Brown's Chemistry Revision
Extra Notes on chemical bonding for advanced A level
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)
CHEMICAL BONDING INDEX
Electronegativity is defined as the power of an
atom to attract the pair of electrons in a covalent bond situation.
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
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.
Electronegativity is the power of an atom to attract electron
charge from another atom it is covalently bonded to. Some
values of electronegativity are quoted below.
speaking electronegativity increases from left to right across a
period of the periodic table and decreases down a group of the
2. Bond polarity and
type of chemical bonding
Three situations are considered
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)
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+
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:
a picture of HCl
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
e.g. in tetrahedral shaped tetrachloromethane CCl4,
is cancelled out by the 4 symmetrically arranged Clδ–
but in chloromethane
there is a permanent dipole due to the polar C-Cl bond (Cδ+-Clδ–).
Carbon dioxide O=C=O (δ–O=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.
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.
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
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–
(2.8), these give rise to giant ionic lattices with a high percentage of
Details of ionic compound examples
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
(i) essentially ionic with some covalent
(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.
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
It basically amounts to a measure of
a positive ion's ability in attracting the electrons clouds in
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
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
ΔHLE: the lattice
enthalpy (energy required to convert one mole of the crystal lattice into
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!
|relative polarising power of cation
(electric charge/ionic radius)
||relative polarisability of anion
||Pauling electronegativity difference
||Na 0.9, Cl 3.0
Δ elec. = 2.1
||Na 1.0, Cl 3.0
Δ elec. = 2.0
||larger I > Cl
||Na 0.9, I 2.5
Δ elec. = 1.5
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
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
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
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
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.
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
CHEMICAL BONDING INDEX
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