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4.1. INTRODUCTION
4.1.1. When it comes to bonding, perhaps the most important
section in chapter 2 was section 2.3.3.: it is essential to imagine
the electrostatic forces involved in bonding.Energy is also a
useful concept, but on its own, it does not always lead to such a
clear picture of the way things are.
It is electrostatic attraction which holds atoms and ions
together, not energy. However, complex energy calculations provide
essential information (for example the shapes of atomic and
molecular orbitals). We do not need to examine the calculations in
detail, but we do need to include in our models the information
they provide (section 4.8.).
Moreover, energy gives us a quantifiable measure of the forces.
A measure which is easily determined from experiments (chapter 9).
See also sections 6.1.1. and 14.4.2.v.
4.2. BONDING AND THE PERIODIC TABLE
4.2.1. Section 2.4.2. was also important preparation for this
chapter.
There are few facts at this level which must simply be learned.
However, you must be able to instantly picture the place in the
periodic table of any element mentioned in your syllabus. If you
have these positions stamped on your mind as the starting point for
all your inorganic theory, there will be far fewer other facts that
you simply need to remember.
Here are the elements whose positions you must know for the
majority of syllabi at this level (FIG. 4.1.):
A valuable part of inorganic chemistry is learning how to relate
a vast mass of theory to two simple observations, and how to use
the relationship to make predictions:
i) First observation: Physical and chemical properties of
elements and their compounds follow largely predictable patterns
down groups in the periodic table.
ii) Second observation: Physical and chemical properties of
elements and their compounds follow largely predictable patterns
across periods in the periodic table.
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If you come to grips with this approach to inorganic theory, you
will have mastered a powerful, and far more widely applicable,
technique.
4.2.2. Patterns down a group: Changes down a group follow from
the increasing number of electron shells down a group. This causes
the outer electrons to become further (*and more screened from) the
nucleus.
These changes are abbreviated as being due to increasing
size.
(* Quantum mechanics reveals that, surprisingly, screening does
not increase down a group as we might predict from the simple
"shell" model of atomic structure. However, this fact is not widely
accepted at this level, and the shell model prevails.)
4.2.3. Patterns across a period: Changes across a period follow
from the increasing number of protons in the nucleus which have to
hold only the same number of shells in place.
These changes are summarised as being due to increasing
effective nuclear charge.
There are also more obvious changes associated with the
increasing number of outer electrons.
4.2.4. Labouring the point: These patterns are fundamental.
Inorganic facts should be learned only as examples of these
patterns, or should at least be understood and related to each
other in the light of the patterns. There are unlimited pieces of
information you could learn.
Note also, that decreasing size across a period should never be
given as a fundamental cause of any other pattern. Decreasing size
itself follows from increasing effective nuclear charge.
This is so important that it is probably better not to use the
concept of effective nuclear charge when explaining how patterns
down a group tie together. (The nuclear charge down a group
increases, but the effect of the nucleus on the outer electrons
decreases.) It is best to refer back to the increased distance (and
screening according to most examiners) of the outer electrons from
the nucleus.
4.3. ELECTRONEGATIVITY AND ELECTROPOSITIVITY
4.3.1. The type of bonding between two elements can be
conveniently predicted from their relative electronegativities.
In this context, electronegativity can be regarded as the
ability of an atom to gain electrons and form negative ions.
Conversely, electropositivity can be regarded as the ability of an
atom to lose electrons and form positive ions.
Often the single characteristic of electronegativity is used: an
electropositive element has a very low electronegativity.
Electronegativity follows group and periodic trends and
changes.
4.3.2. Electronegativity down a group: Down a group, the outer
electrons become further ("and more screened") from the nucleus.
They therefore become less strongly
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attracted by the nucleus and more easily lost: the elements
become more electropositive. It also becomes more difficult for the
nucleus to attract additional electrons into the outer shell: the
elements become less electronegative.
4.3.3. Electronegativity across a period: Across a period, the
increasing effective nuclear charge means that the outer electrons
are more tightly held, and that additionalelectrons are more
strongly attracted into the outer shell: the elements become less
electropositive and more electronegative.
