Inorganic Chemistry
By
Dr. Khalil K. Abid
Lecture 9
Ionic Bonding
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Ionic Compounds
– ionic bond: electrostatic attraction holding together positively charged metal cations and negatively
charged nonmetal anions. Thus, an ionic compound is actually a three dimensional network of ions,
with each cation surrounded by anions, and vice versa.
As an ionic compound forms, like-charged ions repel each other and opposite-charged ions attract
each other. After all the pushing and pulling, the ions settle into alternating positions because this is
the most stable arrangement . Strong electrostatic forces of attraction between the positively charged
and negatively charged ions hold the particles together. Because these forces bind the ions together,
this is known as ionic bonding.
A lot of energy is required to move ions out of their positions because the electrostatic forces are
so strong. This means ionic compounds are hard to melt. At room temperature, they are in the form of
hard, brittle crystals. The most commonly known example of an ionic compound is sodium chloride
(table salt). Its melting point is 801°C. If you use a salt grinder at home, you will be aware of how hard
and brittle salt crystals are. Ionic compounds are formed from the bonding of ions.
Let’s consider sodium chloride, which is produced when sodium and chlorine meet and react. In
this compound, the metal sodium is present in the form of positively charged ions (Na+) and the nonmetal chlorine is present as negatively charged ions (Cl–).
Some transition metals can form more than one ion. In these cases, a Roman numeral is used to
show the charge on the ion. For example, copper forms two ions: one with a 1+ charge and one with
a 2+ charge. These ions are called copper (I) and copper (II) ions, respectively.
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• Prototypical bonding example: NaCl (Sodium Chloride)
• When NaCl is heated to the point of melting, one can demonstrate that the resulting solution
conducts electricity.
• This observation demonstrates that the solution (liquid NaCl) contains charged species. Those
species are Na+ and Cl –.
• Why the formation of Na+ and Cl – ?
• In short, Na+ and Cl – are more energetically stable than atomic Na and Cl.
• With the transfer of an electron from Na to Cl, two ions of opposite charge are produced. The
Coulombic attraction between these ions is largely responsible for the stabilization.
Properties of Ionic Materials
Although there is no sharp boundary between ionic bonding and covalent bonding. it is convenient
to consider each of these as a separate entity before attempting to discuss molecules and lattices, in
which both are important. The simplicity of the electrostatic model has caused chemists to think of
many solids as systems of ions. We shall see that this view needs some modification, and there are,
of course, many solids, ranging from diamond to metals, which require alternative theories of
bonding. Several properties distinguish ionic compounds from covalent compounds. These may be
related rather simply to the crystal structure of ionic compounds, namely, a lattice composed of
positive and negative ions in such a way that the attractive forces between oppositely charged ions
are maximized and the repulsive forces between ions of the same charge are minimized.
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Ionically bonded materials usually have moderate hardness and fairly high melting points.
They are generally soluble and are poor conductors of electricity because their constituent ions are
fairly stable and neither lose nor gain electrons easily. They are usually highly symmetric and non
directional (isotropic), because their cations tend to evenly surround themselves with as many
anions as space permits.
1. Ionic compounds tend to have very low electrical conductivities as solids but conduct
electricity quite well when molten. This conductivity is attributed to the presence of ions, atoms
charged either positively or negatively, which are free to move under the influence of an electric
field. In the solid, the ions are bound tightly in the lattice and are not free to migrate and carry
electrical current. It should be no ted that we have no absolute proof of the existence of ions in
solid sodium chloride, for example. The fact that ions are found when sodium chloride is melted or
dissolved in water does not prove that they existed in the solid crystal However, their existence in
the solid is usually assumed, since the properties of these materials may readily be interpreted in
terms of electrostatic attractions.
2. Ionic compounds tend to have high melting points. Ionic bonds usually are quite strong and
they are omnidirectional. The second point is quite important, since ignoring it could lead one to
conclude that ionic bonding was much stronger than covalent bonding which is not the case. We
shall see that substances containing strong. multidirectional covalent bonds, such as diamond,
also have very high melting points.
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The high melting point of sodium chloride, for example, results from the strong electrostatic
attractions between the sodium cations and the chloride anions, and from the lattice structure, in
which each sodium ion attracts six chloride ions. each of which in turn attracts six sodium ions, etc..
throughout the crystal.
