Bonding Chemistry

8/4/2019 Bonding Chemistry

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ATOMIC STRUCTURE AND BONDING

IONIC (ELECTROVALENT) BONDING

A simple view of ionic bonding

The importance of noble gas structures

At a simple level (like GCSE) a lot of importance is attached tothe electronic structures of noble gases like neon or argon which

have eight electrons in their outer energy levels (or two in thecase of helium). These noble gas structures are thought of asbeing in some way a "desirable" thing for an atom to have.

You may well have been left with the strong impression thatwhen other atoms react, they try to organise things such thattheir outer levels are either completely full or completely empty.

Note: The central role given to noble gasstructures is very much an over-simplification. We shall have to spend some

time later on demolishing the concept!

Ionic bonding in sodium chloride

Sodium (2,8,1) has 1 electron more than a stable noble gasstructure (2,8). If it gave away that electron it would becomemore stable.

Chlorine (2,8,7) has 1 electron short of a stable noble gasstructure (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 achlorine atom, both become more stable.


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The sodium has lost an electron, so it no longer has equalnumbers of electrons and protons. Because it has one moreproton than electron, it has a charge of 1+. If electrons are lostfrom an atom, positive ions are formed.

Positive ions are sometimes called cations.

The chlorine has gained an electron, so it now has one moreelectron than proton. It therefore has a charge of 1-. If electronsare 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 thestrong electrostatic attractions between the positive andnegative charges.

The formula of sodium chloride

You need one sodium atom to provide the extra electron for onechlorine atom, so they combine together 1:1. The formula istherefore NaCl.

Some other examples of ionic bonding

magnesium oxide

Again, noble gas structures are formed, and the magnesiumoxide is held together by very strong attractions between theions. The ionic bonding is stronger than in sodium chloridebecause this time you have 2+ ions attracting 2- ions. Thegreater the charge, the greater the attraction.


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The formula of magnesium oxide is MgO.

calcium chloride

This time you need two chlorines to use up the two outerelectrons in the calcium. The formula of calcium chloride istherefore CaCl2.

potassium oxide

Again, noble gas structures are formed. It takes two potassiumsto supply the electrons the oxygen needs. The formula ofpotassium oxide is K2O.

THE A'LEVEL VIEW OF IONIC BONDING

Electrons are transferred from one atom to anotherresulting in the formation of positive and negative ions.

The electrostatic attractions between the positive andnegative ions hold the compound together.

So what's new? At heart - nothing. What needs modifying is theview that there is something magic about noble gas structures.There are far more ions which don't have noble gas structuresthan there are which do.

Some common ions which don't have noble gas structures

You may have come across some of the following ions in abasic course like GCSE. They are all perfectly stable , but not


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one of them has a noble gas structure.

Fe3+ [Ar]3d5Cu2+ [Ar]3d9Zn

2+[Ar]3d

10Ag+ [Kr]4d10Pb2+ [Xe]4f145d106s2

Noble gases (apart from helium) have an outer electronicstructure ns2np6.

Note: If you aren't happy aboutwritingelectronic structuresusing of s, p andd notation, follow this link before you go on.Return to this page via the menus or byusing the BACK button on your browser.

Apart from some elements at the beginning of a transition series(scandium forming Sc3+ with an argon structure, for example), alltransition elements and any metals following a transition series(like tin and lead in Group 4, for example) will have structureslike those above.

That means that the only elements to form positive ions withnoble gas structures (apart from odd ones like scandium) arethose in groups 1 and 2 of the Periodic Table and aluminium ingroup 3 (boron in group 3 doesn't form ions).

Negative ions are tidier! Those elements in Groups 5, 6 and 7which form simple negative ions all have noble gas structures.

If elements aren't aiming for noble gas structures when theyform ions, what decides how many electrons are transferred?The answer lies in the energetics of the process by which thecompound is made.

Warning! From here to the bottom of thispage goes beyond anything you are likely toneed for A'level purposes. It is included forinterest only.
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What determines what the charge is on an ion?

Elements combine to make the compound which is as stable as

possible - the one in which the greatest amount of energy isevolved in its making. The more charges a positive ion has, thegreater the attraction towards its accompanying negative ion.The greater the attraction, the more energy is released when theions come together.

That means that elements forming positive ions will tend to giveaway as many electrons as possible. But there's a down-side tothis.

Energy is needed to remove electrons from atoms. This is

calledionisation energy. The more electrons you remove, thegreater the total ionisation energy becomes. Eventually the totalionisation energy needed becomes so great that the energyreleased when the attractions are set up between positive andnegative ions isn't large enough to cover it.

The element forms the ion which makes the compound moststable - the one in which most energy is released over-all.

For example, why is calcium chloride CaCl2 rather than CaCl orCaCl3?

If one mole of CaCl (containing Ca+ ions) is made from itselements, it is possible to estimate that about 171 kJ of heat isevolved.

However, making CaCl2 (containing Ca2+ ions) releases more

heat. You get 795 kJ. That extra amount of heat evolved makesthe compound more stable, which is why you get CaCl2 ratherthan CaCl.

What about CaCl3 (containing Ca3+ ions)? To make one mole of

this, you can estimate that you would have to put in1341 kJ.This makes this compound completely non-viable. Why is somuch heat needed to make CaCl3? It is because the thirdionisation energy (the energy needed to remove the thirdelectron) is extremely high (4940 kJ mol-1) because the electronis being removed from the 3-level rather than the 4-level.Because it is much closer to the nucleus than the first twoelectrons removed, it is going to be held much more strongly.


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Note: It would pay you to readaboutionisation energiesif you really wantto understand this.

You could also go to a standard text book

and investigate Born-Haber Cycles.

A similar sort of argument applies to the negative ion. Forexample, oxygen forms an O2- ion rather than an O- ion or an O3-

ion, because compounds containing the O2- ion turn out to bethe most energetically stable.

COVALENT BONDING - SINGLE BONDS

A simple view of covalent bonding

The importance of noble gas structures

At a simple level (like GCSE) a lot of importance is attached tothe electronic structures of noble gases like neon or argon which

have eight electrons in their outer energy levels (or two in thecase of helium). These noble gas structures are thought of asbeing in some way a "desirable" thing for an atom to have.

You may well have been left with the strong impression thatwhen other atoms react, they try to achieve noble gasstructures.

As well as achieving noble gas structures by transferringelectrons from one atom to another as in ionic bonding, it is alsopossible 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 stablestructures by sharing their single unpaired electron as in the
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diagram.

The fact that one chlorine has been drawn with electronsmarked as crosses and the other as dots is simply to showwhere all the electrons come from. In reality there is nodifference 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 theshared pair of electrons is attracted to the nucleus of both

chlorine atoms.

Hydrogen

Hydrogen atoms only need two electrons in their outer level toreach the noble gas structure of helium. Once again, thecovalent bond holds the two atoms together because the pair ofelectrons is attracted to both nuclei.

Hydrogen chloride

The hydrogen has a helium structure, and the chlorine an argonstructure.

Covalent bonding at A'level

Cases where there isn't any difference from the simple view

If you stick closely to modern A'level syllabuses, there is little


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need to move far from the simple (GCSE) view. The only thingwhich must be changed is the over-reliance on the concept ofnoble gas structures. Most of the simple molecules you draw doin fact have all their atoms with noble gas structures.

For example:

Even with a more complicated molecule like PCl3, there's noproblem. In this case, only the outer electrons are shown forsimplicity. Each atom in this structure has inner layers ofelectrons of 2,8. Again, everything present has a noble gasstructure.

