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The electron pair repulsion theory

The shape of a molecule or ion is governed by the arrangement of the electron pairs around the central atom. All you need to do is to work out how many electron pairs there are at the bonding level, and then arrange them to produce the minimum amount of repulsion between them. You have to include both bonding pairs and lone pairs.

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How to work out the number of electron pairs

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

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How to work out the number of electron pairs (2)

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

  • Write down the number of electrons in the outer level of the central atom. That will be the same as the Periodic Table group number, except in the case of the noble gases which form compounds, when it will be 8.

  • Add one electron for each bond being formed. (This allows 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, deduct an electron.

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How to work out the number of electron pairs (3)

Now work out how many bonding pairs and lone pairs of electrons there are:

  • Divide by 2 to find the total number of electron pairs around the central atom.

  • Work out how many of these are bonding pairs, and how many are lone pairs. You know how many bonding pairs there are because you know how many other atoms are joined to the central atom (assuming that only single bonds are formed).

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

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How to work out the number of electron pairs (4)

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

  • Arrange these electron pairs in space to minimise repulsions. How this is done will become clear in the examples which follow

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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't enough to allow the formation of ions.

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

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

(http://www.chemguide.co.uk/atoms/bonding/shapebecl2.GIF)

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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 3 bonds, adding another 3 electrons. There is no charge, so the total is 6 electrons - in 3 pairs.

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

(http://www.chemguide.co.uk/atoms/bonding/shapebf3.GIF)

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Four electron pairs around the central atom

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

Carbon is in group 4, and so has 4 outer electrons. It is forming 4 bonds to hydrogens, adding another 4 electrons - 8 altogether, in 4 pairs. Because it is forming 4 bonds, these must all be bonding pairs.

Four electron pairs arrange themselves in space in what is called a tetrahedral arrangement. A tetrahedron is a regular triangularly-based pyramid. The carbon atom would be at the centre and the hydrogens at the four corners. All the bond angles are 109.5°.

 

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Four electron pairs around the central atom (2)

(http://www.chemguide.co.uk/atoms/bonding/shapech4.GIF)

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Other examples with four electron pairs around the

(http://www.chemguide.co.uk/atoms/bonding/shapenh3.GIF)

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REMEMBER!!!

Remember this:

                         Greatest repulsion-     Lone Pair-Lone Pair

                                                            Lone Pair-Bond Pair

                         Least Repulsion-        Bond Pair-Bond Pair

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water

Water, H2O

(http://www.chemguide.co.uk/atoms/bonding/shapeh2o.GIF)

Following the same logic as before, you will find that the oxygen has four pairs of electrons, two of which are lone pairs. These will again take up a tetrahedral arrangement. This time the bond angle closes slightly more to 104°, because of the repulsion of the two lone pairs.

The shape isn't described as tetrahedral, because we only "see" the oxygen and the hydrogens - not the lone pairs. Water is described as bent or V-shaped.

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Five electron pairs around the central atom

A simple example: phosphorus(V) fluoride, PF5

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

The 5 electron pairs take up a shape described as a trigonal bipyramid - three of the fluorines are in a plane at 120° to each other; the other two are at right angles to this plane. The trigonal bipyramid therefore has two different bond angles - 120° and 90°.

(http://www.chemguide.co.uk/atoms/bonding/shapepf5.GIF)

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What is electronegativity

Definition

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

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

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What happens if two atoms of equal electronegativi

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

(http://www.chemguide.co.uk/atoms/bonding/ab1.GIF)

If the atoms are equally electronegative, both have the same tendency to attract the bonding pair of electrons, and so it will be found on average half way between the two atoms. To get a bond 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.

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

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What happens if B is slightly more electronegative

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

(http://www.chemguide.co.uk/atoms/bonding/ab2.GIF)

That means that the B end of the bond has more than its fair share of electron density and so becomes slightly negative. At the same time, the A end (rather short of electrons) becomes slightly positive. In the diagram, "(http://www.chemguide.co.uk/atoms/bonding/delta.GIF)" (read as "delta") means "slightly" - so (http://www.chemguide.co.uk/atoms/bonding/delta.GIF)+ means "slightly positive".

