ATOMIC AND PHYSICAL PROPERTIES OF THE 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 melting point.

Even if you aren't currently interested in all these things, it would probably pay you to read the whole page. The same ideas tend to recur throughout the atomic properties, and you may find that earlier explanations help to you understand later ones.

Trends in Atomic Radius

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.

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:

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.

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:

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

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.

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.

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.

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.

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.

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.

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).

Trends in Melting Point

You will see that (with the exception of magnesium) the melting point falls as you go down the Group.

Explaining the trends in melting point (or not!)

Thankfully, none of the current A'level syllabuses ask you to explain the pattern in the melting points. Trying to explain patterns in melting point is notoriously difficult, and simple explanations usually come to grief if you look at them closely.

It would be tempting to say that (apart from magnesium) the lower melting points as you go down the Group reflect weaker metallic bonds, but this is dangerous. The metallic bond isn't fully broken on melting. Boiling point is usually a better guide to the strengths of the forces involved.

You will sometimes come across an explanation for the low melting point of magnesium in terms of the way the atoms pack in the crystal. It isn't very difficult to pick holes in this argument and, in any case, the boiling point of magnesium is also surprisingly low - and packing is completely irrelevant when you are thinking about a liquid. In the case of magnesium, there must be some factor producing a relatively weak metallic bond. (I have no idea what that might be!)

You will see that there is no obvious pattern in the boiling points - implying that there is no simple pattern in the strengths of the metallic bonds.

Another measure of the strengths of the metallic bonds might be 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. Don't worry about any of this for A'level purposes, but be very wary of any explanation which looks deceptively simple!