A Level Chemistry: More on electronegativity, bond polarity, polarisation, atomic/ionic radii

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Extra Notes on chemical bonding for advanced A level chemistry students

 Extra notes on electronegativity, bond polarity, polarisation by cations, polarisability of anions, atomic and ionic radii and the influence of these factors on the type of chemical bonding - bond character

 (also notes on dipoles)



Electronegativity is defined as the power of an atom to attract the pair of electrons in a covalent bond situation.

The electron distribution in a covalent bond between elements with different electronegativities will be unsymmetrical.

This produces a polar covalent bond, and may cause a molecule to have a permanent dipole which is indicated by partial charges (δ+ and δ) to show that a bond is polar.

You should be able to recognise when a molecule is likely to be polar and explain why some molecules with polar bonds do not have a permanent dipole.

1. Electronegativity

Electronegativity is the power of an atom to attract electron charge from another atom it is covalently bonded to. Some Pauling values of electronegativity are quoted below.

element Na Mg Al Mn Fe H Si P C S I Br Cl N O F
electronegativity 0.9 1.2 1.5 1.5 1.8 2.1 1.8 2.1 2.5 2.5 2.5 2.8 3.0 3.0 3.5 4.0

Generally speaking electronegativity increases from left to right across a period of the periodic table and decreases down a group of the periodic table.

2. Bond polarity and type of chemical bonding

Three situations are considered

(a) Non-polar covalent bond

 If the two atoms of a covalent bond have the same or very similar electronegativity, the bond is referred to as non-polar.

Examples of non-polar bonds (electronegativity difference): H-H (0), C-C (0), C-H (0.4)

(b) Polar covalent bond

If two atoms of a covalent bond have an appreciable difference in electronegativity, the result is a polar bond AND the molecule usually has a permanent dipole due to a permanent partial charge separation. This asymmetry in the electron cloud distribution is shown a delta + and a delta - (δ+ and δ).

Examples of polar bonds (electronegativity difference): H-Cl (0.9), O-H (1.4), N-H (0.9), C-Cl (0.5)

The polar bonds in polar molecules with a permanent dipole can be represented as:

Hδ+-Clδ, δO-Hδ+, δN-Hδ+, Cδ+-Cl δ

a picture of HCl , water

Note that in some circumstances a molecule with polar bonds may NOT have a permanent dipole because the individual dipoles cancel each other out because of the symmetry of the molecule.

e.g. in tetrahedral shaped tetrachloromethane CCl4, the Cδ4+ is cancelled out by the 4 symmetrically arranged Clδ– atoms.

but in chloromethane (c) doc b there is a permanent dipole due to the polar C-Cl bond (Cδ+-Clδ–).

Carbon dioxide O=C=O (δO=Cδ2+=Oδ) has two polar bonds, but the molecule is linear and the two dipoles cancel each other out.

Simple laboratory experiment to detect permanent polarity in a molecule

By running jets of various liquids from burettes past an electrostatically charged plastic rod you find highly polar molecules like water are quite dramatically deflected due to the permanent dipole but non-polar molecules like hexane show no deflection at all.

The plastic rod is charged by rubbing with a wool or line cloth. A 'bar' of poly(ethene) is quite good for this deflection experiment.

Dipole moment

Its unlikely you need to know about this unit, but some advanced pre–university courses do mention dipole moments.

Dipole moment (μ) is the measure of net molecular polarity, and is the magnitude of the charge (Q) at either end of the molecular dipole multiplied by the distance between the charges (d).

μ = Q x d

Dipole moments are quoted in D, debye units, 1D = 3.34 x 10–30Cm

This is a non–SI unit. C = coulomb, m = metre.

Dipole moments informs about the charge separation in a molecule. The greater the difference in electronegativities of the bonded atoms, the larger the dipole moment. Also, the larger the charge separation, the greater the dipole moment.

In both the situations of (a) and (b) described above you are looking at the formation of covalent molecules.

Details of covalent element/compound molecule examples

(c) doc b(c) Ionic bond

If two atoms constituting a bond have a big difference in electronegativity then an ionic bond may be formed.

In other words, rather than sharing electrons to form a covalent bond, electron transfer is completed giving rise to positive and negative ions whose electrostatic attraction constitutes the ionic bond.

Examples of ionic bonds (electronegativity difference): Na+Cl (2.1), Mg2+(F)2 (2.8), these give rise to giant ionic lattices with a high percentage of ionic character.

Details of ionic compound examples

3. Bond character complications If only it was a simple as a, b, c above!

Even before considering other factors, there are going to be cases where the electronegativity difference will be somewhere between that of a very covalent compound and a very ionic compound.

This will inevitably lead to compounds which are:

(i) essentially ionic with some covalent character or

(ii) essentially covalent compounds with some ionic character.

I'll illustrate some cases of (i) ionic compound with some covalent character, via the data table below, but first the various terms used are defined and explained below.

Terms used:

ionic radius: the radius of an that encloses the majority of the electron clouds.

relative polarising power of a cation: The ability of a cation to distort an anion is known as its polarising (polarizing) power.

It basically amounts to a measure of a positive ion's ability in attracting the electrons clouds in neighbouring ions.

A simple measure of polarising power is to divide the electrical charge by the ionic radius (c+/r)

The polarising (polarizing) power of a cation increases with increase in charge and decrease in radius.