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Defining polar bonds

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

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What happens if B is a lot more electronegative th

 

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

(http://www.chemguide.co.uk/atoms/bonding/ab3.GIF)

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Polar bonds and polar molecules

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

In CCl4, each bond is polar.

(http://www.chemguide.co.uk/atoms/bonding/ccl4.GIF)

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Polar bonds and polar molecules (2)

The molecule as a whole, however, isn't polar - in the sense that it doesn't have an end (or a side) which is slightly negative and one which is slightly positive. The whole of the outside of the molecule is somewhat negative, but there is no overall separation of charge from top to bottom, or from left to right.

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Polar bonds and polar molecules (3)

By contrast, CHCl3 is polar.

(http://www.chemguide.co.uk/atoms/bonding/chcl3.GIF)

The hydrogen at the top of the molecule is less electronegative than carbon and so is slightly positive. This means that the molecule 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.

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Patterns of electronegativity in the Periodic Tabl

 

The most electronegative element is fluorine. If you remember that fact, everything becomes easy, because electronegativity must always increase towards fluorine in the Periodic Table.(http://www.google.co.uk/search?hl=en&site=imghp&tbm=isch&source=hp&biw=1024&bih=673&q=electronegativity&oq=electronegativity&gs_l=img.3..0l10.1644.9736.0.9877.17.12.0.5.5.0.351.1355.9j1j1j1.12.0...0.0...1ac.1.12.img.hSwqbx8IBA8#imgrc=dJLhHvMFbaQKLM%3A%3BzbsQML0-xqanEM%3Bhttp%253A%252F%252Fwww.kentchemistry.com%252Flinks%252Fbonding%252Felectronegativity.jpg%3Bhttp%253A%252F%252Fwww.kentchemistry.com%252Flinks%252Fbonding%252FElectronegativity.htm%3B432%3B189)

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Trends in electronegativity across a period

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

(http://www.chemguide.co.uk/atoms/bonding/p3eneg.GIF)

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Trends in electronegativity down a group

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

(http://www.chemguide.co.uk/atoms/bonding/g1g7eneg.GIF)

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Explaining the patterns in electronegativity

The attraction that a bonding pair of electrons feels for a particular nucleus depends on: the number of protons in the nucleus; the distance from the nucleus; the amount of screening by inner electrons.

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Why does electronegativity increase across a perio

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

(http://www.chemguide.co.uk/atoms/bonding/nacleneg.GIF)

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 protons in it. It is no wonder the electron pair gets dragged so far towards the chlorine that ions are formed.

Electronegativity increases across a period because the number of charges on the nucleus increases. That attracts the bonding pair of electrons more strongly

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Why does electronegativity fall as you go down a g

Think of hydrogen fluoride and hydrogen chloride.

(http://www.chemguide.co.uk/atoms/bonding/fvcleneg.GIF)

The bonding pair is shielded from the fluorine's nucleus only by the 1s2 electrons. In the chlorine case it is shielded by all the 1s22s22p6 electrons.

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

As you go down a group, electronegativity decreases because the bonding pair of electrons is increasingly distant from the attraction of the nucleus.

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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 a covalent bond if one atom is more electronegative than the other. But you can also look at the formation of polar covalent bonds by imagining that you start from ions.

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

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

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The polarising ability of positive ions (2)

(http://www.chemguide.co.uk/atoms/bonding/alionpolar.GIF)

We say that the aluminium ions polarise the chloride ions.

In the case of aluminium chloride, the electron pairs are dragged back towards the aluminium to such an extent that the bonds become covalent. But because the chlorine is more electronegative than aluminium, the electron pairs won't be pulled half way between the two atoms, and so the bond formed will be polar.

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Factors affecting polarising ability

Positive ions can have the effect of polarising (electrically distorting) nearby negative ions. The polarising ability depends on the charge density in the positive ion.

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

As a negative ion gets bigger, it becomes easier to polarise. For example, 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 easily dragged away from the iodide ion. On the other hand, aluminium fluoride is ionic because the aluminium ion can't polarise the small fluoride ion sufficiently to form a covalent bond.