3. Ionic compounds usually are very hard but brittle substances. The hardness of ionic substances
follows naturally from the argument presented above, except in this case we are relating the
multivalent attractions between the ions with mechanical separation rather than separation through
thermal energy. The tendency toward brittleness result from the nature of ionic bonding. If one can
apply sufficient force to displace the ions slightly (e.g~ the length of one half of the unit cell in NaO),
the formerly attractive forces become repulsive as anion-anion and cation - cation contacts occur;
hence the crystal flies apart. This accounts for the well-known cleavage properties of many minerals.
4. Ionic compounds are often soluble in polar solvents with high permittivities (dielectric
constants). The energy of interaction of two charged particles is given by:
E = q+q-
4π r εo
where q+ and q- are the charges, r is the distance of separation, and εo is the permittivity of the
medium. The permittivity of a vacuum, εo , is 8.85 x 10-11 C2 m-1 J - 1
For common polar solvents, however, the permittivity values are considerably higher. For example,
the permittivity is 7.25 X 10—10 C2 m- 1 J -1 for water, 2.9 X 10—10 C2 m- 1 J -1 for acetonitrile,
and 22 X 10—10 C2 m- 1 J -1 for ammonia, giving relative permittivities of 82 εo (H2O). 33 to (CH3CN), and
25 εo (NH3).
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Since the permittivity of ammonia is 25 times that of a vacuum. the attraction between ions
dissolved in ammonia, for example, is only 4% as great as in the absence of solvent. For solvents
with higher permittivities the effect is even more pronounced. Another way of looking at this
phenomenon is to consider the interaction between the dipole moments of the polar solvent and the
ions.
Simple ionic compounds form only between very active metallic elements and very active non
metals. Two important requisites are that the ionization energy to form the cation, and the electron
affinity to form the anion, must be energetically favorable. This does not mean that these two
reactions must be exothermic means, rather, that they must not cost too much energy.
Thus the requirements for ionic bonding are:
(1) The atoms of one element must be able to lose one or two (rarely three) electrons without
undue energy input.
(2) The atoms of the other clement must be able to accept one or two electrons (almost never
three) without undue energy input.
This restricts ionic bonding to compounds between the most active metals: Groups IA(1), IIA(2),
part of IVA(3) and some lower oxidation states of the transition metals (forming cations), and the most
active nonmetals: Groups VIIA(17). VIA(I6), and nitrogen (forming anions). All ionization energies are
endothermic, but for the metals named above they are not prohibitively so. For these elements,
electron affinities are exothermic only for the halogens. but they are not excessively endothermic for
the chalcogens and nitrogen.
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Ionic Solids
The crystal structures of ionic solids composed solely of monatomic ions is governed by two
relatively simple factors.
1. The most efficient packing of the ions is desired.
2. Because like charges repel, it is necessary to prevent two anions or two cations from coming in
contact with each other.
Just as covalent bonding and non-spherical atoms created exceptions to the closest-packing rule
for metals, these same issues can be important in the packing of ionic solids. To pack monatomic ions
together to form an ionic solid, first select the largest ion, which is frequently the anion. Pack these
large ions together in a closest - packed fashion (either hcp or ccp). Next insert the smaller ion
(usually the cation) in the holes in the closest-packed structure for the larger ion. It is important that
ions of like charge do not come in contact; therefore the smaller ion must be larger than the size of the
hole in which it is being place. If the smaller ion is smaller than the hole, it will fit into the hole and the
larger ions will remain in physical contact. If the smaller ion is larger than the hole, the smaller ion will
push the larger ions apart and thus prevent the high-energy situation of having two ions of like charge
in physical contact.
Holes in Closest-Packed Structures
Both ccp and hcp have many trigonal holes. However, the rhole/r ratio for a tetrahedral hole (0.2247) is
very small and relatively few ionic solids have cation-anion radius ratios smaller than this.
Consequently it is rarely necessary to utilize trigonal holes. The vast majority of ionic solids have ions
in tetrahedral, octahedral, or cubic environments.
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Sizes of Holes: In viewing the exercises involving trigonal, tetrahedral, octahedral, and cubic holes, it
should have been apparent that the holes differ in size. Specifically, the ratio rhole/rwas smallest for
the trigonal hole and largest for the cubic hole. The virtual reality display below shows the relative
sizes of the largest spheres that can fit into a hole of each geometry. The light blue sphere on the far
right shows the size of the lattice atoms whose packing creates the hole.
Structure of Ionic Solids: Consider an ionic solid composed of monatomic cations and anions and
assume the anion is larger than the cation. In the interests of efficient packing, the anions will pack
together in a ccp or hcp structure. The cations will then be inserted into holes in the closest-packed
structure. It is important that the cation be larger than the hole into which it is inserted so that the
anions in the structure are pushed apart and thus are not in physical contact. In order to achieve the
most efficient packing, the cation will go in the smallest hole possible so long as the cation
is larger than the hole. For a tetrahedral hole, rhole/r = 0.225. For an octahedral hole, this ratio is 0.414.
