IONIC (ELECTROVALENT) BONDING
A simple view of ionic bonding
The importance of noble gas structures
At a simple level a lot of importance is attached to the electronic structures of noble gases
like neon or argon which have eight electrons in their outer energy levels (or two in the
case of helium). These noble gas structures are thought of as being in some way a
"desirable" thing for an atom to have.
You may well have been left with the strong impression that when other atoms react, they
try to organize things such that their outer levels are either completely full or completely
empty.
Ionic bonding in sodium chloride
Sodium (2,8,1) has 1 electron more than a stable noble gas structure (2,8). If it gave away
that electron it would become more stable.
Chlorine (2,8,7) has 1 electron short of a stable noble gas structure (2,8,8). If it could gain
an electron from somewhere it too would become more stable.
The answer is obvious. If a sodium atom gives an electron to a chlorine atom, both
become more stable.
The sodium has lost an electron, so it no longer has equal numbers of electrons and
protons. Because it has one more proton than electron, it has a charge of 1+. If electrons
are lost from an atom, positive ions are formed.
Positive ions are sometimes called cations.
The chlorine has gained an electron, so it now has one more electron than proton. It
therefore has a charge of 1-. If electrons are gained by an atom, negative ions are formed.
A negative ion is sometimes called an anion.
The nature of the bond
The sodium ions and chloride ions are held together by the strong electrostatic attractions
between the positive and negative charges.
The formula of sodium chloride
You need one sodium atom to provide the extra electron for one chlorine atom, so they
combine together 1:1. The formula is therefore NaCl.
THE IB VIEW OF IONIC BONDING
 Electrons are transferred from one atom to another resulting in the formation of
positive and negative ions.
 The electrostatic attractions between the positive and negative ions hold the
compound together.
So what's new? At heart - nothing. What needs modifying is the view that there is
something magic about noble gas structures. There are far more ions which don't have
noble gas structures than there are which do.
Some common ions which don't have noble gas structures
You may have come across some of the following ions. They are all perfectly stable, but
not one of them has a noble gas structure.
Fe3+
[Ar]3d5
Cu2+
[Ar]3d9
Zn2+
[Ar]3d10
Ag+
[Kr]4d10
Pb2+
[Xe]4f145d106s2
Noble gases (apart from helium) have an outer electronic structure ns2np6.
Apart from some elements at the beginning of a transition series (scandium forming Sc3+
with an argon structure, for example), all transition elements and any metals following a
transition series (like tin and lead in Group 4, for example) will have structures like those
above.
That means that the only elements to form positive ions with noble gas structures (apart
from odd ones like scandium) are those in groups 1 and 2 of the Periodic Table and
aluminum in group 3 (boron in group 3 doesn't form ions).
Negative ions are tidier! Those elements in Groups 5, 6 and 7 which form simple
negative ions all have noble gas structures.
If elements aren't aiming for noble gas structures when they form ions, what decides how
many electrons are transferred? The answer lies in the energetics of the process by which
the compound is made.
***What determines what the charge is on an ion? (for interest only)
Elements combine to make the compound which is as stable as possible - the one in
which the greatest amount of energy is evolved in its making. The more charges a
positive ion has, the greater the attraction towards its accompanying negative ion. The
greater the attraction, the more energy is released when the ions come together.
That means that elements forming positive ions will tend to give away as many electrons
as possible. But there's a down-side to this.
Energy is needed to remove electrons from atoms. This is called ionization energy. The
more electrons you remove, the greater the total ionization energy becomes. Eventually
the total ionization energy needed becomes so great that the energy released when the
attractions are set up between positive and negative ions isn't large enough to cover it.
The element forms the ion which makes the compound most stable - the one in which
most energy is released over-all.
For example, why is calcium chloride CaCl2 rather than CaCl or CaCl3?
If one mole of CaCl (containing Ca+ ions) is made from its elements, it is possible to
estimate that about 171 kJ of heat is evolved.