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What are intermolecular attractions?

Intermolecular attractions are attractions between one molecule and a neighbouring molecule. The forces of attraction which hold an individual molecule together (for example, the covalent bonds) are known as intramolecular attractions. These two words are so confusingly similar that it is safer to abandon one of them and never use it.

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van der Waals forces: dispersion forces

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

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Temporary fluctuating dipoles

Attractions are electrical in nature. In a symmetrical molecule like hydrogen, however, there doesn't seem to be any electrical distortion to produce positive or negative parts. But that's only true on average.

(http://www.chemguide.co.uk/atoms/bonding/fluctuate0.GIF)

The lozenge-shaped diagram represents a small symmetrical molecule - H2, perhaps, or Br2. The even shading shows that on average there is no electrical distortion.

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Temporary fluctuating dipoles (2)

But the electrons are mobile, and at any one instant they might find themselves towards one end of the molecule, making that end (http://www.chemguide.co.uk/atoms/bonding/delta.GIF)-. The other end will be temporarily short of electrons and so becomes (http://www.chemguide.co.uk/atoms/bonding/delta.GIF)+.

(http://www.chemguide.co.uk/atoms/bonding/fluctuate1.GIF)

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

(http://www.chemguide.co.uk/atoms/bonding/fluctuate2.GIF)

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Temporary fluctuating dipoles (3)

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

If both the helium electrons happen to be on one side of the atom at the same time, the nucleus is no longer properly covered by electrons for that instant.

(http://www.chemguide.co.uk/atoms/bonding/atompolar.GIF)

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How temporary dipoles give rise to intermolecular

(http://www.chemguide.co.uk/atoms/bonding/approach.GIF)

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

This sets up an induced dipole in the approaching molecule, which is orientated in such a way that the (http://www.chemguide.co.uk/atoms/bonding/delta.GIF)+ end of one is attracted to the (http://www.chemguide.co.uk/atoms/bonding/delta.GIF)- end of the other.

(http://www.chemguide.co.uk/atoms/bonding/induced1.GIF)

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How temporary dipoles give rise to intermolecular

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

(http://www.chemguide.co.uk/atoms/bonding/induced2.GIF)

The polarity of both molecules reverses, but you still have (http://www.chemguide.co.uk/atoms/bonding/delta.GIF)+ attracting (http://www.chemguide.co.uk/atoms/bonding/delta.GIF)-. As long as the molecules stay close to each other the polarities will continue to fluctuate in synchronisation so that the attraction is always maintained.

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How temporary dipoles give rise to intermolecular

The polarity of both molecules reverses, but you still have (http://www.chemguide.co.uk/atoms/bonding/delta.GIF)+ attracting (http://www.chemguide.co.uk/atoms/bonding/delta.GIF)-. As long as the molecules stay close to each other the polarities will continue to fluctuate in synchronisation so that the attraction is always maintained.

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How temporary dipoles give rise to intermolecular

There is no reason why this has to be restricted to two molecules. As long as the molecules are close together this synchronised movement of the electrons can occur over huge numbers of molecules.

(http://www.chemguide.co.uk/atoms/bonding/lattice.GIF)

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How molecular shape affects the strength of the di

The shapes of the molecules also matter. Long thin molecules can develop bigger temporary dipoles due to electron movement than short fat ones containing the same numbers of electrons.

Long thin molecules can also lie closer together - these attractions are at their most effective if the molecules are really close.

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How molecular shape affects the strength of the di

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

(http://www.chemguide.co.uk/atoms/bonding/butanes.GIF)

Butane has a higher boiling point because the dispersion forces are greater. The molecules are longer (and so set up bigger temporary dipoles) and can lie closer together than the shorter, fatter 2-methylpropane molecules.

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van der Waals forces: dipole-dipole interactions

A molecule like HCl has a permanent dipole because chlorine is more electronegative than hydrogen. These permanent, in-built dipoles will cause the molecules to attract each other rather more than they otherwise would if they had to rely only on dispersion forces.