4.4. PREDICTING BOND TYPE
4.4.1. By combining our understanding of the periodic table and
of electronegativity, we can predict the type of bonding between
pairs of elements.
4.4.2. Bonding between elements widely separated in the periodic
table: Elements towards the left of the table and towards the
bottom will tend to be the (leastelectronegative and) the most
electropositive. They will be the most likely to lose electrons and
form positive ions.
Conversely, those towards the top right hand corner of the
periodic table will be the most electronegative (and least
electropositive), and most likely to form negative ions.
These opposite types of element, which differ widely in
electronegativity, are therefore most likely to form electrovalent
(ionic) bonding with each other.
Definition: ionic bonding is the electrostatic attraction
between oppositely charged ions which are arranged in a crystal
lattice and which are formed by the transfer of electrons from one
atom (giving positive ions) to another (giving negative ions).
Such a transfer of electrons to form ions can be represented for
sodium chloride as follows:
.Note that we cannot define a single ionic bond in a lattice,
because an ion is attracted to all the surrounding (oppositely
charged) ions in the lattice, as seen in the sodium chloride
lattice (FIG. 4.2.):
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4.4.3. Bonding between elements close in the table:
i) The more electropositive elements: When two elements occur in
the more electropositive region of the periodic table, neither is
able to completely gain electrons from the other.
Moreover, neither is able to maintain sufficient control over
the bonding electrons to form a covalent bond (section 4.4.4.).
In fact, those outer electrons which are available for bonding,
form a cloud of electron density in which the resultant positive
ions are arranged as a lattice. The bonding is
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metallic bonding, and it occurs between atoms of the same
element as well as in alloys (see below).
Definition: metallic bonding is the electrostatic attraction of
positive ions arranged in a lattice for the electrons in a single
metallic bonding orbital which permeates the lattice (FIG.
4.3.)
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This type of bonding, unlike electrovalent or covalent bonding,
does not require atoms of different elements to be present in
fixed, whole number, (stoichiometric) ratios. True compounds are
not formed, but alloys (section 18.7.). Alloys are intimate
associations, somewhere between mixtures and compounds, in which a
degree of separation can sometimes be achieved by physical
means.
4.4.4. Bonding between elements close in the table:
ii) The more electronegative elements: When two atoms occur in
the more electronegative part of the periodic table, neither will
be able to completely lose control of electrons to the other.
Moreover, neither is able to lose sufficient control over the
available bonding electrons to form a metallic bonding orbital.
In fact, the bonding electrons are held in a well defined space
between the two bondingatoms.
The precise definition of a covalent bond may surprise you,
because at this stage it is important to emphasise that all bonds
are electrostatic. Too often, covalent bonds are made to sound like
diplomatic pacts based on an agreement between two atoms to share a
pair of electrons.
Definition: A covalent bond between two atoms is the
electrostatic attraction of the two nuclei for a shared pair of
electrons in a bonding orbital formed by the overlap of two singly
occupied atomic orbitals (or, in the case of a dative covalent
bond, sometimes known as a coordinate bond, formed by the overlap
with a vacant orbital of an atomic orbital containing a pair of
electrons).
Note that for the first time we have been able to define a
single bond. Note also that the key difference between one type of
bond and another is where the electrostatic forces occur. All bonds
are electrostatic forces of attraction between positive charge
andnegative charge, and therefore (ultimately) between protons and
electrons.
A simple example of a covalent bond is shown for hydrogen in
FIG. 4.4.
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The formation of a dative covalent bond is illustrated by the
following example:
The arrow in the ammonium ion represents the direction in which
the electrons have been donated to the bond. In practice, once
formed, the new bond is indistinguishable from the other three
covalent bonds. Dative covalent bonds are in general
indistinguishable from ordinary covalent bonds. It is their history
which differs.
4.4.5. When two hydrogen atoms collide they do not always form a
covalent bonding orbital like that shown in FIG. 4.4. For example,
if the two atomic orbitals contain electrons with parallel spins,
they will tend to repel each other. If they have opposite spins,
their magnetic moments cancel out in a bonding orbital.
When repulsion occurs, it is sometimes said that an anti-bonding
orbital has formed.This may seem like an elaborate construct for a
non-event, and you will have to read elsewhere if you want to be
convinced otherwise.
