side of the element symbol could be provided with a max of two electrons
  • Begin the filing of electrons in these respective domains as per the total number of electrons available in the valence shell
  • In Lewis dot structure it’s only the valence shell electrons which could be represented

After the Lewis dot structure is drawn the participating atoms valence shell electrons are checked as well and paired accordingly. Once the respective pairing is over the balance number of electrons are checked and lone pairs are isolated. These lone pairs then are crossed checked with the total number of bonded pairs to make sure no bonding / sharing is incomplete. Most of the time the larger of the two or more participating atoms of the respective elements remain with lone pairs as they are not shared by available electrons from the participating atoms valence shell.

The difference in repulsion or repulsive forces between electron pairs shows that when lone pairs are present, the geometric changes in molecules are obvious. In case of ammonia NH3, the presence of five electrons in valence shell is observed in case of Nitrogen. Nitrogen has 7 electrons in all and hence the electron configuration is 2, 5. The 5 electrons are arranged in accordance of the Lewis dot structure rule.

Hydrogen whereas has one electron in valence orbit and hence as per the pairing rule only three of Nitrogen’s valence electrons get paired. Out of 5 valence electrons of Nitrogen only three electrons get pairing and the rest two remain as lone pairing. 

Ammonia Valence Electron

Moreover, we need to understand it’s the bonded pairs which are usually found all across the central atom and it’s only in specific electron domains we could see the lone pair electrons.

The more the lone pairs the binding with hydrogen with nitrogen gets squeezed and the bonding angles also gets affected and it reduces from the original geometry. Molecules always obey octet rule except Helium and Hydrogen which follows duplet. These molecules have participating atoms is surrounded by eight bonding or lone pair electrons as it is closely associated with stability configuration.

Ammonia, methane and water are the best examples of such lewis forms where the ground state configuration and hybridized state exist. 
Nitrogen
 
Nitrogen has two electrons in the 2s and another three in 2px, 2py and 2pz respectively to bond with each of Hydrogen’s 1s.

Lone Pairs of Electrons in Water

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In the molecule of water, the participating atoms are one from Oxygen and another two from Hydrogen. The electron configuration of Oxygen is 2, 6

As far the orbital distribution is concerned Oxygen has $1s^{2}$, $2s^{2}$, $2p^{4}$ with six electrons in valence shell. The two hydrogens are separately bonded with the central atom Oxygen in covalent manner.

The bonding electrons in this molecule are specific with oxygen and hydrogen on either side which leaves two pairs of unused electron pairs. These unused pairs which constitute oxygen octet also termed as lone pairs or non-bonded electrons is seen as not having any involvement in covalent bonding. 
Lone Pairs of Electrons in Water

Lone Pair Vs Bonding Pair

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As far the difference between the lone pair and bonded pairs are concerned the comparison and common ground for both types of pair is given as below.

 Lone pair   Bonding pair 
 Observed only in covalent bonded molecules and ions   Can be a common ground for hydrogen and covalent bonding
 The greater charge density of electrons  Relatively lesser charge density of electrons than lone pairs
 These pairs remain unused even after the bonding is over   These pairs are the ones which are shared between the participating atoms of the elements in a molecule
 These are usually termed as non-bonding electron pairs  These are considered as bonding electron pair usually in valence sub orbitals
 These are responsible for the overall repulsion over the bonded pairs and hence give the final geometrical shape to the molecules or ions   These are not directly involved for the final geometrical shapes to the molecules and ions
 For each lone pairs presence in a molecule the normal angle squeezes by almost 3 degrees  No such phenomenon is observed for bonded pair
 To identify the lone pairs the lewis electron dot structure is very much necessary  The bonded pairs are identified on the basis of lewis electron dot structure
 The largest repulsion is observed between lone pair- lone pair or multiple or shorter bonds, or any bonds which are polarized towards the central atom  The smallest repulsion is observed when the bonding takes place and has single longer bonds, the bond pairs are far away from central atom causes least repulsion

Why Do Lone Pairs Repel More?

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The lone pairs repel more because these are free electron cloud and not bounded by any force of attraction which usually is seen in bonded pairs. The absence of any binding force these free cloud of electrons carrying same negative charge begin to repel each other and hence the repulsive force.

Although each participating electrons in the formation of molecule carry same charge but the absence of any kind of sharing in between subjects these electrons push each other out from the electron domains.

These lone pairs are observed to have more repulsion force compared to the bonded pairs and finally reach out to give a squeezed bonded angles than its normal geometrical domains.
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