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van der Waals forces: dipole-dipole interactions

It's important to realise that all molecules experience dispersion forces. Dipole-dipole interactions are not an alternative to dispersion forces - they occur in addition to them. Molecules which have permanent dipoles will therefore have boiling points rather higher than molecules which only have temporary fluctuating dipoles.

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

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van der Waals forces: dipole-dipole interactions

(http://www.chemguide.co.uk/atoms/bonding/c2h6vch3f.GIF)

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van der Waals forces: dipole-dipole interactions (

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

The higher boiling point of fluoromethane is due to the large permanent dipole on the molecule because of the high electronegativity of fluorine. However, even given the large permanent polarity of the molecule, the boiling point has only been increased by some 10°.

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van der Waals forces: dipole-dipole interactions (

Here is another example showing the dominance of the dispersion forces. Trichloromethane, CHCl3, is a highly polar molecule because of the electronegativity of the three chlorines. There will be quite strong dipole-dipole attractions between one molecule and its neighbours.

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The evidence for hydrogen bonding

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

(http://www.chemguide.co.uk/atoms/bonding/bptgp4hyd.GIF)

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The evidence for hydrogen bonding (2)

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

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

(http://www.chemguide.co.uk/atoms/bonding/bptgp567hyd.GIF)

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The evidence for hydrogen bonding (3)

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 compound of hydrogen with 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.

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The origin of hydrogen bonding

The molecules which have this extra bonding are:

(http://www.chemguide.co.uk/atoms/bonding/nh3h2ohf.GIF)

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The origin of hydrogen bonding (2)

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.

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The origin of hydrogen bonding (3)

Consider two water molecules coming close together.

(http://www.chemguide.co.uk/atoms/bonding/h2ohbonds.GIF)

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The origin of hydrogen bonding (4)

The (http://www.chemguide.co.uk/atoms/bonding/delta.GIF)+ 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.

Water as a "perfect" example of hydrogen bonding

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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.

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More complex examples of hydrogen bonding (2)

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.

(http://www.chemguide.co.uk/atoms/bonding/clhbonds.GIF)

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More complex examples of hydrogen bonding (3)

However complicated the negative ion, there will always be lone pairs that the hydrogen atoms from the water molecules can hydrogen bond to.

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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.

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Hydrogen bonding in alcohols (2)

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

(http://www.chemguide.co.uk/atoms/bonding/olvether.GIF)

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Trends in Atomic Radius

(http://www.chemguide.co.uk/inorganic/group2/atradiuschart.GIF)

You can see that the atomic radius increases as you go down the Group. Notice that beryllium has a particularly small atom compared with the rest of the Group.

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Explaining the increase in atomic radius

The radius of an atom is governed by

  • the number of layers of electrons around the nucleus

  • the pull the outer electrons feel from the nucleus.

Compare beryllium and magnesium:

Be (http://www.chemguide.co.uk/inorganic/group2/padding.GIF) 1s22s2 Mg (http://www.chemguide.co.uk/inorganic/group2/padding.GIF) 1s22s22p63s2

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Explaining the increase in atomic radius (2)

In each case, the two outer electrons feel a net pull of 2+ from the nucleus. The positive charge on the nucleus is cut down by the negativeness of the inner electrons.

(http://www.chemguide.co.uk/inorganic/group2/bemgscreening.gif)

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Explaining the increase in atomic radius (3)

This is equally true for all the other atoms in Group 2. Work it out for calcium if you aren't convinced.

The only factor which is going to affect the size of the atom is therefore the number of layers of inner electrons which have to be fitted in around the atom. Obviously, the more layers of electrons you have, the more space they will take up - electrons repel each other. That means that the atoms are bound to get bigger as you go down the Group.

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Trends in First Ionisation Energy

First ionisation energy is the energy needed to remove the most loosely held electron from each of one mole of gaseous atoms to make one mole of singly charged gaseous ions - in other words, for 1 mole of this process:

(http://www.chemguide.co.uk/inorganic/group2/padding.GIF)(http://www.chemguide.co.uk/inorganic/group2/ieeqtn.gif)

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Trends in First Ionisation Energy (2)

(http://www.chemguide.co.uk/inorganic/group2/ieschart.gif)

Notice that first ionisation energy falls as you go down the group.

