9.11 Chemical Bonding in
Ionic bonding * (2) Covalent
bonding * (3)
Oxyanions and ligands
bonding in halide salts =
Factors affecting bond formation
differences in electronegativity between metals and non–metal
halogens often results ionic bonding
Ex 1 ONE
combines with ONE
electron configurations [1s22s22p63s1]
+ [1s22s22p63s23p5] ==>
[1s22s22p6]=[Ne] + [1s22s22p63s23p6]=[Ar]
showing the electron transfer
from the electropositive metal to the electronegative non–metal and the ensuing
Ex 2 ONE
combines with TWO
electron configurations [1s22s22p63s2]
+ 2[1s22s22p63s23p5] ==>
[1s22s22p6]=[Ne] + 2[1s22s22p63s23p6]=[Ar]
Ex 3 ONE
combines with THREE
+ 3[1s22s22p5] ==> [1s22s22p6]=[Ne]
- The greater the
electronegativity between two bonded atoms, the greater the likelihood of an
ionic bond being formed.
- However, the electronegativity of an element
itself is not the only factor in deciding bond character!
- The different results of the 'tug of war'
between two positive nuclei acting on the intermediate bonding electrons
produces a range of bond character from complete electron transfer in ionic bond
formation (e.g. M+ and X–), to a highly polar covalent
bond (Mδ+–Xδ–) of partially charged atoms and at
the other extreme a virtually non–polar bond (X–Y) of two atoms, neither of
which carries a significant partial charge.
- Electronegativity, the power of an element to
attract bonding electrons towards it in a bonding situation, is just one, albeit
important, factor in deciding the outcome of the character of an individual
- The stronger the polarising power of the
cation and the higher the polarizability of the anion the more
covalent character is expected in a bond.
- If a cation has appreciable polarising power to
draw bonding electron clouds towards it OR the bonding electron clouds of an
anion are attracted towards the cation, then covalent bonding character is more
- Polarising power for cations is very much
a case of increasing with increased ion charge/ionic radius.
- So, the smaller the ionic radius or the bigger
the positive charge, the greater the polarising power of the cation.
- e.g. in terms of polarising power Al3+
> Mg2+ >Na+ for the series of Period 3 positive ions
where you have both coincident decreasing radii and increasing charge.
- For Group ions, polarising power will decrease
down the group with increasing ionic radius and constant charge.
- therefore in polarising power Li+
> Na+ >K+ for Group 1 Alkali Metals etc.
- or Be2+ > Mg2+ > Ca2+
for Group 2 Alkaline Earth Metals etc. (see also s–block
- Examples of the outcome of this factor are
- increasing ionic character NaCl > MgCl2
- The latter is just ionic in the lattice, but
vaporises to a covalent dimer.
- increasing ionic character KCl > NaCl > LiCl
- As the cation gets larger for same charge its
polarising power diminishes.
- Change in oxidation state can also change the
bond character significantly. Iron(II) chloride is essentially ionic in
character and iron(III) chloride is basically covalent because the polarising
power of the smaller and more highly charged iron(III) ion.
- For anions, the larger the ionic radius and
the greater its charge, the more polarisable it is.
- So in terms of polarizability I– >
Br– > Cl– > F– (for halide ions for
constant charge and decreasing radius)
- or polarizability of Si4– > P3–
> S2– > Cl– (for a series of Period 3 anions of
decreasing charge and decreasing ionic radius)
- therefore you expect for ...
- A series of Group 2 halides the ionic
character CaCl2 > MgCl2 > BeCl2
- Calcium chloride is essentially ionic and
beryllium chloride is essentially covalent.
- The series of Period 3 chlorides the ionic
character be NaCl > Na2S > Na3P > SiCl4
- In fact sodium chloride is very ionic high
melting lattice and silicon(IV) chloride is very covalent low boiling liquid!
- A series of Group 1 halide salts the ionic
character trend is KF > KCl > KBr > KI
- Potassium iodide is essentially ionic, but its
'partial' covalent character is shown by the fact that it dissolves in polar
solvents like propanone (acetone) whereas highly ionic potassium chloride is
~insoluble in polar organic solvents.
- What this set of paragraphs illustrates is a
much wider and deeper approach to electronegativity than e.g. the
electronegativity number quoted on the Pauling scale.
Properties of ionic compounds
diagram above shows a typical of the giant ionic crystal structure of ionic
compounds like sodium chloride and magnesium oxide.
