In depth bonding essay edexcel 2008 chemistry

A more in depth version of prior essay on all types of bonding needed for edexcel syllabus from 2008 chemistry

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Naomi Cockrill
Intramolecular and Intermolecular Bonding
Bonding, whether intramolecular or intermolecular, is present in all chemical compounds or
molecules. Intramolecular bonding comprises a broad spectrum of bonding types, of which ionic and
covalent bonding are the two extremes. Often, intramolecular bonds affect which intermolecular forces are
present and together these account for many of the physical properties of chemicals.
Ionic (electrovalent) bonding is a result of the electrostatic attraction between cations (positive ions)
and anions (negative ions) following the transferring of some of the cation's outer shell electrons (electron
loss) to the outer shell of the anion (electron gain) in order maximise attraction between the atoms. The
greater the charge on the ions and the closer they can get (e.g. smaller the cation), the greater the attraction
and the stronger the bond. These two factors give a cation its charge density (charge / volume). The energy
released from the formation of an ionic bond has to be greater than that needed for its formation (must be
energetically favourable) so that the compound can move to a lower energy state. Therefore, it follows that
the stronger the ionic bond, the more energy is released from its formation. This is how many cations can
afford to donate more than one electron (giving them a greater charge density), which requires a lot of
energy (first and second ionisation energies) because the bond (or bonds) formed are stronger and so
release more energy, compensating for the additional input of energy.
An example of this type of bonding is found in the simple ionic compound sodium chloride and can be
represented by a `dot-and-cross' diagram (Figure 1.) In this example, the
charges on the cation and anion are opposite and equal so a sodium chloride
compound is made up of a 1:1 ratio of sodium and chloride ions. However, in
magnesium chloride, two chloride ions are needed to balance the 2+ charge
of the magnesium ion. Although Figure 1 shows that both anion and cation
achieve their lowest energy state by achieving a full outer shell (ns2np6), this is
not the case for all ionic compounds. For example, the transition metals can
form ionic bonds with non-metals yet they do not have this outer shell
structure as the copper 2+ ion, with an electron structure of [Ar]3d9,
demonstrates. Here, not only are the `p' orbitals empty but also this ion has
more than eight electrons in its outer shell. This shows that ionic bonds form
do in order to achieve the lowest energy state possible and not necessarily to get a full outer shell of
The theory of ionic bonding has been accepted due to the extensive supportive evidence. For
example, all ionic compounds giant or simple) have very high boiling and melting points, so a lot of energy is
needed to overcome the strong attraction between the ions, which corresponds to the ionic model as
positive ions are strongly attracted to negative ions. Moreover, the melting point of magnesium oxide
(double charges) is twice that of sodium chloride (single charges), showing how closely the strength of the
attraction between the ions relates to the melting and boiling points. Secondly, many ionic lattices are
soluble in water (which has permanent dipole ­ permanent dipole interactions) but not in non-polar solvents
(which only have London forces). This is because a substance will only usually dissolve if the solvent ­ solute
bonds formed are equally or more strong than the bonds that are broken. In the case of magnesium oxide,
the cation is attracted to the lone pair of electrons on the oxygen molecule and hydrogen bonding, which is
stronger than the London forces present in the ionic crystal, could occur between the oxide ions and
hydrogen atoms. Some ionic lattices cannot be hydrated because their intermolecular forces are stronger
than those that would from with water. (NB. There is more on solubility and intermolecular forces later.)
Another property of ionic compounds is that they do no conduct electricity when solid but do when molten
or in solution( as seen in electrolysis as the ions are free to move). This is especially seen in the migration of
ions of wet filter paper. When a current is passed through the paper, the negative ions and positive ions are
attracted to oppositely charged ends and these ions are often coloured, so the appearance of colours,
indicating ions moved due to attraction, supports the ionic theory. Finally, ionic lattices ( their structure also
supports the ionic theory ­ later explained) are very brittle because once a force is applied, the `layers' of

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Naomi Cockrill
atoms slide past one another and as soon as the ions are aligned with ions of the same charge, repulsion
causes the layers to break apart.
In addition to properties of ionic compounds, other evidence can be seen
by analysing the electron density diagrams of a highly ionic compound such
as sodium chloride ( figure 2.) The circular lines represent the likelihood of
electrons being in that location, so, the closer the lines the higher the
electron density.…read more

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Naomi Cockrill
Haber values to the theoretical values, showing that these compounds fit the `purely ionic' model. However,
the Born-Haber values for magnesium halides are much lower than the theoretical values, indicating the
bonding is stronger than the `purely ionic' model accounts for and that the charge is not evenly distributed.
This means that there is some covalent character in the ionic bonds, which perfectly demonstrates the idea
of a spectrum of bonding. The uneven distribution of electrons is known as polarisation.…read more

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Naomi Cockrill
forces have to be overcome and so are generally liquids and gases at room temperature. On an electron
density map, there is merging between the two atoms, demonstrating the sharing of electrons. Giant
covalent molecules, like silicon dioxide or graphite, contain a network of covalent bonds (as silicon and
carbon can form four bonds each) and so have very high melting points, being generally solid at room
temperature.…read more

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the ionic bond is polarised and if the difference is only small, then it is a covalent bond that is polar ( has a
negative and positive end.) This is the case for the bonds in carbon dioxide as the oxygen molecules are
much more electronegative than the carbon.
Although all covalent bonds ( except ones between two of the same atom) are at least slightly polar, not all
covalent molecules are polar.…read more

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Naomi Cockrill
molecules are aligned and attracted to one another. However, these instantaneous-induced dipoles are only
temporary but, on average, there is an attraction across the whole molecule. As London forces depend on
the presence of electrons, the bigger the molecule, the more electrons there are, the greater the electrical
imbalance that can be caused, the greater the dipole and the stronger the force.…read more


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