For a cubic hole, the ratio is 0.732.
Suppose rcation/ranion = 0.38. If the cation is placed in a cubic or octahedral hole, the anions will not
be pushed apart, because the cation is smaller than the hole. In this case the cation will go into a
tetrahedral hole. Suppose rcation/ranion = 0.68. If the cation is placed in a cubic hole, the anions will not
be pushed apart, because the cation is smaller than the hole. The cation must go into a tetrahedral or
octahedral hole. Placing the cation in a tetrahedral hole will result in less efficient packing (fewer
atoms per unit volume) than placing the cation in an octahedral hole, because the tetrahedral hole is
smaller and thus the anions will be pushed further apart. As a result the cation will end up in
octahedral holes.
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If rcation/ranion > 0.732, the cation will prefer to go in cubic holes. Observant readers, however, will
recall that there are no cubic holes in either the hexagonal closest-packed structure or the cubic
closest-packed structure. Two possibilities exist for coping with this situation. First, the anions
might form a cubic structure rather than a ccp or hcp structure. Second, although there are no cubic
holes in the ccp structure, you may recall that if all tetrahedral holes are filled with cations, the
cations themselves create a cubic hole for the anion. The ccp structure might, therefore, be used
if rcation/ranion is near unity.
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Coordination of Ions
For ionic bonding,
ion geometry ~
spherical. Spherical
ions will geometrically
pack (coordinate)
oppositely charged
ions around them as
tightly as possible
while maintaining
charge neutrality. For a
particular ion, the
surrounding
coordination ions
define the apices
(corners) of a
polyhedron. The
number of surrounding
ions is the Coordination
Number
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The Structure of Ionic Solid Crystals
Sodium chloride is taken as a typical ionic compound. Compounds like this consist of a giant
(endlessly repeating) lattice of ions. So sodium chloride (and any other ionic compound) is
described as having a giant ionic structure. You should be clear that giant in this context doesn't just
mean very large. It means that you can't state exactly how many ions there are. There could be
billions of sodium ions and chloride ions packed together, or trillions, or whatever - it simply
depends how big the crystal is. That is different from, say, a water molecule which always contains
exactly 2 hydrogen atoms and one oxygen atom - never more and never less. A small representative
bit of a sodium chloride lattice looks like this:
If you look at the diagram carefully, you will see that the sodium ions and chloride ions alternate
with each other in each of the three dimensions. Only those ions joined by lines are actually
touching each other. The sodium ion in the centre is being touched by 6 chloride ions. By chance we
might just as well have centered the diagram around a chloride ion - that, of course, would be
touched by 6 sodium ions. Sodium chloride is described as being 6:6 coordinated. You must
remember that this diagram represents only a tiny part of the whole sodium chloride crystal. The
pattern repeats in this way over countless ions.
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Why is sodium chloride 6:6-co-ordinated?
The more attraction there is between the positive and negative ions, the more energy is released.
The more energy that is released, the more energetically stable the structure becomes. That means
that to gain maximum stability, you need the maximum number of attractions.
The different structure of cesium chloride
We'll look first at the arrangement of the ions and then talk about why the structures of sodium
chloride and cesium chloride are different afterwards. How the ions are arranged in cesium chloride.
Imagine a layer of chloride ions as shown below. The individual chloride ions aren't touching
each other. That's really important - if they were touching, there would be repulsion. If you now think
about a cesium ion sandwiched between the two layers of chloride ions, it is touching four chloride
ions in the bottom layer, and another four in the top one. Each cesium ion is touched by eight
chloride ions. We say that it is 8-co-ordinated. If we added another layer of cesium ions, you could
similarly work out that each chloride ion was touching eight cesium ions. The chloride ions are also
8-co-ordinated.
Overall, then, cesium chloride is 8:8-co-ordinated. The final diagram in this sequence takes a
slightly tilted view of the structure so that you can see how the layers build up. These diagrams are
quite difficult to draw without it looking as if ions of the same charge are touching each other. They
aren't!. Diagrams of ionic crystals are usually simplified to show the most basic unit of the repeating
pattern.
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For cesium chloride, draw a simple diagram showing the arrangement of the chloride ions around
each cesium ion, you would have an exactly equivalent diagram for the arrangement of cesium ions
around each chloride ion. Now imagine what would happen if you replaced the cesium ion with the
smaller sodium ion. Sodium ions are, of course, smaller than cesium ions because they have fewer
layers of electrons around them. You still have to keep the chloride ions in contact with the sodium.
