Electronegativity and the Pauling Scale

As we have seen, there are three types of primary bond - covalent, ionic and metallic - all of which involve the exchange or sharing of electrons. At one end of the spectrum, we have ionic bonding, in which one atom essentially gives up one or more of its valence electrons to another. Then we have covalent bonding, in which valence alectrons are "shared" between two atoms. Finally, we have metallic bonding, in which the electrons don’t belong to any particular atom, but act as a kind of "glue", holding the atoms of the metallic substance together. The factor that determines the nature of the primary bonds to be found in any elementary substance or chemical compound is something called electronegativity.

Electronegativity (or electron affinity) is the degree to which an atom of a particular element can attract a shared electron when participating in a chemical bond. Electronegativity is a dimensionless property, and values are usually assigned according to the Pauling Scale. This scale gives flourine - the most electronegative element - a value of 3.98. All other elements are assigned a value depending on their electronegativity relative to flourine. The version of the periodic table shown below gives the Pauling Scale electronegativy values for each element.


The Periodic Table - Pauling Scale Electronegativity
 123456 789101112 131415161718
1 H
2.20
He
 
2 Li
0.98
Be
1.57
B
2.04
C
2.55
N
3.04
O
3.44
F
3.98
Ne
 
3 Na
0.93
Mg
1.31

Al
1.61
Si
1.90
P
2.19
S
2.58
Cl
3.16
Ar
 
4 K
0.82
Ca
1.00
Sc
1.36
Ti
1.54
V
1.63
Cr
1.66
Mn
1.55
Fe
1.83
Co
1.88
Ni
1.91
Cu
1.90
Zn
1.65
Ga
1.81
Ge
2.01
As
2.18
Se
2.55
Br
2.96
Kr
3.00
5 Rb
0.82
Sr
0.95
Y
1.22
Zr
1.33
Nb
1.60
Mo
2.16
Tc
1.90
Ru
2.20
Rh
2.28
Pd
2.20
Ag
1.93
Cd
1.69
Ln
1.78
Sn
1.96
Sb
2.05
Te
2.10
I
2.66
Xe
2.60
6 Cs
0.79
Ba
0.89
La-Lu
 
Hf
1.30
Ta
1.50
W
2.36
Re
1.90
Os
2.20
Ir
2.20
Pt
2.28
Au
2.54
Hg
2.00
Tl
1.62
Pb
1.87
Bi
2.02
Po
2.00
At
2.20
Rn
2.20
7 Fr
0.79
Ra
0.90
Ac-Lr
 
Rf
 
Db
 
Sg
 
Bh
 
Hs
 
Mt
 
Ds
 
Rg
 
Cn
 
Nh
 
Fl
 
Mc
 
Lv
 
Ts
 
Og
 
Down arrow
La
1.10
Ce
1.12
Pr
1.13
Nd
1.14
Pm
-
Sm
1.17
Eu
-
Gd
1.20
Tb
1.10
Dy
1.22
Ho
1.23
Er
1.24
Tm
1.25
Yb
-
Lu
1.27
Ac
1.10
Th
1.30
Pa
1.5
U
1.38
Np
1.36
Pu
2.28
Am
1.13
Cm
1.28
Bk
1.30
Cf
1.30
Es
1.30
Fm
1.30
Md
1.30
No
1.30
Lr
1.30

We stated above that electronegativity is responsible for determining the type of bond we can expect to find in a given elemental substance or chemical compound, but it would be more accurate to say that it is the difference in electronegativity between two atoms that will determine the kind of bond that forms between them. If you look at the electronegativity values of the main group elements in the periodic table, you will see that the values increase from left to right within each period.

The difference in electronegativity also reflects the difference in the number of valence electrons. For example, Na (sodium) in Group 1 has an electronegativity of 0.93 and one valence electron, whereas Cl (chlorine) in Group 17 has an electronegativity of 3.16 and seven valence electrons. When a sodium atom combines with a chlorine atom to form a sodium chloride formula unit (NaCl), the difference in electronegativity (2.23) is sufficiently great to effect the complete transfer of the sodium atom's lone valence electron to the chlorine atom, leaving both atoms with a full outer electron shell, and creating an ionic bond.


A sodium atom and a chlorine atom form an ionic bond

A sodium atom and a chlorine atom form an ionic bond


At the other end of the scale, we have metallic bonding, which occurs only in metallic elements and metal alloys. Metallic substances tend to have relatively low electronegativities, allowing valence electrons to break away from their parent atom and preventing them from being captured by a neighbouring atom. The result is a kind of "soup" consisting of positively charged metal ions surrounded by delocalised and negatively charged ions.

In covalent bonds, the difference in electronegativity between the bonded atoms is usually far less pronounced than is the case for ionic bonds. We say that the valence electrons are "shared" between two atoms. Even so, if one of the atoms is more electronegative than the other, the valence electrons will spend more time around the more electronegative atom. So how do we differentiate between ionic and covalent bonds? As a general guide, ionic bonds are characterised by a difference in electronegativity of 2.0 or more on the Pauling scale, while a difference of less than 2.0 indicates a covalent bond.

A carbon atom has four valence electrons and so requires four additional electrons to complete its outer electron shell. In a methane (CH4 ) molecule, the carbon atom achieves this by forming covalent bonds with four hydrogen atoms, as shown below. In each covalent bond, two valence electrons are “shared” between the carbon atom and one of the hydrogen atoms. Hydrogen and carbon have similar electronegativities - 2.20 and 2.35 respectively - so the difference in electronegativity is only 0.35, well under the 2.0 threshold.


Covalent bonds in a methane molecule

Covalent bonds in a methane molecule


Covalent bonds can be either polar or non-polar. In a purely non-polar covalent bond, the shared electrons can be thought of as residing half way between the bonded atoms. This requires the atoms in question to have the same electronegativity, so purely non-polar covalent bonds only occur in simple molecules involving atoms of the same type - for example hydrogen (H2 ), oxygen (O2 ), or chlorine (Cl2 ) molecules.

A polar covalent bond occurs when one of the bonded atoms has a greater electronegativity, and thus a greater attraction attraction for the shared electrons, than the other bonded atom. The result is that one end of the bond achieves a greater electron density and becomes slightly negative, while the other end, with a lower electron density, becomes slightly positive. The difference in electron desnsity will be proportional to the difference in electronegativity between the two bonded atoms.

Hydrogen chloride (HCl) is a colourless gas at standard temperature and pressure. HCl provides a good example of polar covalent bonding because each HCl molecule consists of a hydrogen atom covalently bonded to a chlorine atom. The hydrogen atom shares its single valence atom with the chlorine atom, giving both atoms a complete outer electron shell, but the difference in electronegativity (2.20 for hydrogen against 3.16 for chlorine) means that the valence electrons spend more time around the chlorine atom than they do around the hydrogen atom, although the difference in electronegativity (0.96) is still well within bounds for a covalent bond.


Polar covalent bonding in a hydrogen chloride (HCl) molecule

Polar covalent bonding in a hydrogen chloride (HCl) molecule


We have already mentioned that purely non-polar covalent bonds occur only in molecules in which the atoms involved are all of the same type - for example, the diatomic molecules O2  or Cl2 . Essentially, all covalent bonds formed between atoms of different types are polar in nature. The degree of polarity, however, will depend on the difference in negativity between the atoms participating in the bond, which will determine how the electrons are distributed between the two bonds. Some covalent bonds are only nominally polar in nature, whilst other bonds are strongly polar.