However, making CaCl2 (containing Ca2+ ions) releases more heat. You get 795 kJ. That
extra amount of heat evolved makes the compound more stable, which is why you get
CaCl2 rather than CaCl.
What about CaCl3 (containing Ca3+ ions)? To make one mole of this, you can estimate
that you would have to put in 1341 kJ. This makes this compound completely non-viable.
Why is so much heat needed to make CaCl3? It is because the third ionisation energy (the
energy needed to remove the third electron) is extremely high (4940 kJ mol-1) because the
electron is being removed from the 3-level rather than the 4-level. Because it is much
closer to the nucleus than the first two electrons removed, it is going to be held much
more strongly. ***
COVALENT BONDING - SINGLE BONDS
This page explains what covalent bonding is. It starts with a simple picture of the single
covalent bond, and then modifies it slightly for IB purposes. It also takes a more
sophisticated view (beyond IB) if you are interested.
A simple view of covalent bonding
The importance of noble gas structures
At a simple level a lot of importance is attached to the electronic structures of noble gases
like neon or argon which have eight electrons in their outer energy levels (or two in the
case of helium). These noble gas structures are thought of as being in some way a
"desirable" thing for an atom to have.
You may well have been left with the strong impression that when other atoms react, they
try to achieve noble gas structures.
As well as achieving noble gas structures by transferring electrons from one atom to
another as in ionic bonding, it is also possible for atoms to reach these stable structures
by sharing electrons to give covalent bonds.
Some very simple covalent molecules
Chlorine
For example, two chlorine atoms could both achieve stable structures by sharing their
single unpaired electron as in the diagram.
The fact that one chlorine has been drawn with electrons marked as crosses and the other
as dots is simply to show where all the electrons come from. In reality there is no
difference between them.
The two chlorine atoms are said to be joined by a covalent bond. The reason that the two
chlorine atoms stick together is that the shared pair of electrons is attracted to the nucleus
of both chlorine atoms.
Hydrogen
Hydrogen atoms only need two electrons in their outer level to reach the noble gas
structure of helium. Once again, the covalent bond holds the two atoms together because
the pair of electrons is attracted to both nuclei.
Hydrogen chloride
The hydrogen has a helium structure, and the chlorine an argon structure.
Covalent bonding (IB)
Cases where there isn't any difference from the simple view
The only thing which must be changed is the over-reliance on the concept of noble gas
structures. Most of the simple molecules you draw do in fact have all their atoms with
noble gas structures.
For example:
Even with a more complicated molecule like PCl3, there's no problem. In this case, only
the outer electrons are shown for simplicity. Each atom in this structure has inner layers
of electrons of 2,8. Again, everything present has a noble gas structure.
Cases where the simple view causes problems
Boron trifluoride, BF3
A boron atom only has 3 electrons in its outer level, and there is no possibility of it
reaching a noble gas structure by simple sharing of electrons. Is this a problem? No. The
boron has formed the maximum number of bonds that it can in the circumstances, and
this is a perfectly valid structure.
Energy is released whenever a covalent bond is formed. Because energy is being lost
from the system, it becomes more stable after every covalent bond is made. It follows,
therefore, that an atom will tend to make as many covalent bonds as possible. In the case
of boron in BF3, three bonds is the maximum possible because boron only has 3 electrons
to share.
Phosphorus(V) chloride, PCl5
In the case of phosphorus 5 covalent bonds are possible - as in PCl5.
Phosphorus forms two chlorides - PCl3 and PCl5. When phosphorus burns in chlorine
both are formed - the majority product depending on how much chlorine is available.
We've already looked at the structure of PCl3.
The diagram of PCl5 (like the previous diagram of PCl3) shows only the outer electrons.
Notice that the phosphorus now has 5 pairs of electrons in the outer level - certainly not a
noble gas structure. You would have been content to draw PCl3 at GCSE, but PCl5 would
have looked very worrying.