Certainly, at this level, the concept is likely to lead to
meaningless statements like: "He2 molecules therefore do not exist
because their formation is energetically unfavourable". This leaves
unanswered real questions like, "What makes their formation
energetically unfavourable?"
An explanation in terms of electrostatic forces would be closer
to experience. If you could get down to atomic level and pull an
electron away from an atom you would experience forces due to
electrostatic attraction, not energy - at least, not in models
which are easy to grasp!
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The concept of energy is more useful in chemistry when
considering something you would experience as energy, such as heat.
It also provides a unifying concept, for example between
electrostatic forces and heat or light. Moreover, it is useful as a
tool for developing better models, as mentioned in section 4.1.1.,
and as we shall see in the section on hybrid orbitals (section
4.8.).
Energy is also an extremely important concept for checking
predictions, and for making them when further explanation (or more
detailed description) is unnecessary (section 14.4.2.v.).
4.5. OTHER TYPES OF BONDING
4.5.1. Hydrogen bonds: Hydrogen bonds are well defined by a
description of the hydrogen bonds in hydrogen fluoride (FIG.
4.5).
4.5.2. Van der Waals bonds are very weak forces of electrostatic
attraction amongst instantaneous dipoles and induced dipoles.
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At a given moment in time, the electron cloud surrounding even
an inert gas atom may not be perfectly symmetrical. Thus one side
will carry a relative positive charge and the other a relative
negative charge. The separation of charge (dipole) can be almost
negligible in small atoms such as helium. In addition to these
instantaneous dipoles, there are dipoles induced by the original
dipoles.
Since the dipoles are very small, the forces which exist between
them are also very small. Even at low temperatures, the random
motion of group VIII atoms is vigorous enough to keep them gaseous.
Moreover, on collision, repulsion between like charges is just as
likely to occur as attraction between opposite charges.
However, as extremely low temperatures are reached, the random
motion is not enoughto break van der Waals attractions more often
than they occur. Eventually, at low enough temperatures, the gas
may liquefy, and the liquid may then even solidify. In energy
terms, the kinetic energy is less than the bonding energy.
The larger an atom or molecule, the easier it is for dipoles to
exist. In large atoms this is because the outer electrons are a
long way from the control of the nucleus. Relativelysimple large
molecules (e.g. iodine) may solidify even at room temperature. In
contrast,helium must be reduced to 4K, and a pressure of 103
atmospheres must be applied, in order for solidification to
occur.
Van der Waals forces are particularly important in determining
the shapes which macromolecules such as proteins adopt (section
25.2.5.). Here, the large number of bonds compensates for their
extreme weakness. This may also be the case between very large
molecules. Note that in DNA, large numbers of hydrogen bonds are
the important bonds holding the two strands together (section
25.4.6.).
4.6. BONDING CHARACTERISTICS
4.6.1. The main type of bonding has effects on properties, as
summarised in TABLE. 4.1.
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The fact that ionic compounds have high melting and boiling
points whereas covalent compounds have low melting and boiling
points does not indicate that ionic bonds are stronger than
covalent bonds.
The high melting and boiling points of ionic compounds do
indicate that ionic bonds are strong i.e. the temperature must be
high before ionic motion is vigorous enough to overcome the
electrostatic forces of attraction.
However, covalent bonds are also generally very strong, but it
is not these which break when small molecular compounds melt or
boil. It is extremely weak van der Waals forces between one
molecule and another which break.
4.7. INTERMEDIATE TYPES OF BONDING
4.7.1. The three main types of bonding, electrovalent, metallic,
and covalent are extreme cases. As often as not, bonding is
intermediate between two types.
4.7.2. Electrovalent vs. covalent: There is an initial problem
here of deciding whether to treat the bonding as ionic with
covalent character or covalent with ionic character. This must be
judged from an overall assessment of properties.
For example, any compound which conducts electricity in the
molten state is more likelyto be judged as principally ionic than
it is to be judged as principally covalent, and vice versa.
i) Ionic bonding with covalent character: When two types of ion
come together toform a compound, there is likely to be a high
degree of covalent character when:
First, the positive ion is: a)
small...................................b) highly charged
Second, the negative ion is: a)
large........................................b) highly charged
(These empirical rules are known as Fajan's rules.)