Cases where the simple view throws up problems

Boron trifluoride, BF3

A boron atom only has 3 electrons in its outer level, and there isno possibility of it reaching a noble gas structure by simplesharing of electrons. Is this a problem? No. The boron hasformed the maximum number of bonds that it can in the


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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 morestable after every covalent bond is made. It follows, therefore,

that an atom will tend to make as many covalent bonds aspossible. In the case of boron in BF3, three bonds is themaximum possible because boron only has 3 electrons to share.

Note: You might perhaps wonder why boron doesn't formionic bonds with fluorine instead. Boron doesn't form ionsbecause the total energy needed to remove three electrons toform a B

3+ion is simply too great to be recoverable when

attractions are set up between the boron and fluoride ions.

Phosphorus(V) chloride, PCl5

In the case of phosphorus 5 covalent bonds are possible - as inPCl5.

Phosphorus forms two chlorides - PCl3 and PCl5. Whenphosphorus burns in chlorine both are formed - the majorityproduct depending on how much chlorine is available. We'vealready 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 theouter level - certainly not a noble gas structure. You would havebeen content to draw PCl3 at GCSE, but PCl5 would have lookedvery worrying.

Why does phosphorus sometimes break away from a noble gas


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structure and form five bonds? In order to answer that question,we need to explore territory beyond the limits of current A'levelsyllabuses. Don't be put off by this! It isn't particularly difficult,and is extremely useful if you are going to understand thebonding in some important organic compounds.

A more sophisticated view of covalent bonding

The bonding in methane, CH4

Warning! If you aren't happy with describing electronarrangements in s and p notation, and with the shapes of s andp orbitals, you need to read aboutorbitalsbefore you go on.

Use the BACK button on your browser to return quickly to thispoint.

What is wrong with the dots-and-crosses picture of bondingin methane?

We are starting with methane because it is the simplest casewhich illustrates the sort of processes involved. You willremember 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. Themodern structure shows that there are only 2 unpaired electronsfor hydrogens to share with, instead of the 4 which the simple

view requires.

You can see this more readily using theelectrons-in-boxes notation. Only the 2-
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level electrons are shown. The 1s2electrons are too deep insidethe atom to be involved in bonding. The only electrons directlyavailable for sharing are the 2p electrons. Why then isn'tmethane CH2?

Promotion of an electron

When bonds are formed, energy isreleased and the system becomes morestable. If carbon forms 4 bonds ratherthan 2, twice as much energy is releasedand so the resulting molecule becomeseven more stable.

There is only a small energy gapbetween the 2s and 2p orbitals, and so it

pays the carbon to provide a smallamount of energy to promote an electronfrom the 2s to the empty 2p to give 4unpaired electrons. The extra energy

released when the bonds form more than compensates for theinitial input.

The carbon atom is now said to be in an excited state.

Note: People sometimes worry that the promoted electron isdrawn as an up-arrow, whereas it started as a down-arrow.The reason for this is actually fairly complicated - well beyondthe level we are working at. Just get in the habit of writing it likethis because it makes the diagrams look tidy!


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Now that we've got 4 unpaired electrons ready for bonding,another problem arises. In methane all the carbon-hydrogenbonds are identical, but our electrons are in two different kindsof orbitals. You aren't going to get four identical bonds unlessyou start from four identical orbitals.

Hybridisation

The electrons rearrange themselvesagain in a process called hybridisation.This reorganises the electrons into fouridentical hybrid orbitals called sp3hybrids(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 porbital, and they arrange themselves in spaceso that they are as far apart as possible. Youcan picture the nucleus as being at the centreof a tetrahedron (a triangularly basedpyramid) with the orbitals pointing to the corners. For clarity, thenucleus 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 - aspherically symmetric region of space surrounding the nucleuswhere there is some fixed chance (say 95%) of finding theelectron. When a covalent bond is formed, the atomic orbitals(the orbitals in the individual atoms) merge to produce a newmolecular orbital which contains the electron pair which creates

the bond.


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Four molecular orbitals are formed, looking rather like theoriginal sp3 hybrids, but with a hydrogen nucleus embedded ineach lobe. Each orbital holds the 2 electrons that we'vepreviously drawn as a dot and a cross.

The principles involved - promotion of electrons if necessary,then hybridisation, followed by the formation of molecularorbitals - can be applied to any covalently-bound molecule.

Note: You will find this bit on methane repeated in the organicsection of this site. That article onmethanegoes on to look atthe formation of carbon-carbon single bonds in ethane.

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 theshape of the molecule properly.) If you were going to take a
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more modern look at it, the argument would go like this:

Phosphorus has the electronic structure1s22s22p63s23px

13py13pz

1. If we look only at the outer electronsas "electrons-in-boxes":

There are 3 unpaired electrons that can be used to form bondswith 3 chlorine atoms. The four 3-level orbitals hybridise toproduce 4 equivalent sp3 hybrids just like in carbon - except thatone of these hybrid orbitals contains a lone pair of electrons.

Each of the 3 chlorines then forms a covalent bond by mergingthe atomic orbital containing its unpaired electron with one of thephosphorus unpaired electrons to make 3 molecular orbitals.

You might wonder whether all this is worth the bother! Probablynot! It isworth it with PCl5, though.

What's wrong with the simple view of PCl5?

You will remember that the dots-and-crosses picture ofPCl5 looks awkward because the phosphorus doesn't end upwith a noble gas structure. This diagram also shows only theouter electrons.


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In this case, a more modern view makes things look better byabandoning any pretence of worrying about noble gasstructures.

If the phosphorus is going to form PCl5 it has first to generate 5unpaired electrons. It does this by promoting one of the

electrons in the 3s orbital to the next available higher energyorbital.

Which higher energy orbital? It uses one of the 3d orbitals. Youmight have expected it to use the 4s orbital because this is theorbital that fills before the 3d when atoms are being built fromscratch. Not so! Apart from when you are building the atoms inthe first place, the 3d always counts as the lower energy orbital.

This leaves the phosphorus with this arrangement of itselectrons:

The 3-level electrons now rearrange (hybridise) themselves togive 5 hybrid orbitals, all of equal energy. They would be called


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sp3d hybrids because that's what they are made from.

The electrons in each of these orbitals would then share spacewith electrons from five chlorines to make five new molecularorbitals - and hence five covalent bonds.

Why does phosphorus form these extra two bonds? It puts in anamount of energy to promote an electron, which is more thanpaid back when the new bonds form. Put simply, it isenergetically profitable for the phosphorus to form the extrabonds.

The advantage of thinking of it in this way is that it completelyignores the question of whether you've got a noble gasstructure, and so you don't worry about it.

A non-existent compound - NCl5

Nitrogen is in the same Group of the Periodic Table asphosphorus, and you might expect it to form a similar range ofcompounds. In fact, it doesn't. For example, the compound

NCl3exists, but there is no such thing as NCl5.

Nitrogen is 1s22s22px12py

12pz1. The reason that NCl5 doesn't

exist is that in order to form five bonds, the nitrogen would haveto promote one of its 2s electrons. The problem is that therearen't any 2d orbitals to promote an electron into - and theenergy gap to the next level (the 3s) is far too great.

In this case, then, the energy released when the extra bonds aremade isn'tenough to compensate for the energy needed topromote an electron - and so that promotion doesn't happen.

Atoms will form as many bonds as possible provided it isenergetically profitable.


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CO-ORDINATE (DATIVE COVALENT) BONDING

Co-ordinate (dative covalent) bonding

A covalent bond is formed by two atoms sharing a pair ofelectrons. The atoms are held together because the electronpair is attracted by both of the nuclei.

