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Science

Covalent bonding

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Dot-and-cross diagrams – compounds

You will also need to be able to draw dot-and-cross diagrams representing the covalent bonds [covalent bond: A covalent bond between atoms forms when atoms share electrons to achieve a full outer shell of electrons. ] in the molecules [molecule: A molecule is a collection of two or more atoms held together by chemical bonds. It is the smallest part of a substance that displays the properties of the substance. ] of some common compounds [compound: A compound is a substance formed by the chemical union (involving bond formation) of two or more elements. ]:

Hydrogen chloride, HCl

Bonding in hydrogen chloride. A hydrogen atom and chlorine atom each share one electron

Hydrogen atoms and chlorine atoms can each form one covalent bond. One pair of electrons [electron: An electron is a very small negatively-charged particle found in an atom in the space surrounding the nucleus. ] is shared in a hydrogen chloride molecule (HCl).

Water, H2O

Bonding in water. Two hydrogen atoms each share one electron, and an oxygen atom shares two electrons

Hydrogen atoms can each form one covalent bond, while oxygen atoms can each form two covalent bonds. Two pairs of electrons are shared in a water molecule (H2O).

Ammonia, NH3

Bonding in ammonia. Three hydrogen atoms each share one electron, and a nitrogen atom shares three electrons

Hydrogen atoms can each form one covalent bond, while nitrogen atoms can each form three covalent bonds. Three pairs of electrons are shared in an ammonia molecule (NH3).

Methane, CH4

Bonding in methane. Four hydrogen atoms each share one electron, and a carbon atom shares four electrons

Hydrogen atoms can each form one covalent bond, while carbon atoms can each form four covalent bonds. Four pairs of electrons are shared in a methane molecule (CH4).

Now try a Test Bite .

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Co-ordinate 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 (such as GCSE), 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 aluminium chloride

Aluminium 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 aluminium and boron are in the same group of the Periodic Table, as are fluorine and chlorine.


Measurements of the relative formula mass of aluminium 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.

 

Note:  The uninteresting electrons on the chlorines have been 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 separate AlCl3 molecules.

 

Note:  Aluminium chloride is complicated because of the way it keeps changing its bonding as the temperature increases. If you are interested in exploring this in more detail, you could have a look at the page about the Period 3 chlorides . It isn’t particularly relevant to the present page, though.

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

 

 

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.


Aluminium 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 aluminium re-organises (hybridises) 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 aluminium 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 aluminium 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.

 

Note:  Dotted arrows represent lone pairs coming from water molecules behind the plane of the screen or paper. Wedge shaped arrows represent bonds from water molecules in front of the plane of the screen or paper.

 

 

Two more molecules

 

Note:  Only one current A’level syllabus wants these two. Check yours! If you haven’t got a copy of your syllabus , follow this link to find out how to get one.

 

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 delocalisation.

INTERMOLECULAR BONDING – HYDROGEN BONDS

 

This page explains the origin of hydrogen bonding – a relatively strong form of intermolecular attraction. If you are also interested in the weaker intermolecular forces (van der Waals dispersion forces and dipole-dipole interactions), there is a link at the bottom of the page.

 

The evidence for hydrogen bonding

Many elements form compounds with hydrogen – referred to as "hydrides". If you plot the boiling points of the hydrides of the Group 4 elements, you find that the boiling points increase as you go down the group.

The increase in boiling point happens because the molecules are getting larger with more electrons, and so van der Waals dispersion forces become greater.

 

Note:  If you aren’t sure about van der Waals dispersion forces , it would pay you to follow this link before you go on.

 

If you repeat this exercise with the hydrides of elements in Groups 5, 6 and 7, something odd happens.

Although for the most part the trend is exactly the same as in group 4 (for exactly the same reasons), the boiling point of the hydride of the first element in each group is abnormally high.

In the cases of NH3, H2O and HF there must be some additional intermolecular forces of attraction, requiring significantly more heat energy to break. These relatively powerful intermolecular forces 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 screen or paper away from you, and wedge-shaped ones are coming out towards you.