It is predicatable that a small, highly charged positive ion
will have a high surface charge density, and will tend to attract
electrons away from the negative ion into a covalent bond.
It is also predictable that it will be easiest to remove
electrons from negative ions when those electrons are a long way
from the negative ion's nucleus, and when there is a large number
of "extra" electrons in the outer shell.
ii) Covalent bonding with ionic character: When a largely
covalent bond forms between elements of differing electronegativity
(most cases), the electron pair will not be shared equally.
The more electronegative element will have a greater share of
the bonding pair, as in the case of HCl molecules:
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If it seems like fantasy to consider this molecular gas as
partially ionic, remember that itdissolves in water to produce
hydrated H+ and Cl- ions.
Sometimes the polarity of a molecule will be enhanced by its
shape and the existence oflone pairs:
The assymetric ammonia molecule is polar. Apart from discrete
polar molecules, molecular solids may also exhibit ionic character.
In fact, the mere existence of covalentmolecules in a solid lattice
is a step towards ionic character: the van der Waals bonding
between molecules is becoming less distinguishable from the
covalent bonding within the molecules (e.g. section 14.3.1.).
4.7.3. Covalent vs. metallic bonding: This is most relelvant to
the case of pure elements and is well illustrated by changes down
group IV (section 17.2.1) and group VII (section 16.2.1.) and by
changes across the period from sodium to chlorine. We shall discuss
the changes across a period here.
Sodium (LHS) is electropositive and forms clear-cut metallic
bonding. Chlorine (RHS) is electronegative and forms diatomic
molecules with a precise covalent bond between thetwo atoms.
From sodium to aluminium, the increasing effective nuclear
charge (as well as the increasing number of electrons contributed
by each atom to the metallic bonding orbital) causes the metallic
bonding to become stronger. (Note that this is not the same as
saying the bonding becomes become more metallic.)
By silicon, the effective nuclear charge is high enough to hold
the bonding electrons in fairly concentrated regions of electron
density on the axes between the nuclei i.e. the bonds are more
covalent than metallic, but silicon still has a lattice structure
more
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comparable with a metal than with the discrete molecules
normally associated with covalent bonding.
Moreover, silicon has a metallic lustre, suggesting that the
bonding electrons are diffuseenough to reflect light. Some of
silicon's semi-conductive properties may also be considered as
intermediate between metallic and covalent properties.
The higher effective nuclear charges of phosphorus and sulphur
result in even more coavlent character. Discrete molecules are
formed, P4 and S8, but van der Waals forces between one molecule
and another are strong enough for both elements to be solids at
room temperature.
4.8. PREDICTING THE NUMBER OF BONDS
4.8.1. Having predicted the type of bonding, it is necessary to
predict the number of bonds (charge in the case of ions).
4.8.2. The octet fable: In its least controversial form the
octet "rule" observes that atoms often end up with eight outer
electrons when they form bonds. In its most controversial form it
suggests that atoms form bonds in order to achieve this
configuration. The configuration acquires the status of "stable
octet" which is attributed with magical properties resulting from
its similarity to the outer electron configuration ofthe "noble"
gas atoms.
Perhaps it is the wishy-washy use of stability as an explanation
which has allowed the continuation of the myth. In more cases than
this, use of the word "stable" is an attempt to make "low energy"
sound like a complete explanation.
The facts are even more disturbing. For one thing, in energy
terms the octet is usually less stable than the electron
configuration of the parent atom. This is indicated by the fact
that even group I metal atoms (which are amongst the most able to
lose electrons and form positive ions with an outer octet) require
energy to remove the outer electron.
It is true that Cl-, with its outer octet, is more "stable" than
the parent atom (electron affinities are often negative) but this
is not because forming negative ions achieves the octet.
This is clearly illustrated by oxygen. An O- ion is more
"stable" than an oxygen atom. The first electron affinity is
negative i.e. the ion has less energy than the parent atom and
energy is given out as heat when it forms. However, when O- goes on
to achieve the octet by forming O2-, energy is required.