Why does phosphorus sometimes break away from a noble gas structure and form five
bonds? In order to answer that question, we need to explore territory beyond the limits of
current IB syllabuses. Don't be put off by this! It isn't particularly difficult, and is
extremely useful if you are going to understand the bonding in some important organic
compounds.
A more sophisticated view of covalent bonding
The bonding in methane, CH4
What is wrong with the dots-and-crosses picture of bonding in methane?
We are starting with methane because it is the simplest case which illustrates the sort of
processes involved. You will remember that the dots-and-crossed picture of methane
looks like this.
There is a serious mis-match between this structure and the modern electronic structure
of carbon, 1s22s22px12py1. The modern structure shows that there are only 2 unpaired
electrons for hydrogens to share with, instead of the 4 which the simple view requires.
You can see this more readily using the electrons-in-boxes
notation. Only the 2-level electrons are shown. The 1s2
electrons are too deep inside the atom to be involved in
bonding. The only electrons directly available for sharing are
the 2p electrons. Why then isn't methane CH2?
Promotion of an electron
When bonds are formed, energy is released and the system
becomes more stable. If carbon forms 4 bonds rather than 2,
twice as much energy is released and so the resulting
molecule becomes even more stable.
There is only a small energy gap between the 2s and 2p
orbitals, and so it pays the carbon to provide a small amount
of energy to promote an electron from the 2s to the empty
2p to give 4 unpaired electrons. The extra energy released
when the bonds form more than compensates for the initial
input.
Now that we've got 4 unpaired electrons ready for bonding, another problem arises. In
methane all the carbon-hydrogen bonds are identical, but our electrons are in two
different kinds of orbitals. You aren't going to get four identical bonds unless you start
from four identical orbitals.
Hybridization
The electrons rearrange themselves again in a process called
hybridization. This reorganizes the electrons into four identical
hybrid orbitals called sp3 hybrids (because they are made from
one s orbital and three p orbitals). You should read "sp3" as "s p
three" - not as "s p cubed".
sp3 hybrid orbitals look a bit like half a p orbital, and they arrange
themselves in space so that they are as far apart as possible. You
can picture the nucleus as being at the centre of a tetrahedron (a
triangularly based pyramid) with the orbitals pointing to the
corners. For clarity, the nucleus is drawn far larger than it really is.
What happens when the bonds are formed?
Remember that hydrogen's electron is in a 1s orbital - a spherically symmetric region of
space surrounding the nucleus where there is some fixed chance (say 95%) of finding the
electron. When a covalent bond is formed, the atomic orbitals (the orbitals in the
individual atoms) merge to produce a new molecular orbital which contains the electron
pair which creates the bond.
Four molecular orbitals are formed, looking rather like the original sp3 hybrids, but with a
hydrogen nucleus embedded in each lobe. Each orbital holds the 2 electrons that we've
previously drawn as a dot and a cross.
The principles involved - promotion of electrons if necessary, then hybridization,
followed by the formation of molecular orbitals - can be applied to any covalently-bound
molecule.
The bonding in the phosphorus chlorides, PCl3 and PCl5
What's wrong with the simple view of PCl3?
This diagram only shows the outer (bonding) electrons.
Nothing is wrong with this! (Although it doesn't account for the shape of the molecule
properly.) If you were going to take a more modern look at it, the argument would go like
this:
Phosphorus has the electronic structure 1s22s22p63s23px13py13pz1. If we look only at the
outer electrons as "electrons-in-boxes":
There are 3 unpaired electrons that can be used to form bonds with 3 chlorine atoms. The
four 3-level orbitals hybridize to produce 4 equivalent sp3 hybrids just like in carbon except that one of these hybrid orbitals contains a lone pair of electrons.
Each of the 3 chlorines then forms a covalent bond by merging the atomic orbital
containing its unpaired electron with one of the phosphorus unpaired electrons to make 3
molecular orbitals.
You might wonder whether all this is worth the bother! Probably not! It is worth it with
PCl5, though.
What's wrong with the simple view of PCl5?