In the formation of a simple covalent bond, each atom suppliesone electron to the bond - but that doesn't have to be the case.A co-ordinate bond (also called a dative covalent bond) is acovalent bond (a shared pair of electrons) in which bothelectrons 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 aproblem!

The reaction between ammonia and hydrogen chloride

If these colourless gases are allowed to mix, a thick white smokeof 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 onthe ammonia molecule.

When the ammonium ion, NH4+, is formed, the fourth hydrogen


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is attached by a dative covalent bond, because only thehydrogen's nucleus is transferred from the chlorine to thenitrogen. The hydrogen's electron is left behind on the chlorineto form a negative chloride ion.

Once the ammonium ion has been formed it is impossible to tellany difference between the dative covalent and the ordinarycovalent bonds. Although the electrons are shown differently inthe 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 theatom accepting it.

Dissolving hydrogen chloride in water to make hydrochloricacid

Something similar happens. A hydrogen ion (H+) is transferredfrom 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 (such as GCSE), wheneveryou have talked about hydrogen ions (for example in acids), youhave actually been talking about the hydroxonium ion. A raw


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hydrogen ion is simply a proton, and is far too reactive to existon its own in a test tube.

If you write the hydrogen ion as H+(aq), the "(aq)" represents thewater molecule that the hydrogen ion is attached to. When it

reacts with something (an alkali, for example), the hydrogen ionsimply becomes detached from the water molecule again.

Note that once the co-ordinate bond has been set up, all thehydrogens attached to the oxygen are exactly equivalent. Whena 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, youmay remember boron trifluoride as a compound which doesn'thave a noble gas structure around the boron atom. The borononly has 3 pairs of electrons in its bonding level, whereas therewould be room for 4 pairs. BF3 is described as being electrondeficient.

The lone pair on the nitrogen of an ammonia molecule can beused to overcome that deficiency, and a compound is formedinvolving a co-ordinate bond.

Using lines to represent the bonds, this could be drawn moresimply as:


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The second diagram shows another way that you might find co-ordinate bonds drawn. The nitrogen end of the bond hasbecome positive because the electron pair has moved awayfrom the nitrogen towards the boron - which has thereforebecome negative. We shan't use this method again - it's more

confusing than just using an arrow.

The structure of aluminium chloride

Aluminium chloride sublimes (turns straight from asolid to a gas) at about 180C. If it simplycontained ions it would have a very high meltingand boiling point because of the strong attractionsbetween the positive and negative ions. The

implication is that it when it sublimes at thisrelatively 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 asimilarity, because aluminium and boron are in the same groupof the Periodic Table, as are fluorine and chlorine.

Measurements of the relative formula mass of aluminiumchloride show that its formula in the vapour at the sublimationtemperature is not AlCl3, but Al2Cl6. It exists as a dimer (twomolecules joined together). The bonding between the twomolecules is co-ordinate, using lone pairs on the chlorine atoms.Each chlorine atom has 3 lone pairs, but only the two importantones are shown in the line diagram.

Note: The uninteresting electrons on the chlorines have been


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faded in colour to make the co-ordinate bonds show up better.There's nothing special about those two particular lone pairs -they just happen to be the ones pointing in the right direction.

Energy is released when the two co-ordinate bonds are formed,and so the dimer is more stable than two separateAlCl3molecules.

Note: Aluminium chloride is complicated because of the wayit keeps changing its bonding as the temperature increases. Ifyou are interested in exploring this in more detail, you couldhave a look at the page about thePeriod 3 chlorides. It isn'tparticularly relevant to the present page, though.

If you choose to follow this link, use the BACK button on yourbrowser to return quickly to this page later.

The bonding in hydrated metal ions

Water molecules are strongly attracted to ions in solution - thewater 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. Ionswith water molecules attached are described as hydrated ions.

Although aluminium chloride is covalent, when it dissolves inwater, ions are produced. Six water molecules bond to thealuminium to give an ion with the formula Al(H2O)6

3+. It's calledthe hexaaquaaluminium ion - which translates as six ("hexa")water molecules ("aqua") wrapped around analuminium 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 watermolecules.

Aluminium is 1s22s22p63s23px1. When it forms an Al3+ ion it loses
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the 3-level electrons to leave 1s22s22p6.

That means that all the 3-level orbitals are now empty. Thealuminium re-organises (hybridises) six of these (the 3s, three3p, and two 3d) to produce six new orbitals all with the same

energy. These six hybrid orbitals accept lone pairs from sixwater molecules.

You might wonder why it chooses to use six orbitals rather thanfour or eight or whatever. Six is the maximum number of watermolecules it is possible to fit around an aluminium ion (and mostother metal ions). By making the maximum number of bonds, itreleases most energy and so becomes most energeticallystable.

Only one lone pair is shown on each water molecule. The otherlone pair is pointing away from the aluminium and so isn'tinvolved in the bonding. The resulting ion looks like this:

Because of the movement of electrons towards the centre of theion, the 3+ charge is no longer located entirely on the

aluminium, but is now spread over the whole of the ion.

Note: Dotted arrows represent lone pairs coming from watermolecules behind the plane of the screen or paper. Wedgeshaped arrows represent bonds from water molecules in frontof the plane of the screen or paper.


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Two more molecules

Note: It looks as if only one current UK syllabus wants these

two. Check yours! If you haven't got a copy of yoursyllabus,follow this link to find out how to get one.

Carbon monoxide, CO

Carbon monoxide can be thought of as having two ordinarycovalent 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 lonepair on the nitrogen atom.

In fact this structure is misleading because it suggests that thetwo oxygen atoms on the right-hand side of the diagram are

joined to the nitrogen in different ways. Both bonds are actuallyidentical in length and strength, and so the arrangement of theelectrons must be identical. There is no way of showing this
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using a dots-and-crosses picture. The bonding involvesdelocalisation.

ELECTRONEGATIVITY

What is electronegativity

Definition

Electronegativity is a measure of the tendency of an atom toattract a bonding pair of electrons.

The Pauling scale is the most commonly used. Fluorine (themost electronegative element) is assigned a value of 4.0, andvalues range down to caesium and francium which are the leastelectronegative at 0.7.

What happens if two atoms of equal electronegativity bondtogether?

Consider a bond between two atoms, A and B. Each atom maybe forming other bonds as well as the one shown - but these areirrelevant to the argument.

If the atoms are equally electronegative, both have the sametendency to attract the bonding pair of electrons, and so it will befound on averagehalf way between the two atoms. To get abond like this, A and B would usually have to be the same atom.You will find this sort of bond in, for example, H2 or

Cl2 molecules.

Note: It's important to realise that this is an averagepicture.The electrons are actually in a molecular orbital, and aremoving around all the time within that orbital.


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This sort of bond could be thought of as being a "pure" covalentbond - where the electrons are shared evenly between the twoatoms.

What happens if B is slightly more electronegative than A?

B will attract the electron pair rather more than A does.

That means that the B end of the bond has more than its fairshare of electron density and so becomes slightly negative. Atthe same time, the A end (rather short of electrons) becomes

slightly positive. In the diagram, " " (read as "delta") means"slightly" - so + means "slightly positive".

Defining polar bonds

This is described as a polar bond. A polar bond is a covalentbond in which there is a separation of charge between one endand the other - in other words in which one end is slightlypositive and the other slightly negative. Examples include mostcovalent bonds. The hydrogen-chlorine bond in HCl or thehydrogen-oxygen bonds in water are typical.