These trends effectively increase the intensity of the attracting electric field on neighbouring electron clouds.

relative polarisability of an anion: the tendency of the anion to become polarised (polarized) by a cation is known as its polarisability (polarizability).

The larger the anion and the greater its negative charge, more easily its outer electron clouds are polarised - more easily attracted-distorted towards the cation.

ΔHmelt: the enthalpy of fusion (melting)

ΔHLE: the lattice enthalpy (energy required to convert one mole of the crystal lattice into separated gaseous ions at 298K, 1 atm pressure).

Pauling electronegativity: is a commonly used measure of the ability of an atom to attract the electron cloud of the other atom in a chemical bond.

The greater the difference in electronegativity between the two atoms in a chemical bond, the greater the ionic character, though this is only one factor that influences the character of the bond - read on!

Compound formula mpt/oC ionic radii


relative polarising power of cation (electric charge/ionic radius) relative polarisability of anion Pauling electronegativity difference ΔHmelt




sodium chloride NaCl 801 Na+ (0.095)

Cl- (0.181)

10.5   Na 0.9, Cl 3.0

Δ elec. = 2.1

28.9 766
lithium chloride LiCl 617 Li+ (0.060)

Cl- (0.181)

16.7   Na 1.0, Cl 3.0

Δ elec. = 2.0

13.4 833
lithium iodide LiI 447 Li+(0.060)

I- (0.216)

16.7 larger I > Cl Na 0.9, I 2.5

Δ elec. = 1.5

6.3 744

In an exam question, you might be provided with data, in the absence of data you need to be able to explain the situation on a descriptive conceptual basis!

(i) Cases of ionic compounds with varying covalent character

Sodium chloride: An appreciable electronegativity difference leads to a well defined giant ionic lattice which has a relatively high melting point and enthalpy of fusion.

It is a good conductor when molten and in aqueous solution where in both cases the ions are free to move and carry an electrical current.

It is a classic ionic compound which is considered to have a high degree of ionic character and very little covalent character in the lattice bonding.

Lithium chloride: Now for the complications! On face value you might expect lithium chloride to be quite similar in properties to sodium chloride. Since the lithium ion has a smaller ionic radius than sodium, the attractive force between the 'closer' ions should be greater, giving LiCl a greater melting point and greater enthalpy of fusion than NaCl. BUT this is not the case. In fact lithium chloride has a significantly lower melting point and enthalpy of fusion compared to sodium chloride. So, what's going on!

We need to consider other factors affecting the situation!

The relatively smaller radius of the lithium ion compared to the sodium ion gives this cation a much greater polarising power (charge/radius).

The lithium ion polarises the chloride ion much more than the sodium ion. The lithium ion attracts the electron clouds of the chloride ion towards itself reducing the ionic character of the bond and inducing some covalent character to the bond - inducing a degree of electron sharing (but take care how you use this phrase, it does not necessarily involve orbital overlap as in a covalent bond!).

It should be added that molten lithium chloride is a good conductor, though not as good as sodium chloride, but as far as I know aqueous lithium chloride is as good as a conductor as brine!

Lithium iodide: Following on from the LiCl discussion, in lithium iodide we have an even more polarisable anion in iodide.

So we can compare lithium chloride with lithium iodide, but this time we compare the polarisability of the anion.

Although the difference in electronegativity is still appreciable, lithium iodide shows even more covalent character than lithium chloride.

The larger iodide ion, compared to the chloride ion, is more easily polarised than the chloride anion, particularly bearing in mind the relatively small radius of the lithium ion.

You might expect the melting point and enthalpy of fusion to be a bit lower than that of lithium chloride (larger inter-ionic distance) but the increased polarisation of the iodide ion produces quite a lowering of these values compared to what might be expected for a 'pure' ionic crystal of lithium iodide.

The curious case of aluminium chloride

Aluminium chloride is very ambiguous! X-ray diffraction shows that solid, AlCl3, 'appears' to consist of an ionic lattice of Al3+ ions, each surrounded by six Cl ions, BUT on heating, at about 180oC, the thermal kinetic energy of vibration of the ions in the lattice is sufficient to cause it break down and sublimation takes place (s ==> g) to form the readily vapourised covalent dimer molecule, Al2Cl6.

This suggests the highly polarising aluminium ion affects the chloride ion, distorting the anion's electron clouds towards it and inducing considerable covalent character even in the crystal lattice. At 180oC there is insufficient ionic character to maintain a strong ionic lattice, hence ionic bonding, and the bonding becomes highly covalent and discrete covalent molecules are formed.

The electrical conductivity of molten aluminium chloride (under high pressure to prevent sublimation) is much less than that you would expect for an ionic compound like molten sodium chloride, again showing a lack of 'full' ionic character.

(ii) Cases of covalent compounds showing some ionic character

The O-H bond in water is a highly polar molecule (δ+H–Oδ2––Hδ+). The relatively large difference in electronegativity between hydrogen and oxygen (H 2.1, O 3.5, Δ elec. = 1.4)  introducing some ionic character into the bond, and a tiny fraction of molecules can undergo self-ionisation giving water a tiny, but real and measurable electrical conductivity.

2H2O(l)  H3O+(aq)  + OH(aq)

Note the difference in electronegativity is only slightly smaller than in lithium iodide, but the tiny proton has quite a high polarising power effect even on the very electronegative oxygen, so covalency dominates!



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