Elements which form negative ions are elements with high
effective nuclear charges, and this is "why" first electron
affinities can be negative. However, when a second electron is
added to O-, it is repelled by the negative charge of the ion and
must be added to a small shell which already contains seven,
negatively charged, electrons.
Note that we have found it necessary to explain energy changes
in terms of electrostatic forces.
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4.8.3. Rubbing it in: In addition to all this, there are plenty
of cases where bonding does not result in an outer octet:
i) More than the octet: a) Electrovalent: Fe3+ has 13
outer..........................................electrons in
Fe2(SO4)3......................................b) Covalent: Sulphur
has 12 outer..........................................electrons in
SF6
ii) Fewer than the octet: a) Electrovalent: Pb2+ has 2
outer............................................electrons in
PbSO4........................................b) Covalent: Al has 6
outer electrons............................................in
AlCl3
The situation is more involved than the octet rule implies.
There are far more events occurring than merely the formation, or
otherwise, of octets. It is convenient to tackle the problem in
terms of a more detailed look at energy changes, and to refer to
the electrostatic forces involved in order to understand them.
4.8.4. Predicting the charge on a positive ion: Obviously this
step comes after deciding that a compound is electrovalent and
after deciding which element constitutes the positive ion.
i) The charge on a positive ion is equal to the number of
electrons lost by the parent atom.
An atom is unlikely to lose more electrons than it has in the
outer shell because the next shell is closer to ("and less screened
from") the attraction of the nucleus.
An atom is unlikely to lose fewer electrons than it has present
in the outer shell. This is because complete loss of the outer
shell electrons not only makes the ion more highly charged, but
also makes it much smaller. The resultant ion will therfore *form a
much stronger lattice (in aqueous solution, it will be much more
strongly hydrated).
With respect to calcium, try to picture the events and imagine
the changing electrostaticforces when calcium atoms react with
oxygen atoms. Then try to imagine the even more complex events when
calcium metal reacts with oxygen gas.
(*In energy terms, the much higher lattice energy compensates
for the extra ionisation energies.)
ii) Exceptions: Unlike the octet rule, this approach provides a
basis for understanding exceptions. For example, elements at the
bottom of group IV can form 2+ ions as well as the expected 4+ ions
(section 17.3.2.iii).
The behaviour of d-block elements can also be understood. They
can lose penultimate shell electrons as well as outer electrons
because there is little difference in the strengthwith which their
penultimate d-electrons and outer s-electrons are held by the
nucleus (section 2.4.3.).
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4.8.5. Predicting charge on a negative ion: When an
electronegative atom gains electrons to form negative ions it fills
any singly occupied orbitals in its outer shell.
It is unlikely to gain more than this number of electrons, since
this would involve attracting electrons into a shell further ("and
more screened from") the attraction of thenucleus.
It is unlikely to gain fewer electrons, because filling the
singly occupied orbitals increases the charge but results in a
relatively small increase in the overall size of the ion. The
resultant ion will therefore *form a stronger lattice than an ion
which still has asingly occupied orbital.
With respect to oxygen, try to picture the events and imagine
the changing electrostaticforces when calcium atoms react with
oxygen atoms. Then try to imagine the even more complex events when
calcium metal react with oxygen gas.
(* In energy terms, the extra lattice energy compensates for the
extra electron affinity.)
4.8.6. Predicting the number of covalent bonds:
i) Covalent bonds: Obviously, predicting the number of covalent
bonds comes after deciding that the bonding is covalent. Covalent
bonds are usually formed by the overlapof singly occupied atomic
orbitals, and the number of singly occupied orbitals in atom gives
a rough indication of the number of bonds it will form.
ii) Electron promotion: There are, as you will recall,
complications. For example, carbon in its "ground state" has two
singly occupied orbitals:
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However, carbon almost exclusively forms four bonds. This is
described in terms of a model in which one of the s electrons is
promoted into the vacant 2p orbital, giving it four singly occupied
orbitals:
The advantage of forming 4 bonds can be seen by considering
methane. In each C-H bond the carbon and hydrogen nuclei are
attracted to the high electron density of a bonding pair of
electrons. This attraction is greater than the total attraction of
a carbon nucleus for one of its outer electrons plus that of a
hydrogen nucleus for its outer electron in an H2 molecule. In other
words, the formation of two extra bonds outweighs the advantages of
leaving carbon's s-electrons paired and closer to the nucleus than
they would be in a p-orbital. Moreover, the electrons actually
enter bonding orbitals rather than p-orbitals, as we shall see in
section 4.8.