What happens if B is a lot more electronegative than A?

In this case, the electron pair is dragged right over to B's end ofthe bond. To all intents and purposes, A has lost control of itselectron, and B has complete control over both electrons. Ionshave been formed.

A "spectrum" of bonds

The implication of all this is that there is no clear-cut divisionbetween covalent and ionic bonds. In a pure covalent bond, the


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electrons are held on average exactly half way between theatoms. In a polar bond, the electrons have been dragged slightlytowards one end.

How far does this dragging have to go before the bond counts

as ionic? There is no real answer to that. You normally think ofsodium chloride as being a typically ionic solid, but even herethe sodium hasn't completelylost control of its electron.Because of the properties of sodium chloride, however, we tendto count it as if it were purely ionic.

Note: Don't worry too much about the exact cut-off pointbetween polar covalent bonds and ionic bonds. At A'level,examples will tend to avoid the grey areas - they will beobviously covalent or obviously ionic. You will, however, beexpected to realise that those grey areas exist.

Lithium iodide, on the other hand, would be described as being"ionic with some covalent character". In this case, the pair ofelectrons hasn't moved entirely over to the iodine end of thebond. Lithium iodide, for example, dissolves in organic solventslike ethanol - not something which ionic substances normally do.

Summary

No electronegativity difference between two atoms leadsto a pure non-polar covalent bond.

A small electronegativity difference leads to a polarcovalent bond.

A large electronegativity difference leads to an ionicbond.

Polar bonds and polar molecules

In a simple molecule like HCl, if the bond is polar, so also is thewhole molecule. What about more complicated molecules?

In CCl4, each bond is polar.


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Note: Ordinary lines represent bonds in the plane of thescreen or paper. Dotted lines represent bonds going awayfrom you into the screen or paper. Wedged lines representbonds coming out of the screen or paper towards you.

The molecule as a whole, however, isn't polar - in the sense thatit doesn't have an end (or a side) which is slightly negative and

one which is slightly positive. The whole of the outside of themolecule is somewhat negative, but there is no overallseparation of charge from top to bottom, or from left to right.

By contrast, CHCl3ispolar.

The hydrogen at the top of the molecule is less electronegativethan carbon and so is slightly positive. This means that themolecule now has a slightly positive "top" and a slightly negative"bottom", and so is overall a polar molecule.

A polar molecule will need to be "lop-sided" in some way.

Patterns of electronegativity in the Periodic Table

The most electronegative element is fluorine. If you rememberthat fact, everything becomes easy, because electronegativitymust always increase towards fluorine in the Periodic Table.


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Note: This simplification ignores the noble gases.Historically this is because they were believed not to formbonds - and if they don't form bonds, they can't have anelectronegativity value. Even now that we know that some ofthem do form bonds, data sources still don't quoteelectronegativity values for them.

Trends in electronegativity across a period

As you go across a period the electronegativity increases. Thechart shows electronegativities from sodium to chlorine - youhave to ignore argon. It doesn't have an electronegativity,because it doesn't form bonds.

Trends in electronegativity down a group

As you go down a group, electronegativity decreases. (If itincreases up to fluorine, it must decrease as you go down.) Thechart shows the patterns of electronegativity in Groups 1 and 7.


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Explaining the patterns in electronegativity

The attraction that a bonding pair of electrons feels for aparticular nucleus depends on:

the number of protons in the nucleus; the distance from the nucleus; the amount of screening by inner electrons.

Note: If you aren't happy about the conceptof screeningorshielding, it would pay you to read the pageonionisation energiesbefore you go on. The factorsinfluencing ionisation energies are just the same as thoseinfluencing electronegativities.

Use the BACK button on your browser to return to this page.

Why does electronegativity increase across a period?

Consider sodium at the beginning of period 3 and chlorine at theend (ignoring the noble gas, argon). Think of sodium chloride asif it were covalently bonded.

Both sodium and chlorine have their bonding electrons in the 3-level. The electron pair is screened from both nuclei by the 1s,2s and 2p electrons, but the chlorine nucleus has 6 more
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protons in it. It is no wonder the electron pair gets dragged so fartowards the chlorine that ions are formed.

Electronegativity increases across a period because the numberof charges on the nucleus increases. That attracts the bonding

pair of electrons more strongly.

Why does electronegativity fall as you go down a group?

Think of hydrogen fluoride and hydrogen chloride.

The bonding pair is shielded from the fluorine's nucleus only bythe 1s2 electrons. In the chlorine case it is shielded by all the1s22s22p6 electrons.

In each case there is a net pull from the centre of the fluorine orchlorine of +7. But fluorine has the bonding pair in the 2-levelrather than the 3-level as it is in chlorine. If it is closer to thenucleus, the attraction is greater.

As you go down a group, electronegativity decreases because

the bonding pair of electrons is increasingly distant from theattraction of the nucleus.

Warning! As far as I am aware, none of the UK-based Alevel (or equivalent) syllabuses any longer want the next bit.It used to be on the AQA syllabus, but has been removedfrom their new syllabus. At the time of writing, it does,however, still appear on at least one overseas A levelsyllabus (Malta, but there may be others that I'm not awareof). If in doubt, check your syllabus.

Otherwise, ignore the rest of this page. It is an alternative(and, to my mind, more awkward) way of looking at theformation of a polar bond. Reading it unnecessarily just risksconfusing you.


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The polarising ability of positive ions

What do we mean by "polarising ability"?

In the discussion so far, we've looked at the formation of polar

bonds from the point of view of the distortions which occur in acovalent bond if one atom is more electronegative than theother. But you can also look at the formation of polar covalentbonds by imagining that you start from ions.

Solid aluminium chloride is covalent. Imagine instead that it wasionic. It would contain Al3+ and Cl- ions.

The aluminium ion is very small and is packed with threepositive charges - the "charge density" is therefore very high.That will have a considerable effect on any nearby electrons.

We say that the aluminium ions polarise the chloride ions.

In the case of aluminium chloride, the electron pairs are draggedback towards the aluminium to such an extent that the bondsbecome covalent. But because the chlorine is moreelectronegative than aluminium, the electron pairs won't bepulled half way between the two atoms, and so the bond formedwill be polar.

Factors affecting polarising ability

Positive ions can have the effect of polarising (electricallydistorting) nearby negative ions. The polarising ability dependson the charge density in the positive ion.

Polarising ability increases as the positive ion gets smaller andthe number of charges gets larger.


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As a negative ion gets bigger, it becomes easier to polarise. Forexample, in an iodide ion, I-, the outer electrons are in the 5-level - relatively distant from the nucleus.

A positive ion would be more effective in attracting a pair of

electrons from an iodide ion than the corresponding electrons in,say, a fluoride ion where they are much closer to the nucleus.

Aluminium iodide is covalent because the electron pair is easilydragged away from the iodide ion. On the other hand, aluminiumfluoride is ionic because the aluminium ion can't polarise thesmall fluoride ion sufficiently to form a covalent bond.

SHAPES OF MOLECULES AND IONS

The examples on this page are all simple in the sense that theyonly contain two sorts of atoms joined by single bonds - forexample, ammonia only contains a nitrogen atom joined to threehydrogen atoms by single bonds. If you are given a morecomplicated example, look carefully at the arrangement of theatoms before you start to make sure that there are only singlebonds present.

For example, if you had a molecule such as COCl2, you wouldneed to work out its structure, based on the fact that you knowthat carbon forms 4 covalent bonds, oxygen 2, and chlorine(normally) 1. If you did that, you would find that the carbon is

joined to the oxygen by a double bond, and to the two chlorinesby single bonds.