- The alternate positive and negative ions in an
ionic solid are arranged in an orderly way in a
giant ionic lattice structure.
- The ionic bond is the strong electrical attraction between the positive and negative ions next to each other in the lattice.
- The bonding extends throughout the crystal
in all directions.
- Halide salts are typical ionic compounds.
- This strong bonding force makes the structure hard (if brittle) and have high melting and boiling points,
so they are not very volatile!
- A relatively large amount of energy is
needed to melt or boil ionic compounds. Energy changes for the physical changes of state
of melting and boiling for a range of differently bonded substances are compared in a section of
the Energetics Notes.
- The bigger the charges on the ions the stronger
the bonding attraction e.g. magnesium oxide Mg2+O2–
has a higher melting point than sodium chloride Na+Cl–.
- Unlike covalent molecules, ALL ionic compounds are crystalline solids at room temperature.
- They are hard but brittle, when stressed
the bonds are broken along planes of ions which shear away. They are NOT
malleable like metals.
- Many ionic compounds are soluble in water,
all, so don't make assumptions.
- Salts can dissolve in water because the ions can
separate and become surrounded by water molecules which weakly bond to the ions.
- This reduces the attractive forces between the ions, preventing the crystal
structure to exist. Evaporating the water from a salt solution will eventually
allow the ionic crystal lattice to reform.
- The solid crystals DO NOT conduct electricity because the ions are not free to move to carry an electric current. However, if the ionic compound is melted or dissolved in water, the liquid will now conduct electricity, as the ion particles are now free.
- See also
Notes on ionic bonds and ionic compounds
bonding in small molecules of the halogens and their compounds
differences in electronegativity between two non–metallic elements
usually results in covalent bonding
chlorine atoms (2.8.7) form the molecule of the element chlorine Cl2
combine to form
where both atoms have a pseudo argon structure of 8 outer electrons around each atom.
All the other halogens would be similar e.g. F2, Br2 and I2
etc. Valency of halogens like chlorine is 1 here.
atom of hydrogen (1) combines with one atom of chlorine (2.8.7) to form the molecule of the
hydrogen chloride HCl
combine to form
where hydrogen is electronically like helium and chlorine like argon. All the
other hydrogen halides will be similar e.g. hydrogen fluoride HF, hydrogen
bromide HBr and hydrogen iodide HI. These are polar molecules due to the greater
electronegativity of the halogen compared to hydrogen i.e. Hδ+–Xδ–
Hydrogen chloride gas is a true covalent substance consisting of small
HCl molecules. If the gas is dissolved in a hydrocarbon solvent like hexane or
methylbenzene it remains as HCl molecules and because there are no ions present,
the solution does not conduct electricity. However, if hydrogen chloride gas is
dissolved in water, things are very different and the HCl molecules split
into ions. Hydrochloric acid is formed which consists of a solution
of hydrogen ions (H+) and chloride ions (Cl–).
The solution then conducts electricity and passage of a d.c. current causes
electrolysis to take place forming hydrogen and chlorine.
BeCl2 (X = Be, Q = F, Cl, X = Be)
Linear molecules with a Q–X–Q bond
angle of 180o.
Group 7 F, Cl etc. X
= Group 5 N, P etc.
Pyramidal or trigonal pyramid shaped molecules
with a Q–X–Q bond angle of ~180o
COCl2 is a good example
to practice your dot and cross diagram skills. A trigonal planar molecule with a
Cl–C=O bond angle of ~120o
Ex 9 gaseous Group 3 halides
Q = F, Cl and X =
B, Al for F
Trigonal planar molecules with a
Q–X–Q bond angle of 120o
chloromethane, tetrahedral shape, bond angles of H–C–H and
H–C–Cl are all ~109o
Q = Cl, X = C, Si, Ge etc. in group 4, tetrahedral, bond angles 109o
e.g. PF5 X=P, Q=F,
trigonal bipyramid, bond angles, 180, 120 and 90o
e.g. SF6 X=S, Q=F,
octahedral shape, bond angles 90 and 180o
chloride is a curious substance in its behaviour. The solid,
AlCl3, consists of an ionic lattice
of Al3+ aluminium 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). In doing so the
co–ordination number of the aluminium changes from six to
four to form the readily vapourised covalent dimer
Al2Cl6, shown above.
discussion on electronegativity and cation
polarising power in the ionic bonding section.