In energy terms, the two extra bond energies more than
compensate for any promotionenergy, and the H-H bond energy (c.f.
section 17.3.1.ii.).
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iii) d-orbitals: from the third period onwards, d-orbitals are
available for electron promotion. Thus sulphur can form 2, 4, or 6
bonds:
In this electron promotion model, the number of bonds formed is
dependent on the number of electrons promoted. This in turn depends
on:
a) The sizes of the bonding atoms
b) The effective nuclear charges of the bonding atoms.
These can be deduced from the positions of the elements in the
periodic table. Moreover, they affect two main aspects of covalent
bonding which themselves have a bearing on the number of bonds
formed:
a) First they affect the strengths of the bonds formed: then it
can be predicted that the stronger the bonds, the larger the number
likely to form.
b) Second, they affect the amount of space available for bond
formation: then it can be predicted that the larger the bonding
atoms relative to the central atom, the smaller thenumber of bonds
likely to form, owing to lack of space.
Thus, referring back to sulphur, this element forms SF6 but not
SCl6. This is predicatble for two reasons: First, fluorine's
nucleus is closer to ("and less screened from") the bonding
electrons than is chlorine's nucleus. Fluorine therefore forms
stronger bonds than chlorine with sulphur. Second, chlorine's
larger size means that there is less room for 6 of its atoms to fit
around the small central sulphur atom (FIG. 4.17.)
You are likely to develop a feel for this kind of reasoning only
after considering a good number of examples. Note also:
c) Multiple bonding reduces the number of bonded atoms required
to form the higher oxidation states.
iv) Dative covalent bonds: a final complication is that orbitals
containing pairs of electrons may be involved in dative covalent
bond formation by overlap with vacant atomic orbitals.
Moreover, singly occupied orbitals may become available for
dative bond formation by spin-pairing:
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It is possible to predict the liklihood of dative bond formation
from its strength and the amount of room available, and thus from
sizes and effective nuclear charges of the bonding atoms. Hence
(ultimately) the prediction can be made from the positions of
thebonding elements in the periodic table. Again you will develop a
feel for this after considering a large number of cases.
In the N2O4 molecule cited above, it is found that in practice
all the N-O bonds are equivalent. This emphasises the fact that we
are using a only partially correct model in order to make
predictions. The model is developed by considering one more piece
of bonding theory, hybridisation.
4.9. HYBRIDISATION
4.9.1. The shapes of atomic orbitals come from the solution of
complex wave equations. Such calculations also allow shapes to be
assigned to bonding orbitals, and they rely heavily on energy
considerations.
The wave equations are simplified by considering only the
predicted shapes of orbitals which atoms could theoretically use to
form bonds. Quantum mechanists can actually calculate the space
around a nucleus in which an electron will most probably occur
(FIG.2.3.). The shape is a direct function of the orbital's energy,
and this can change if hybridisation is included in the quantum
mechanical calculations.
Despite the useful outcome of these calculations, it is not wise
to be generally bullied bymathemeticians. A cynic might say that
they resort to mathematical models because they have as much
difficulty understanding reality as we have understanding them.
Mathematical models must, as they are here, be related to reality.
The calculated shapeof an orbital is very useful for developing the
models we use to describe atoms and molecules.
4.9.2. Why we need the concept of hybrids: Even when carbon is
described as having promoted its 2s electron into the vacant 2p
orbital, it is not in an appropriate state to form bonds, according
to our existing models. This is despite having four singly occupied
orbitals.
According to our existing model, the four orbitals are not a
convenient shape, nor are they conveniently arranged, to overlap
with other atomic orbitals (FIG. 4.6.)
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However, orbital shape is calculated as a function of energy. If
the energy of an orbital were different from that considered as
normal in a free atom, we would consequently predict a different
shape. The energies of these orbitals would be different from
normal if the distribution of energies amongst the orbitals were
different from normal.