That means that you couldn't use the techniques on this page,because this page only considers single bonds.

The electron pair repulsion theory

The shape of a molecule or ion is governed by the arrangementof the electron pairs around the central atom. All you need to dois to work out how many electron pairs there are at the bondinglevel, and then arrange them to produce the minimum amount of


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repulsion between them. You have to include both bonding pairsand lone pairs.

How to work out the number of electron pairs

You can do this by drawing dots-and-crosses pictures, or byworking out the structures of the atoms using electrons-in-boxesand worrying about promotion, hybridisation and so on. But thisis all very tedious! You can get exactly the same information in amuch quicker and easier way for the examples you will meet ifyou are doing one of the UK-based exams for 16 - 18 year olds.

Warning: This method won't work without somemodification for many ions containing metals, and no simplemethod gives reliable results where the central atom is atransition metal. The method will, however, cope with all the

substances that you are likely to meet in this section of thesyllabus. When you deal with transition metal chemistry, youwill be expected to know the shapes of some ions formed bytransition metals, but not to work them out. At that point,learn the ones your syllabus wants you to know.

It is important to know exactly which molecules and ionsyour syllabus expects you to be able to work out the shapesfor in this part of the syllabus. You should also check pastexam papers. If you are working to a UK-based syllabus for16 - 18 year olds, and haven't got copies of yoursyllabusand past papersfollow this link to find out how to get them.

First you need to work out how many electrons there are aroundthe central atom:

Write down the number of electrons in the outer level ofthe central atom. That will be the same as the PeriodicTable group number, except in the case of the noblegases which form compounds, when it will be 8.

Add one electron for each bond being formed. (Thisallows for the electrons coming from the other atoms.)

Allow for any ion charge. For example, if the ion has a 1-charge, add one more electron. For a 1+ charge, deductan electron.

Now work out how many bonding pairs and lone pairs ofelectrons there are:
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Divide by 2 to find the total number of electron pairsaround the central atom.

Work out how many of these are bonding pairs, and howmany are lone pairs. You know how many bonding pairsthere are because you know how many other atoms are

joined to the central atom (assuming that only singlebonds are formed).

For example, if you have 4 pairs of electrons but only 3bonds, there must be 1 lone pair as well as the 3 bondingpairs.

Finally, you have to use this information to work out the shape:

Arrange these electron pairs in space to minimiserepulsions. How this is done will become clear in the

examples which follow.

Don't panic! This is all much easier to do in practice than itis to describe in a long list like this one!

Two electron pairs around the central atom

The only simple case of this is beryllium chloride, BeCl2. The

electronegativity difference between beryllium and chlorine isn'tenough to allow the formation of ions.

Beryllium has 2 outer electrons because it is in group 2. It formsbonds to two chlorines, each of which adds another electron tothe outer level of the beryllium. There is no ionic charge to worryabout, so there are 4 electrons altogether - 2 pairs.

It is forming 2 bonds so there are no lone pairs. The two bondingpairs arrange themselves at 180 to each other, because that'sas far apart as they can get. The molecule is described as

beinglinear.


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Three electron pairs around the central atom

The simple cases of this would be BF3 or BCl3.

Boron is in group 3, so starts off with 3 electrons. It is forming 3bonds, adding another 3 electrons. There is no charge, so thetotal is 6 electrons - in 3 pairs.

Because it is forming 3 bonds there can be no lone pairs. The 3pairs arrange themselves as far apart as possible. They all lie inone plane at 120 to each other. The arrangement iscalledtrigonal planar.

In the diagram, the other electrons on the fluorines have beenleft out because they are irrelevant.

Four electron pairs around the central atom

There are lots of examples of this. The simplest is methane,CH4.

Note: Elsewhere on the site, you will find the shape ofmethane worked out in detail using modern bonding theory.Here we are doing it the quick and easy way!

If you are interested in thebonding in methaneyou can findit in the organic section by following this link, or in a pageoncovalent bondingby following this one.

Carbon is in group 4, and so has 4 outer electrons. It is forming4 bonds to hydrogens, adding another 4 electrons - 8 altogether,in 4 pairs. Because it is forming 4 bonds, these must all bebonding pairs.
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Four electron pairs arrange themselves in space in what iscalled a tetrahedralarrangement. A tetrahedron is a regulartriangularly-based pyramid. The carbon atom would be at thecentre and the hydrogens at the four corners. All the bondangles are 109.5.

Note: It is important that you understand the use of varioussorts of line to show the 3-dimensional arrangement of thebonds. In diagrams of this sort, an ordinary line represents abond in the plane of the screen or paper. A dotted line

shows a bond going away from you into the screen or paper.A wedge shows a bond coming out towards you.

It is my habit to draw diagrams like this with the bond at thetop in the plane of the paper, the middle bond at the bottomcoming out towards you, and the other two going back in.But that's all it is - a habit! You can equally well draw itdifferently if you rotate the molecule a bit. This is alldescribed in some detail about half-way down the pageaboutdrawing organic molecules.

Use the BACK button on your browser to return here later ifyou choose to follow this link.

Other examples with four electron pairs around the centralatom

Ammonia, NH3

Nitrogen is in group 5 and so has 5 outer electrons. Each of the3 hydrogens is adding another electron to the nitrogen's outerlevel, making a total of 8 electrons in 4 pairs. Because thenitrogen is only forming 3 bonds, one of the pairs must be a lonepair. The electron pairs arrange themselves in a tetrahedralfashion as in methane.
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In this case, an additional factor comes into play. Lone pairs arein orbitals that are shorter and rounder than the orbitals that thebonding pairs occupy. Because of this, there is more repulsionbetween a lone pair and a bonding pair than there is betweentwo bonding pairs.

That forces the bonding pairs together slightly - reducing thebond angle from 109.5 to 107. It's not much, but the examinerswill expect you to know it.

Remember this:

Greatest repulsion lone pair - lone pairlone pair - bond pair

Least repulsion bond pair - bond pairBe very careful when you describe the shape of ammonia.Although the electron pair arrangement is tetrahedral, when youdescribe the shape, you only take notice of the atoms. Ammoniaispyramidal- like a pyramid with the three hydrogens at thebase and the nitrogen at the top.

Water, H2O

Following the same logic as before, you will find that the oxygenhas four pairs of electrons, two of which are lone pairs. Thesewill again take up a tetrahedral arrangement. This time the bondangle closes slightly more to 104, because of the repulsion ofthe two lone pairs.

The shape isn't described as tetrahedral, because we only "see"


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the oxygen and the hydrogens - not the lone pairs. Water isdescribed as bentor V-shaped.

The ammonium ion, NH4+

The nitrogen has 5 outer electrons, plus another 4 from the fourhydrogens - making a total of 9.

But take care! This is a positive ion. It has a 1+ charge becauseit has lost 1 electron. That leaves a total of 8 electrons in theouter level of the nitrogen. There are therefore 4 pairs, all ofwhich are bonding because of the four hydrogens.

The ammonium ion has exactly the same shape as methane,

because it has exactly the same electronic arrangement. NH4+

istetrahedral.

Note: To simplify diagrams, bonding electrons won't beshown from now on. Each line, of course, represents abonding pair. It is essential, however, to draw lone pairs.

Methane and the ammonium ion are said tobe isoelectronic. Two species (atoms, molecules or ions) areisoelectronic if they have exactly the same number andarrangement of electrons (including the distinction betweenbonding pairs and lone pairs).