Covalent Bonding – small simple molecules and properties
Properties of small covalent molecules
The electrical forces of attraction, that
chemical bond, between atoms in a molecule are usually very strong,
so, most covalent molecules do not change
chemically on moderate heating.
e.g. although a covalent molecule
like iodine, I2, is readily vapourised on heating, it does NOT
break up into iodine atoms I. The I–I covalent bond is strong enough to
withstand the heating and the purple vapour still consists of the same I2
molecules as the dark coloured solid is made up of.
So why the ease of vaporisation on
These weak electrical attractions are known as intermolecular
forces and in particular instantaneous dipole – induced dipole forces.
These intermolecular forces are readily weakened further on
The effect of absorbing heat energy results in increased the
thermal vibration of the molecules in the solid which weakens the intermolecular forces.
In liquids the increase in the average particle kinetic energy makes it
easier for molecules to overcome the intermolecular forces and change into a
gas or vapour.
Consequently, small covalent molecules
tend to be volatile liquids with low boiling points, so easily vapourised, or low melting point solids.
On heating the inter–molecular forces are
easily overcome with the increased kinetic energy of the particles giving
the material a low melting or boiling point and a
relatively small amount of energy is needed to effect these state
for the physical changes of state of melting and boiling for a range of
differently bonded substances are compared in a section of
This contrasts with the high
melting points of
giant covalent structures
with their strong 3D network.
Note: The weak electrical attractive
forces between molecules, the so called intermolecular forces
should be clearly distinguished between the strong covalent bonding
between atoms in molecules (small or giant), and these are sometimes
referred to as intramolecular forces (i.e. internal to the
Covalent structures are
usually poor conductors of electricity because there are
no free electrons or ions in any state to carry electric charge.
Most small molecules will dissolve in
some solvent to form a
'Like dissolves like' is a
useful, if limited, 'rule of thumb'. Chloromethane and other
halogenoalkanes will dissolve in hydrocarbon solvents but not highly
This again contrasts with
covalent structures where the strong bond network stops solvent
molecules interacting with the particles making up the material.
Covalent Bonding – small simple molecules and properties
Other examples of bonding in halogen
compounds – oxyanions and ligands
Sketches of the dot & cross diagrams for
the following anions: chlorate(III) – bent shape, chlorate(V) –
pyramidal (trigonal pyramid) shape and chlorate(VII) – tetrahedral shape.
Although ions, all the internal bonds are covalent and the some of the
electrons are delocalised in orbitals between the Cl and O atoms
constituting the bond.
The chloride ion acting as a
electron pair donor ligand in transition metal complexes forming dative covalent
bonds with the central transition metal ion
Left: The octahedral complex dichlorotetraaquachromium(III) ion.
Right: The geometric (or E/Z) isomers of the
square planar complex platinum(II) complex 'platin'.
In both cases all the bond angles
are 90 or 180o
GCSE Level GROUP 7 HALOGENS NOTES are on a separate webpage
INORGANIC Part 9
Group 7/17 Halogens sub–index:
9.1 Introduction, trends
& Group 7/17 data * 9.2 Halogen displacement
reaction and reactivity trend * 9.3 Reactions of
halogens with other elements - halides * 9.4
Reaction between halide salts and conc.
sulfuric acid *
9.5 Tests for halogens and halide ions *
9.6 Extraction of halogens from natural sources
* 9.7 Uses of halogens & compounds * 9.8
Oxidation & Reduction – more on redox reactions
of halogens & halide ions * 9.9 Volumetric
analysis – titrations involving halogens or halide ions * 9.10
Ozone, CFC's and halogen organic chemistry
links * 9.11 Chemical bonding in halogen
compounds * 9.12
Miscellaneous aspects of
Level Inorganic Chemistry Periodic Table Index:
Periodic Table history
Electron configurations, spectroscopy,
ionisation energies *
Period 1 survey H to He *
Period 2 survey Li to Ne * Part
5 Period 3 survey Na to Ar *
Period 4 survey K to Kr and important
trends down a group *
s–block Groups 1/2 Alkali Metals/Alkaline Earth Metals *
p–block Groups 3/13 to 0/18 *
Group 7/17 The Halogens *
3d block elements & Transition Metal Series
Group & Series data & periodicity plots
Group numbering and the modern periodic
The original group numbers of
the periodic table ran from group 1 alkali metals to group 0
noble gases (= group 8). To account for the d block elements and
their 'vertical' similarities, in the modern periodic table,
group 3 to group 0/8 are numbered 13 to 18. So, the halogen
elements are referred to as group 17 at a higher academic level,
though group 7 is still used, usually at a lower academic level.
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