Moreover, the distribution of energy would be different from
normal if the 2s orbital were equivalent in energy to one or more
of the p-orbitals. As a predictive tool, it is suggested that this
equalisation of energy does exist. It is known as hybridisation,
and associated calculations describe some much more conveniently
shaped (sp hybrid) orbitals (FIG. 4.7.).
There are three possibilities: sp1, sp2, and sp3
hybridisation.
4.9.3. sp3 hybridisation: In this case the s-orbital is
considered as equivalent in energy to all three p-orbitals, giving
four sp3 hybrid orbitals (FIG. 4.8.).
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The four orbitals point to the four corners of a
tetrahedron.
4.9.4. sp2 hybridisation: In this case the s-orbital is
considered as equivalent in energy to two of the p-orbitlas, giving
three sp2 hybrids, and leaving one p-orbital with a higher energy
(FIG. 4.9.).
The three sp2 hybrids point to the three corners of an
equilateral triangle.
4.9.5. sp1 hybridisation: In this case the s-orbital is
considered as equivalent in energy to one of the p-orbitals, giving
two sp1 hybrids, and leaving two p-orbitals with ahigher energy,
but with the same energy as each other (FIG. 4.10).
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The sp1 orbitals are arranged linearly.
4.9.6. sp3 hybrids in methane (FIG. 4.11.):
In an s-bond, the electron density is concentrated along the
axis between the centres ofthe two bonded atoms.
4.9.7. sp2 hybrids in ethane (FIG. 4.12):
4.9.8. sp2 hybrids in buta-1,3-diene (FIG. 4.13.):
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Electrons are said to be delocalised when they are not confined
to the positions indicated in a simple bonding diagram, but are
spread out by, for example, p-orbital andp-bonding system
overlap.
4.9.9. sp2 hybrids in benzene (FIG. 4.14.):
4.9.10. sp1 hybrids in ethyne (FIG. 4.15.):
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4.10. SHAPES OF MOLECULES
4.10.1. Hybridisation theory allows the shapes of molecules to
be predicted. Shape may also be predicted by remembering that
electrons in one bond will repel electrons inanother. The repulsive
effect of multiple bonds will be greater than the repulsive effect
of single bonds.
Also, lone pairs of electrons will repel electrons in bonds, and
repel other lone pairs, as well. Lone pairs exert a greater
repulsive force even than multiple bonds.
Thus the shape of a molecule will tend to achieve maximum
separation of bonds and lone pairs, taking into consideration the
relative strengths of the repulsive forces. Theseare summarised as
decreasing in the order: lone pair > multiple bond > single
bond. (This is a summary of, so-called, Sidgwick-Powell Theory).
Some examples are shown inFIG 4.16 below.
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(Lattice arrangements in ionic compounds, metals, and molecular
crystals are summarised in chapter 5.)
4.11. QUESTIONS
1) In chapter 2 we implied that it was dangerous to talk about
causes, and about answers to the question "why?". How do you equate
this with sections 4.2.2. and 4.2.3.?
2) What makes helium unable to form He2 molecules?
3) A small highly charged positive ion is more likely than a big
positive ion with a single charge to form a strong ionic lattice?
Comment.
4) How can the forces which exist in an ionic lattice be
responsible for the existence of ions within that lattice? Answer
the question with reference to the formation of calcium oxide from
i) calcium atoms and oxygen atoms ii) calcium metal and molecular
oxygen gas.
5) Why is the phrase "less screened from" placed in parentheses
in this chapter?
6) Predict the simplest formulae of all compounds formed between
the following pairs of elements, and predict the type of bonding
found in each case:
copper and zinc, lead and oxygen, potassium and bromine,
chlorine and bromine, oxygen and sulphur, sulphur and chlorine,
beryllium and chlorine, strontium and sulphur, hydrogen and
sulphur, phosphorus and bromine, phosphorus and fluorine.
In three cases, explain your reasoning.
7) Predict the shapes of SO2, CO2, AlCl3, NCl3, the [Cu(H2O)6]2+
ion, and SCl4.
8) Draw simple bonding structures for NO, NO2, N2O4, N2O5,
SO32-, SO42-, S2O32-, S4O62-, S2O82-, and S2O72-.