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Explaining the decrease in first ionisation energy

Ionisation energy is governed by

  • the charge on the nucleus,

  • the amount of screening by the inner electrons,

  • the distance between the outer electrons and the nucleus.

As you go down the Group, the increase in nuclear charge is exactly offset by the increase in the number of inner electrons. Just as when we were talking about atomic radius further up this page, in each of the elements in this Group, the outer electrons feel a net attraction of 2+ from the centre.

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Explaining the decrease in first ionisation energy

However, as you go down the Group, the distance between the nucleus and the outer electrons increases and so they become easier to remove - the ionisation energy falls.

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Trends in Electronegativity

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. It is usually measured on the Pauling scale, on which the most electronegative element (fluorine) is given an electronegativity of 4.0.

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Trends in Electronegativity (2)

(http://www.chemguide.co.uk/inorganic/group2/enegchart.gif)

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Trends in Electronegativity (3)

All of these elements have a low electronegativity. (Remember that the most electronegative element, fluorine, has an electronegativity of 4.0.) Notice that electronegativity falls as you go down the Group. The atoms become less and less good at attracting bonding pairs of electrons.

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Explaining the decrease in electronegativity

Imagine a bond between a magnesium atom and a chlorine atom. Think of it to start with as a covalent bond - a pair of shared electrons. The electron pair will be dragged towards the chlorine end because there is a much greater net pull from the chlorine nucleus than from the magnesium one.

(http://www.chemguide.co.uk/inorganic/group2/mg-clbond.gif)

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Explaining the decrease in electronegativity (2)

The electron pair ends up so close to the chlorine that there is essentially a transfer of an electron to the chlorine - ions are formed.

The large pull from the chlorine nucleus is why chlorine is much more electronegative than magnesium is.

Now compare this with the beryllium-chlorine bond.

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Explaining the decrease in electronegativity (3)

The net pull from each end of the bond is the same as before, but you have to remember that the beryllium atom is smaller than a magnesium atom. That means that the electron pair is going to be closer to the net 2+ charge from the beryllium end, and so more strongly attracted to it.

(http://www.chemguide.co.uk/inorganic/group2/be-clbond.gif)

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Explaining the decrease in electronegativity (4)

In this case, the electron pair doesn't get attracted close enough to the chlorine for an ionic bond to be formed. Because of its small size, beryllium forms covalent bonds, not ionic ones. The attraction between the beryllium nucleus and a bonding pair is always too great for ions to be formed.

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Summarising the trend down the Group

As the metal atoms get bigger, any bonding pair gets further and further away from the metal nucleus, and so is less strongly attracted towards it. In other words, as you go down the Group, the elements become less electronegative.

As you go down the Group, the bonds formed between these elements and other things such as chlorine become more and more ionic. The bonding pair is increasingly attracted away from the Group 2 element towards the chlorine (or whatever).

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Trends in Melting Point, Boiling Point, and Atomis

Melting points

(http://www.chemguide.co.uk/inorganic/group2/mptchart.gif)You will see that (apart from where the smooth trend is broken by magnesium) the melting point falls as you go down the Group.

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Trends in Melting Point, Boiling Point, and Atomis

Boiling points

(http://www.chemguide.co.uk/inorganic/group2/bptchart.gif)

You will see that there is no obvious pattern in boiling points. It would be quite wrong to suggest that there is any trend here whatsoever.

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Atomisation energy

This is the energy needed to produce 1 mole of separated atoms in the gas state starting from the element in its standard state (the state you would expect it to be in at approximately room temperature and pressure).

(http://www.chemguide.co.uk/inorganic/group2/atenergychart.gif)

And again there is no simple pattern. It looks similar to, but not exactly the same as, the boiling point chart.

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LOL HAVE FUN!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!

http://www.chemguide.co.uk/index.html#top

go to this website for information on the few topics missing.

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