Atomic and Physical Properties of Group 2 Elements

This page explores the trends in some atomic and physical properties of the Group 2 elements – beryllium, magnesium, calcium, strontium and barium. You will find separate sections below covering the trends in atomic radius, first ionisation energy, electronegativity and physical properties.

Even if you aren't currently interested in all these things, it would probably pay you to read most of this page. The same ideas tend to recur throughout the atomic properties, and you may find that earlier explanations help to you understand later ones. The physical properties are extremely difficult to explain, however, and you might not want to read about those unless you have to.

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.

The radius of an atom is governed by

Compare beryllium and magnesium:

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.

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.

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:

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

Ionisation energy is governed by

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.

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.

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.

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.

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.

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.

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.

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.

You will see that (apart from where the smooth trend is broken by magnesium) the melting point falls as you go down the Group.

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.

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).

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

The only explanations you are likely to ever come across relate to the melting points. I will give you the most common explanation, and then explain why I think it is completely wrong!

All of these elements are held together by metallic bonds. The melting points get lower as you go down the Group because the metallic bonds get weaker. The oddity of magnesium has to be explained separately.

The atoms in a metal are held together by the attraction of the nuclei to the delocalised electrons. As the atoms get bigger, the nuclei get further away from these delocalised electrons, and so the attractions fall. That means that the atoms are more easily separated to make a liquid and finally a gas.

As you go down the Group, the arrangement of the atoms in the various solid metals changes. Beryllium and magnesium are both hexagonal close-packed; calcium and strontium are face-centred cubic; barium is body-centred cubic. Don't worry if you don't know what this means. All that matters is that there is a change in crystal structure between magnesium and calcium. That is supposed to account for the fact that magnesium is out of line with the rest of the Group.

Let's take this first, because that argument is relatively easy to demolish.

Despite the fact that the first four elements have two different structures, those structures are both 12-co-ordinated. Each atom is touched by 12 surrounding atoms. In that case, you would expect the metallic bond to be similar in each case, because the orbitals are going to overlap and delocalise in the same sort of way. Any differences just due to the structures should only be minor.

By contrast, barium is 8-co-ordinated (like the Group 1 metals). That's a less efficient packing, and you might expect that to be reflected in a much weaker metallic bond. Although the barium melting point is lower than that of strontium, it isn't dramatically lower. It just follows the general trend – suggesting that the major change of structure isn't making much difference. You can't have it both ways! If a minor change of structure at magnesium-calcium makes a huge difference, then a major one at barium should make an even bigger difference. It obviously doesn't.

The strength of the metallic bonds