You will remember that the dots-and-crosses picture of PCl5 looks awkward because the
phosphorus doesn't end up with a noble gas structure. This diagram also shows only the
outer electrons.
In this case, a more modern view makes things look better by abandoning any pretence of
worrying about noble gas structures.
If the phosphorus is going to form PCl5 it has first to generate 5 unpaired electrons. It
does this by promoting one of the electrons in the 3s orbital to the next available higher
energy orbital.
Which higher energy orbital? It uses one of the 3d orbitals. You might have expected it to
use the 4s orbital because this is the orbital that fills before the 3d when atoms are being
built from scratch. Not so! Apart from when you are building the atoms in the first place,
the 3d always counts as the lower energy orbital.
This leaves the phosphorus with this arrangement of its electrons:
The 3-level electrons now rearrange (hybridize) themselves to give 5 hybrid orbitals, all
of equal energy. They would be called sp3d hybrids because that's what they are made
from.
The electrons in each of these orbitals would then share space with electrons from five
chlorines to make five new molecular orbitals - and hence five covalent bonds.
Why does phosphorus form these extra two bonds? It puts in an amount of energy to
promote an electron, which is more than paid back when the new bonds form. Put simply,
it is energetically profitable for the phosphorus to form the extra bonds.
The advantage of thinking of it in this way is that it completely ignores the question of
whether you've got a noble gas structure, and so you don't worry about it.
A non-existent compound - NCl5
Nitrogen is in the same Group of the Periodic Table as phosphorus, and you might expect
it to form a similar range of compounds. In fact, it doesn't. For example, the compound
NCl3 exists, but there is no such thing as NCl5.
Nitrogen is 1s22s22px12py12pz1. The reason that NCl5 doesn't exist is that in order to form
five bonds, the nitrogen would have to promote one of its 2s electrons. The problem is
that there aren't any 2d orbitals to promote an electron into - and the energy gap to the
next level (the 3s) is far too great.
In this case, then, the energy released when the extra bonds are made isn't enough to
compensate for the energy needed to promote an electron - and so that promotion doesn't
happen.
Atoms will form as many bonds as possible provided it is energetically profitable.
COVALENT BONDING - DOUBLE BONDS
A simple view of double covalent bonds
A double covalent bond is where two pairs of electrons are shared between the atoms
rather than just one pair.
Some simple molecules containing double bonds
Oxygen, O2
Two oxygen atoms can both achieve stable structures by sharing two pairs of electrons as
in the diagram.
The double bond is shown conventionally by two lines joining the atoms. Each line
represents one pair of shared electrons.
Carbon dioxide, CO2
Ethene, C2H4
Ethene has a double bond between the two carbon atoms.
A more sophisticated view of the bonding in ethene
It is important to explore the bonding in ethene in more detail because it has a direct
impact on its chemistry. Unless you have some understanding of the true nature of the
double bond, you can't really understand the way that ethene behaves.
An orbital view of the bonding in ethene
Ethene is built from hydrogen atoms (1s1) and carbon atoms (1s22s22px12py1).
The carbon atom doesn't have enough unpaired electrons to form the required number of
bonds, so it needs to promote one of the 2s2 pair into the empty 2pz orbital. This is
exactly the same as happens whenever carbon forms bonds - whatever else it ends up
joined to.
Now there's a difference, because each carbon is only joining to three other atoms
(another carbon and two hydrogens) rather than four - as, for example, in methane.
The carbon re-organizes the s orbital and two of the p orbitals to give three new orbitals
with exactly the same energy. This process is called hybridization. The other p orbital is
left unchanged.
The new orbitals formed are called sp2 hybrids, because
they are made by an s orbital and two p orbitals
reorganizing themselves. sp2 orbitals look rather like sp3
orbitals that we discussed in the bonding in methane in
the page on single bonds, except that they are shorter
and fatter. The three sp2 hybrid orbitals arrange
themselves as far apart as possible - which is at 120° to
each other in a plane. The remaining p orbital is at right
angles to them.