The effect of this would be that the whole arrangement would shrink, bringing the chloride ions into
contact with each other - and that introduces repulsion.
Any gain in attractions because you have eight chlorides around the sodium rather than six is more
than countered by the new repulsions between the chloride ions themselves. When sodium chloride is
6:6-co-ordinated, there are no such repulsions - and so that is the best way for it to organize itself.
Which structure a simple 1:1 compound like NaCl or CsCl crystallises in depends on the radius ratio of
the positive and the negative ions. If the radius of the positive ion is bigger than 73% of that of the
negative ion, then 8:8-co-ordination is possible. Less than that (down to 41%) then you get 6:6-coordination. In CsCl, the cesium ion is about 93% of the size of the chloride ion - so is easily within the
range where 8:8-co-ordination is possible. But with NaCl, the sodium ion is only about 52% of the size
of the chloride ion. That puts it in the range where you get 6:6-co-ordination.
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NaCl Crystal Structure
CsCl structure
Circles represent the Na+ and Cl- ions.
In CsCl, metal ions are shifted into the
Each ion is surrounded by six other ions
center of each cubic element of the Cl–
of opposite charge, therefore NaCl is
ion lattice. Each cesium ion has eight
described as having (6,6) coordination.
nearest-neighbor chloride ions, while
each chloride ion is also surrounded by
eight cesium ions in (8,8) coordination.
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we can now complete our table for predicting cation coordinations from the radius ratio rules:
1 – Radius Na+ = 1.02Å, radius Cl- = 1.81Å; radius ratio = 0.563. Therefore Na octahedral.
Coordination # Na = 6; coordination # Cl = 6.
2 – Radius Zn2+ = 0.6Å, radius S2- = 1.84Å; radius ratio = 0.33 Æ Zn tetrahedral.
Have 2 tetrahedral sites/ anion, therefore from formula of ZnS only 50% of the tetrahedral sites can
be filled. Coordination # Zn = 4; coordination # S = 4. Simple cubic arrangement of anions - 50%
3 – ionic radius Ca2+ = 1.12Å; radius F- = 1.31Å; radius ratio = 0.85 Æ Ca2+ cubic coordination.One
cubic site per F anion; from stoichiometry only 50% cubic sites filled by Ca cations.
Arrangement of the filled cubic sites is such that the Ca-Ca distances are as large as possible .C. N.:
Ca2+ surrounded by 8 F- 's; F- surrounded by 4 Ca2+'s.
4 – CsCl: radius Cs+ = 1.74Å, radius Cl- = 1.81Å: radius ratio = 0.96 Æ predict cubic coordination.
All cubic sites are filled by Cs cations. C. N.: Cs = 8; Cl = 8.
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The Ionic Lattice Crystal Packing
In an ionic solid, the ions are packed together into a repeating array called a crystal lattice. The
concept of crystal packing assumes that the ions are hard spheres. The easiest way to picture such
an array is to arrange one layer of spheres and then place successive layers over it.
The simplest arrangement is one in which the spheres in the base are packed side by side. The
successive layers of spheres are then added directly over the spheres of the layer below. This type
of array is known as:
(1) simple cubic packing. In this arrangement each atom is touched by four other atoms in its own
plane plus one atom above and one below. Thus, each sphere is touched by six neighboring spheres
giving it a coordination number of six.
(2) An alternative packing arrangement can be obtained by placing the second layer of spheres
over the holes (or interstices) of the base layer. The third layer of spheres are placed over the holes
of layer two. Successive layers are added in the same fashion. This type of array is known as bodycentered cubic (bcc). This is a more compact arrangement than the simply cubic packing array. In
this arrangement, each sphere is touched by four atoms above and four atoms below its plane giving
a coordination number of eight.
(3)Another arrangement has a base layer of spheres organized in a hexagonal arrangement in
which each each sphere is surrounded by six neighbors in the plane. In this arrangement the holes
between spheres are smaller than in the cubic arrangement.
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(4)When placing the second hexagonal layer over the first, it is physically impossible for spheres to
be placed over all the holes in the first layer -- only half of the holes can be covered. If a third layer is
placed over the holes of the second layer so it is superimposed over over the base layer,
the hexagonal close-packed array is obtained. Each sphere has a coordination number of twelve.
(5) If the third layer of spheres instead of being placed over the hole of the second layer,is placed
over the holes not covered from the first layer, the arrangement is called cubic close-packed.
(1)
(2)
(4)
(3)
(5)
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