The hydroxonium ion, H3O+

Oxygen is in group 6 - so has 6 outer electrons. Add 1 for eachhydrogen, giving 9. Take one off for the +1 ion, leaving 8. Thisgives 4 pairs, 3 of which are bond pairs. The hydroxonium ion isisoelectronic with ammonia, and has an identical shape -


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pyramidal.

Five electron pairs around the central atom

A simple example: phosphorus(V) fluoride, PF5

(The argument for phosphorus(V) chloride, PCl5, would be

identical.)

Phosphorus (in group 5) contributes 5 electrons, and the fivefluorines 5 more, giving 10 electrons in 5 pairs around thecentral atom. Since the phosphorus is forming five bonds, therecan't be any lone pairs.

The 5 electron pairs take up a shape described as a trigonalbipyramid- three of the fluorines are in a plane at 120 to eachother; the other two are at right angles to this plane. The trigonalbipyramid therefore has two different bond angles - 120 and

90.

A tricky example, ClF3

Chlorine is in group 7 and so has 7 outer electrons. The threefluorines contribute one electron each, making a total of 10 - in 5pairs. The chlorine is forming three bonds - leaving you with 3bonding pairs and 2 lone pairs, which will arrange themselvesinto a trigonal bipyramid.


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But don't jump to conclusions. There are actually three differentways in which you could arrange 3 bonding pairs and 2 lonepairs into a trigonal bipyramid. The right arrangement will be theone with the minimum amount of repulsion - and you can'tdecide that without first drawing all the possibilities.

These are the only possible arrangements. Anything else youmight think of is simply one of these rotated in space.

We need to work out which of these arrangements has theminimum amount of repulsion between the various electronpairs.

A new rule applies in cases like this:

If you have more than four electron pairs arrangedaround the central atom, you can ignorerepulsions at angles of greater than 90.

One of these structures has a fairly obvious large amount ofrepulsion.

In this diagram, two lone pairs are at 90 to each other, whereasin the other two cases they are at more than 90, and so theirrepulsions can be ignored. ClF3 certainly won't take up thisshape because of the strong lone pair-lone pair repulsion.

To choose between the other two, you need to count up eachsort of repulsion.

In the next structure, each lone pair is at 90 to 3 bond pairs,


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and so each lone pair is responsible for 3 lone pair-bond pairrepulsions.

Because of the two lone pairs there are therefore 6 lone pair-bond pair repulsions. And that's all. The bond pairs are at anangle of 120 to each other, and their repulsions can be ignored.

Now consider the final structure.

Each lone pair is at 90 to 2 bond pairs - the ones above andbelow the plane. That makes a total of 4 lone pair-bond pairrepulsions - compared with 6 of these relatively strongrepulsions in the last structure. The other fluorine (the one in theplane) is 120 away, and feels negligible repulsion from the lonepairs.

The bond to the fluorine in the plane is at 90 to the bondsabove and below the plane, so there are a total of 2 bond pair-bond pair repulsions.

The structure with the minimum amount of repulsion is thereforethis last one, because bond pair-bond pair repulsion is less thanlone pair-bond pair repulsion. ClF3 is described as T-shaped.

Warning! If your syllabus expects you to discuss exampleswith more than 4 pairs of electrons around the central atom,

check past exam papers to see if nasty questions like thisone involving ClF3 ever come up. If so, don't leave thisexample until you are sure that you understand it. It is by farthe most complicated one on this page.


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Six electron pairs around the central atom

A simple example: SF6

6 electrons in the outer level of the sulphur, plus 1 each from thesix fluorines, makes a total of 12 - in 6 pairs. Because thesulphur is forming 6 bonds, these are all bond pairs. Theyarrange themselves entirely at 90, in a shape describedas octahedral.

Two slightly more difficult examples

XeF4

Xenon forms a range of compounds, mainly with fluorine oroxygen, and this is a typical one. Xenon has 8 outer electrons,plus 1 from each fluorine - making 12 altogether, in 6 pairs.

There will be 4 bonding pairs (because of the four fluorines) and2 lone pairs.

There are two possible structures, but in one of them the lonepairs would be at 90. Instead, they go opposite each other.XeF4 is described as square planar.

ClF4-

Chlorine is in group 7 and so has 7 outer electrons. Plus the 4from the four fluorines. Plus one because it has a 1- charge.That gives a total of 12 electrons in 6 pairs - 4 bond pairs and 2


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lone pairs. The shape will be identical with that of XeF4.

METALLIC BONDING

What is a metallic bond?

Metallic bonding in sodium

Metals tend to have high melting points and boiling pointssuggesting strong bonds between the atoms. Even a metal like

sodium (melting point 97.8C) melts at a considerably highertemperature than the element (neon) which precedes it in thePeriodic Table.

Sodium has the electronic structure 1s22s22p63s1. When sodiumatoms come together, the electron in the 3s atomic orbital of onesodium atom shares space with the corresponding electron on aneighbouring atom to form a molecular orbital - in much thesame sort of way that a covalent bond is formed.

The difference, however, is that each sodium atom is being

touched by eight other sodium atoms - and the sharing occursbetween the central atom and the 3s orbitals on all of the eightother atoms. And each of these eight is in turn being touched byeight sodium atoms, which in turn are touched by eight atoms -and so on and so on, until you have taken in all the atoms in thatlump of sodium.

All of the 3s orbitals on all of the atoms overlap to give a vastnumber of molecular orbitals which extend over the whole pieceof metal. There have to be huge numbers of molecular orbitals,of course, because any orbital can only hold two electrons.

The electrons can move freely within these molecular orbitals,and so each electron becomes detached from its parent atom.The electrons are said to be delocalised. The metal is heldtogether by the strong forces of attraction between the positivenuclei and the delocalised electrons.


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This is sometimes described as "an array of positive ions in asea of electrons".

If you are going to use this view, beware! Is a metal made up ofatoms or ions? It is made of atoms.

Each positive centre in the diagram represents all the rest of theatom apart from the outer electron, but that electron hasn't beenlost - it may no longer have an attachment to a particular atom,but it's still there in the structure. Sodium metal is thereforewritten as Na - notNa+.

Metallic bonding in magnesium

If you work through the same argument with magnesium, youend up with stronger bonds and so a higher melting point.

Magnesium has the outer electronic structure 3s2. Both of these

electrons become delocalised, so the "sea" has twice theelectron density as it does in sodium. The remaining "ions" alsohave twice the charge (if you are going to use this particularview of the metal bond) and so there will be more attractionbetween "ions" and "sea".

More realistically, each magnesium atom has one more protonin the nucleus than a sodium atom has, and so not only willthere be a greater number of delocalised electrons, but there willalso be a greater attraction for them.

Magnesium atoms have a slightly smaller radius than sodiumatoms, and so the delocalised electrons are closer to the nuclei.Each magnesium atom also has twelve near neighbours ratherthan sodium's eight. Both of these factors increase the strengthof the bond still further.


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Metallic bonding in transition elements

Transition metals tend to have particularly high melting pointsand boiling points. The reason is that they can involve the 3delectrons in the delocalisation as well as the 4s. The more

electrons you can involve, the stronger the attractions tend tobe.

Note: If you aren't happy about theelectronicstructureof transition metals, then you might liketo follow this link to revise it.

The metallic bond in molten metals

In a molten metal, the metallic bond is still present, although theordered structure has been broken down. The metallic bond isn'tfully broken until the metal boils. That means that boiling point isactually a better guide to the strength of the metallic bond thanmelting point is. On melting, the bond is loosened, not broken.