The two carbon atoms and four hydrogen atoms would look like this before they joined
together:
The various atomic orbitals which are pointing towards each other now merge to give
molecular orbitals, each containing a bonding pair of
electrons. Molecular orbitals made by end-to-end overlap of
atomic orbitals are called sigma bonds.
The p orbitals on each carbon aren't pointing towards each other, and so we'll leave those
for a moment. In the diagram, the black dots represent the nuclei of the atoms.
Notice that the p orbitals are so close that they are overlapping sideways.
This sideways overlap also creates a molecular orbital, but of a
different kind. In this one the electrons aren't held on the line
between the two nuclei, but above and below the plane of the
molecule. A bond formed in this way is called a pi bond.
For clarity, the sigma bonds are shown using lines - each line representing one pair of
shared electrons. The various sorts of line show the directions the bonds point in. An
ordinary line represents a bond in the plane of the screen (or the paper if you've printed
it), a broken line is a bond going back away from you, and a wedge shows a bond coming
out towards you.
Be clear about what a pi bond is. It is a region of space in which you can find the two
electrons which make up the bond. Those two electrons can live anywhere within that
space. It would be quite misleading to think of one living in the top and the other in the
bottom.
The pi bond dominates the chemistry of ethene. It is very vulnerable to attack - a very
negative region of space above and below the plane of the molecule. It is also somewhat
distant from the control of the nuclei and so is a weaker bond than the sigma bond joining
the two carbons.
CO-ORDINATE (DATIVE COVALENT) BONDING
A covalent bond is formed by two atoms sharing a pair of electrons. The atoms are held
together because the electron pair is attracted by both of the nuclei.
In the formation of a simple covalent bond, each atom supplies one electron to the bond but that doesn't have to be the case. A co-ordinate bond (also called a dative covalent
bond) is a covalent bond (a shared pair of electrons) in which both electrons come from
the same atom.
For the rest of this page, we shall use the term co-ordinate bond - but if you prefer to call
it a dative covalent bond, that's not a problem!
The reaction between ammonia and hydrogen chloride
If these colourless gases are allowed to mix, a thick white smoke of solid ammonium
chloride is formed.
Ammonium ions, NH4+, are formed by the transfer of a hydrogen ion from the hydrogen
chloride to the lone pair of electrons on the ammonia molecule.
When the ammonium ion, NH4+, is formed, the fourth hydrogen is attached by a dative
covalent bond, because only the hydrogen's nucleus is transferred from the chlorine to the
nitrogen. The hydrogen's electron is left behind on the chlorine to form a negative
chloride ion.
Once the ammonium ion has been formed it is impossible to tell any difference between
the dative covalent and the ordinary covalent bonds. Although the electrons are shown
differently in the diagram, there is no difference between them in reality.
Representing co-ordinate bonds
In simple diagrams, a co-ordinate bond is shown by an arrow. The arrow points from the
atom donating the lone pair to the atom accepting it.
Dissolving hydrogen chloride in water to make hydrochloric acid
Something similar happens. A hydrogen ion (H+) is transferred from the chlorine to one
of the lone pairs on the oxygen atom.
The H3O+ ion is variously called the hydroxonium ion, the hydronium ion or the oxonium
ion.
In an introductory chemistry course, whenever you have talked about hydrogen ions (for
example in acids), you have actually been talking about the hydroxonium ion. A raw
hydrogen ion is simply a proton, and is far too reactive to exist on its own in a test tube.
If you write the hydrogen ion as H+(aq), the "(aq)" represents the water molecule that the
hydrogen ion is attached to. When it reacts with something (an alkali, for example), the
hydrogen ion simply becomes detached from the water molecule again.
Note that once the co-ordinate bond has been set up, all the hydrogens attached to the
oxygen are exactly equivalent. When a hydrogen ion breaks away again, it could be any
of the three.