 

Notice that in each of these molecules:

  • The hydrogen is attached directly to one of the most electronegative elements, causing the hydrogen to acquire a significant amount of positive charge.
  • Each of the elements to which the hydrogen is attached is not only significantly negative, but also has at least one "active" lone pair.

Lone pairs at the 2-level have the electrons contained in a relatively small volume of space which therefore has a high density of negative charge. Lone pairs at higher levels are more diffuse and not so attractive to positive things.

 

Note:  If you aren’t happy about electronegativity , you should follow 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 is almost as if you were beginning to form a co-ordinate (dative covalent) bond. It doesn’t go that far, but the attraction is significantly stronger than an ordinary dipole-dipole interaction.

Hydrogen bonds have about a tenth of the strength of an average covalent bond, and are being constantly broken and reformed in liquid water. If you liken the covalent bond between the oxygen and hydrogen to a stable marriage, the hydrogen bond has "just good friends" status. On the same scale, van der Waals attractions represent mere passing acquaintances!

Water as a "perfect" example of hydrogen bonding

Notice that each water molecule can potentially form four hydrogen bonds with surrounding water molecules. There are exactly the right numbers of + hydrogens and lone pairs so that every 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, the amount of hydrogen bonding is limited by the fact that each nitrogen only has one lone pair. In a group of ammonia molecules, there aren’t enough lone pairs to go around to satisfy all the hydrogens.

In hydrogen fluoride, the problem is a shortage of hydrogens. In water, there are exactly the right number of each. Water could be 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 on molecular structures .

 

 

More complex examples of hydrogen bonding

The hydration of negative ions

When an ionic substance dissolves in water, water molecules cluster around the separated ions. This process is called hydration.

Water frequently attaches to positive ions by co-ordinate (dative covalent) bonds. It bonds to negative ions using hydrogen bonds.

 

Note:  If you are interested in the bonding in hydrated positive ions, you could follow this link to co-ordinate (dative covalent) bonding .

 

The diagram shows the potential hydrogen bonds formed to a chloride ion, Cl-. Although the lone pairs in the chloride ion are at the 3-level and wouldn’t normally be active enough to form hydrogen bonds, in this case they are made more attractive by 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 can hydrogen 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 an oxygen or a nitrogen is capable of hydrogen bonding. Such molecules will always have higher boiling points than similarly sized molecules which don’t have an -O-H or an -N-H group. The hydrogen bonding makes the molecules "stickier", and more heat is necessary to separate them.

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

 

Note:  If you haven’t done any organic chemistry yet, don’t worry about the names.

 

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

However, ethanol has a hydrogen atom attached directly to an oxygen – and that oxygen still has exactly the same two lone pairs as in a water molecule. Hydrogen bonding can occur between ethanol molecules, although not as effectively as in water. The hydrogen bonding is limited by the fact that there is only 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 to form. Except in some rather unusual cases, the hydrogen atom has to be attached directly to the very electronegative element for hydrogen bonding to occur.

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

ethanol (with hydrogen bonding)

78.5°C

methoxymethane (without hydrogen bonding)

-24.8°C

The hydrogen bonding in the ethanol has lifted its boiling point about 100°C.

 

It is important to realise that hydrogen bonding exists in addition to van der Waals attractions. For example, all the following molecules contain the same number of electrons, and the first two are much the same length. The higher boiling point of the butan-1-ol is due to the additional hydrogen bonding.

Comparing the two alcohols (containing -OH groups), both boiling points are high because of the additional hydrogen bonding 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 the butan-1-ol because the branching in the molecule makes the van der Waals attractions less effective than in the longer butan-1-ol.

 

Hydrogen bonding in organic molecules containing nitrogen

Hydrogen bonding also occurs in organic molecules containing N-H groups – in the same sort of way that it occurs in ammonia. Examples range from simple molecules like CH3NH2 (methylamine) to large molecules like proteins and DNA. The two strands of the famous alpha-helix in DNA are held together by hydrogen bonds involving N-H groups.