INTERMOLECULAR BONDING - VAN DER WAALS FORCES

What are intermolecular attractions?

Intermolecular versus intramolecular bonds

Intermolecularattractions are attractions between onemolecule and a neighbouring molecule. The forces of attractionwhich hold an individual molecule together (for example, thecovalent bonds) are known as intramolecularattractions.These two words are so confusingly similar that it is safer toabandon one of them and never use it. The term"intramolecular" won't be used again on this site.

All molecules experience intermolecular attractions, although insome cases those attractions are very weak. Even in a gas likehydrogen, H2, if you slow the molecules down by cooling thegas, the attractions are large enough for the molecules to stick
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together eventually to form a liquid and then a solid.

In hydrogen's case the attractions are so weak that themolecules have to be cooled to 21 K (-252C) before theattractions are enough to condense the hydrogen as a liquid.

Helium's intermolecular attractions are even weaker - themolecules won't stick together to form a liquid until thetemperature drops to 4 K (-269C).

van der Waals forces: dispersion forces

Dispersion forces (one of the two types of van der Waals forcewe are dealing with on this page) are also known as "Londonforces" (named after Fritz London who first suggested how they

might arise).

The origin of van der Waals dispersion forces

Temporary fluctuating dipoles

Attractions are electrical in nature. In a symmetrical moleculelike hydrogen, however, there doesn't seem to be any electricaldistortion to produce positive or negative parts. But that's onlytrue on average.

The lozenge-shaped diagram represents a small symmetricalmolecule - H2, perhaps, or Br2. The even shading shows that onaverage there is no electrical distortion.

But the electrons are mobile, and at any one instant they mightfind themselves towards one end of the molecule, making thatend -. The other end will be temporarily short of electrons andso becomes +.

Note: (read as "delta") means "slightly" - so + means"slightly positive".


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An instant later the electrons may well have moved up to theother end, reversing the polarity of the molecule.

This constant "sloshing around" of the electrons in the moleculecauses rapidly fluctuating dipoles even in the most symmetricalmolecule. It even happens in monatomic molecules - moleculesof noble gases, like helium, which consist of a single atom.

If both the helium electrons happen to be on one side of theatom at the same time, the nucleus is no longer properly

covered by electrons for that instant.

How temporary dipoles give rise to intermolecularattractions

I'm going to use the same lozenge-shaped diagram now torepresent anymolecule which could, in fact, be a much morecomplicated shape. Shape does matter (see below), but keepingthe shape simple makes it a lot easier to both draw the diagramsand understand what is going on.

Imagine a molecule which has a temporary polarity beingapproached by one which happens to be entirely non-polar justat that moment. (A pretty unlikely event, but it makes thediagrams much easier to draw! In reality, one of the molecules islikely to have a greater polarity than the other at that time - and

so will be the dominant one.)

As the right hand molecule approaches, its electrons will tend tobe attracted by the slightly positive end of the left hand one.


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This sets up an induced dipolein the approaching molecule,which is orientated in such a way that the + end of one isattracted to the - end of the other.

An instant later the electrons in the left hand molecule may wellhave moved up the other end. In doing so, they will repel theelectrons in the right hand one.

The polarity of both molecules reverses, but you still have +attracting -. As long as the molecules stay close to each otherthe polarities will continue to fluctuate in synchronisation so thatthe attraction is always maintained.

There is no reason why this has to be restricted to twomolecules. As long as the molecules are close together thissynchronised movement of the electrons can occur over hugenumbers of molecules.

This diagram shows how a whole lattice of molecules could beheld together in a solid using van der Waals dispersion forces.

An instant later, of course, you would have to draw a quitedifferent arrangement of the distribution of the electrons as theyshifted around - but always in synchronisation.

The strength of dispersion forces


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Dispersion forces between molecules are much weaker than thecovalent bonds within molecules. It isn't possible to give anyexact value, because the size of the attraction variesconsiderably with the size of the molecule and its shape.

How molecular size affects the strength of the dispersionforces

The boiling points of the noble gases are

helium -269Cneon -246Cargon -186Ckrypton -152Cxenon -108Cradon -62C

All of these elements have monatomic molecules.

The reason that the boiling points increase as you go down thegroup is that the number of electrons increases, and so alsodoes the radius of the atom. The more electrons you have, andthe more distance over which they can move, the bigger thepossible temporary dipoles and therefore the bigger thedispersion forces.

Because of the greater temporary dipoles, xenon molecules are"stickier" than neon molecules. Neon molecules will break awayfrom each other at much lower temperatures than xenonmolecules - hence neon has the lower boiling point.

This is the reason that (all other things being equal) biggermolecules have higher boiling points than small ones. Biggermolecules have more electrons and more distance over whichtemporary dipoles can develop - and so the bigger moleculesare "stickier".


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How molecular shape affects the strength of the dispersionforces

The shapes of the molecules also matter. Long thin molecules

can develop bigger temporary dipoles due to electron movementthan short fat ones containing the same numbers of electrons.

Long thin molecules can also lie closer together - theseattractions are at their most effective if the molecules are reallyclose.

For example, the hydrocarbon molecules butane and 2-methylpropane both have a molecular formula C4H10, but theatoms are arranged differently. In butane the carbon atoms arearranged in a single chain, but 2-methylpropane is a shorter

chain with a branch.

Butane has a higher boiling point because the dispersion forcesare greater. The molecules are longer (and so set up biggertemporary dipoles) and can lie closer together than the shorter,

fatter 2-methylpropane molecules.

van der Waals forces: dipole-dipole interactions

Warning! There's a bit of a problem here with modernsyllabuses. The majority of the syllabuses talk as if dipole-dipole interactions were quite distinct from van der Waalsforces. Such a syllabus will talk about van der Waalsforces (meaning dispersion forces) and, separately, dipole-dipole interactions.

Allintermolecular attractions are known collectively as vander Waals forces. The various different types were firstexplained by different people at different times. Dispersionforces, for example, were described by London in 1930;dipole-dipole interactions by Keesom in 1912.

This oddity in the syllabuses doesn't matter in the least asfar as understanding is concerned - but you obviously


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must know what your particular examiners mean by theterms they use in the questions. Check your syllabus.

If you are working to a UK-based syllabus for 16 - 18 yearolds, but don't have a copy of it,follow this linkto find outhow to get one.

A molecule like HCl has a permanent dipole because chlorine ismore electronegative than hydrogen. These permanent, in-builtdipoles will cause the molecules to attract each other rathermore than they otherwise would if they had to rely only ondispersion forces.

Note: If you aren't happy aboutelectronegativity and

polar molecules, follow this link before you go on.

It's important to realise that all molecules experience dispersionforces. Dipole-dipole interactions are not an alternative todispersion forces - they occur in addition to them. Moleculeswhich have permanent dipoles will therefore have boiling pointsrather higher than molecules which only have temporaryfluctuating dipoles.

Surprisingly dipole-dipole attractions are fairly minor comparedwith dispersion forces, and their effect can only really be seen ifyou compare two molecules with the same number of electronsand the same size. For example, the boiling points of ethane,CH3CH3, and fluoromethane, CH3F, are

Why choose these two molecules to compare? Both haveidentical numbers of electrons, and if you made models youwould find that the sizes were similar - as you can see in the
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diagrams. That means that the dispersion forces in bothmolecules should be much the same.

The higher boiling point of fluoromethane is due to the largepermanent dipole on the molecule because of the high

electronegativity of fluorine. However, even given the largepermanent polarity of the molecule, the boiling point has onlybeen increased by some 10.