The reaction between ammonia and boron trifluoride, BF3
If you have recently read the page on covalent bonding, you may remember boron
trifluoride as a compound which doesn't have a noble gas structure around the boron
atom. The boron only has 3 pairs of electrons in its bonding level, whereas there would
be room for 4 pairs. BF3 is described as being electron deficient.
The lone pair on the nitrogen of an ammonia molecule can be used to overcome that
deficiency, and a compound is formed involving a co-ordinate bond.
Using lines to represent the bonds, this could be drawn more simply as:
The second diagram shows another way that you might find co-ordinate bonds drawn.
The nitrogen end of the bond has become positive because the electron pair has moved
away from the nitrogen towards the boron - which has therefore become negative. We
shan't use this method again - it's more confusing than just using an arrow.
The structure of aluminum chloride
Aluminum chloride sublimes (turns straight from a solid to a gas) at
about 180°C. If it simply contained ions it would have a very high
melting and boiling point because of the strong attractions between
the positive and negative ions. The implication is that it when it
sublimes at this relatively low temperature, it must be covalent. The
dots-and-crosses diagram shows only the outer electrons.
AlCl3, like BF3, is electron deficient. There is likely to be a similarity, because aluminum
and boron are in the same group of the Periodic Table, as are fluorine and chlorine.
Measurements of the relative formula mass of aluminum chloride show that its formula in
the vapour at the sublimation temperature is not AlCl3, but Al2Cl6. It exists as a dimer
(two molecules joined together). The bonding between the two molecules is co-ordinate,
using lone pairs on the chlorine atoms. Each chlorine atom has 3 lone pairs, but only the
two important ones are shown in the line diagram.
Energy is released when the two co-ordinate bonds are formed, and so the dimer is more
stable than two separate AlCl3 molecules.
The bonding in hydrated metal ions
Water molecules are strongly attracted to ions in solution - the water molecules clustering
around the positive or negative ions. In many cases, the attractions are so great that
formal bonds are made, and this is true of almost all positive metal ions. Ions with water
molecules attached are described as hydrated ions.
Although aluminium chloride is covalent, when it dissolves in water, ions are produced.
Six water molecules bond to the aluminium to give an ion with the formula Al(H2O)63+.
It's called the hexaaquaaluminium ion - which translates as six ("hexa") water molecules
("aqua") wrapped around an aluminium ion.
The bonding in this (and the similar ions formed by the great majority of
other metals) is co-ordinate (dative covalent) using lone pairs on the water
molecules.
Aluminum is 1s22s22p63s23px1. When it forms an Al3+ ion it loses the 3-level electrons to
leave 1s22s22p6.
That means that all the 3-level orbitals are now empty. The aluminum re-organizes
(hybridizes) six of these (the 3s, three 3p, and two 3d) to produce six new orbitals all with
the same energy. These six hybrid orbitals accept lone pairs from six water molecules.
You might wonder why it chooses to use six orbitals rather than four or eight or
whatever. Six is the maximum number of water molecules it is possible to fit around an
aluminum ion (and most other metal ions). By making the maximum number of bonds, it
releases most energy and so becomes most energetically stable.
Only one lone pair is shown on each water molecule. The other lone pair is pointing away
from the aluminum and so isn't involved in the bonding. The resulting ion looks like this:
Because of the movement of electrons towards the centre of the ion, the 3+ charge is no
longer located entirely on the aluminium, but is now spread over the whole of the ion.
Two more molecules
Carbon monoxide, CO
Carbon monoxide can be thought of as having two ordinary covalent bonds between the
carbon and the oxygen plus a co-ordinate bond using a lone pair on the oxygen atom.
Nitric acid, HNO3
In this case, one of the oxygen atoms can be thought of as attaching to the nitrogen via a
co-ordinate bond using the lone pair on the nitrogen atom.
In fact this structure is misleading because it suggests that the two oxygen atoms on the
right-hand side of the diagram are joined to the nitrogen in different ways. Both bonds are
actually identical in length and strength, and so the arrangement of the electrons must be
identical. There is no way of showing this using a dots-and-crosses picture. The bonding
involves delocalization.