Here is another example showing thedominance of the dispersion forces.Trichloromethane, CHCl3, is a highly polarmolecule because of the electronegativity ofthe three chlorines. There will be quite

strong dipole-dipole attractions between onemolecule and its neighbours.

On the other hand, tetrachloromethane, CCl4, is non-polar. Theoutside of the molecule is uniformly - in all directions. CCl4 hasto rely only on dispersion forces.

So which has the highest boiling point? CCl4 does, because it isa bigger molecule with more electrons. The increase in the

dispersion forces more than compensates for the loss of dipole-dipole interactions.

The boiling points are:

CHCl3 61.2CCCl4 76.8C


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INTERMOLECULAR BONDING - HYDROGEN BONDS

This page explains the origin of hydrogen bonding - a relativelystrong form of intermolecular attraction. If you are also interestedin the weaker intermolecular forces (van der Waals dispersionforces and dipole-dipole interactions), there is a link at thebottom of the page.

The evidence for hydrogen bonding

Many elements form compounds with hydrogen. If you plot theboiling points of the compounds of the Group 4 elements withhydrogen, you find that the boiling points increase as you godown the group.

The increase in boiling point happens because the moleculesare getting larger with more electrons, and so van der Waalsdispersion forces become greater.

Note: If you aren't sure aboutvan der Waals dispersionforces, it would pay you to follow this link before you go on.

If you repeat this exercise with the compounds of the elementsin Groups 5, 6 and 7 with hydrogen, something odd happens.
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Although for the most part the trend is exactly the same as ingroup 4 (for exactly the same reasons), the boiling point of thecompound of hydrogen with the first element in each group is

abnormally high.

In the cases of NH3, H2O and HF there must be some additionalintermolecular forces of attraction, requiring significantly moreheat energy to break. These relatively powerful intermolecularforces are described as hydrogen bonds.

The origin of hydrogen bonding

The molecules which have this extra bonding are:

Note: The solid line represents a bond in the plane of the

screen or paper. Dotted bonds are going back into the screenor paper away from you, and wedge-shaped ones are comingout towards you.

Notice that in each of these molecules:

The hydrogen is attached directly to one of the most


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electronegative elements, causing the hydrogen toacquire a significant amount of positive charge.

Each of the elements to which the hydrogen is attached isnot only significantly negative, but also has at least one"active" lone pair.

Lone pairs at the 2-level have the electrons contained ina relatively small volume of space which therefore has ahigh density of negative charge. Lone pairs at higherlevels are more diffuse and not so attractive to positivethings.

Note: If you aren't happy aboutelectronegativity, you shouldfollow this link before you go on.

Consider two water molecules coming close together.

The + hydrogen is so strongly attracted to the lone pair that it isalmost as if you were beginning to form a co-ordinate (dativecovalent) bond. It doesn't go that far, but the attraction issignificantly stronger than an ordinary dipole-dipole interaction.

Hydrogen bonds have about a tenth of the strength of anaverage covalent bond, and are being constantly broken and

reformed in liquid water. If you liken the covalent bond betweenthe oxygen and hydrogen to a stable marriage, the hydrogenbond has "just good friends" status. On the same scale, van derWaals attractions represent mere passing acquaintances!

Water as a "perfect" example of hydrogen bonding

Notice that each water molecule can potentially form four
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hydrogen bonds with surrounding water molecules. There areexactly the right numbers of + hydrogens and lone pairs so thatevery one of them can be involved in hydrogen bonding.

This is why the boiling point of water is higher than that of

ammonia or hydrogen fluoride. In the case of ammonia, theamount of hydrogen bonding is limited by the fact that eachnitrogen only has one lone pair. In a group of ammoniamolecules, there aren't enough lone pairs to go around to satisfyall the hydrogens.

In hydrogen fluoride, the problem is a shortage of hydrogens. Inwater, there are exactly the right number of each. Water couldbe considered as the "perfect" hydrogen bonded system.

Note: You will find more discussion on the effect of hydrogen

bonding on the properties of water in the page onmolecularstructures.

More complex examples of hydrogen bonding

The hydration of negative ions

When an ionic substance dissolves in water, water moleculescluster around the separated ions. This process is calledhydration.

Water frequently attaches to positive ions by co-ordinate (dativecovalent) bonds. It bonds to negative ions using hydrogenbonds.

Note: If you are interested in the bonding in hydrated positiveions, you could follow this link toco-ordinate (dative covalent)bonding.

The diagram shows the potential hydrogen bonds formed to achloride ion, Cl-. Although the lone pairs in the chloride ion are atthe 3-level and wouldn't normally be active enough to formhydrogen bonds, in this case they are made more attractive by
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the full negative charge on the chlorine.

However complicated the negative ion, there will always be lone

pairs that the hydrogen atoms from the water molecules canhydrogen bond to.

Hydrogen bonding in alcohols

An alcohol is an organic molecule containing an -O-H group.

Any molecule which has a hydrogen atom attached directly to anoxygen or a nitrogen is capable of hydrogen bonding. Such

molecules will always have higher boiling points than similarlysized molecules which don't have an -O-H or an -N-H group.The hydrogen bonding makes the molecules "stickier", and moreheat is necessary to separate them.

Ethanol, CH3CH2-O-H, and methoxymethane, CH3-O-CH3, bothhave the same molecular formula, C2H6O.

Note: If you haven't done any organic chemistry yet, don't


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worry about the names.

They have the same number of electrons, and a similar length tothe molecule. The van der Waals attractions (both dispersionforces and dipole-dipole attractions) in each will be much thesame.

However, ethanol has a hydrogen atom attached directly to anoxygen - and that oxygen still has exactly the same two lonepairs as in a water molecule. Hydrogen bonding can occurbetween ethanol molecules, although not as effectively as inwater. The hydrogen bonding is limited by the fact that there isonly one hydrogen in each ethanol molecule with sufficient +charge.

In methoxymethane, the lone pairs on the oxygen are still there,but the hydrogens aren't sufficiently + for hydrogen bonds toform. Except in some rather unusual cases, the hydrogen atomhas to be attached directlyto the very electronegative elementfor hydrogen bonding to occur.

The boiling points of ethanol and methoxymethane show thedramatic effect that the hydrogen bonding has on the stickinessof the ethanol molecules:

ethanol (with hydrogen bonding)

78.5C

methoxymethane (without hydrogen bonding) -24.8CThe hydrogen bonding in the ethanol has lifted its boiling pointabout 100C.

It is important to realise that hydrogen bonding exists inadditionto van der Waals attractions. For example, all thefollowing molecules contain the same number of electrons, and

the first two are much the same length. The higher boiling pointof the butan-1-ol is due to the additional hydrogen bonding.


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Comparing the two alcohols (containing -OH groups), bothboiling points are high because of the additional hydrogenbonding due to the hydrogen attached directly to the oxygen -but they aren't the same.

The boiling point of the 2-methylpropan-1-ol isn't as high as thebutan-1-ol because the branching in the molecule makes thevan der Waals attractions less effective than in the longer butan-

1-ol.

Hydrogen bonding in organic molecules containingnitrogen

Hydrogen bonding also occurs in organic molecules containingN-H groups - in the same sort of way that it occurs in ammonia.Examples range from simple molecules likeCH3NH2(methylamine) to large molecules like proteins and DNA.

The two strands of the famous double helix in DNA are heldtogether by hydrogen bonds between hydrogen atoms attachedto nitrogen on one strand, and lone pairs on another nitrogen oran oxygen on the other one.

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Bonding Chemistry