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.

Trends in Atomic Radius

Elements included in this group include the beryllium, magnesium, calcium, strontium, barium and radium. Among all the elements, radium is the radioactive element. Abundant amounts of oxides of these elements are found in the earth’s crust. So, group IIA elements are also termed as alkaline earth metals.

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:

Be 1s22s2
Mg 1s22s22p63s2

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, Boiling Point, and Atomisation Energy

The facts

Melting points

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

Boiling points

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.

Physical properties :

    • Physical nature:
These elements have two electrons in their outermost orbital. They have a silvery luster. But, it soon disappears upon exposure to air. They are malleable and ductile but very less when compared to alkali metals.
    • Atomic Volume and Radius:
As we go down the group, the number of electrons is less when compared to the increase in the number of available orbitals. Hence, the atomic radius increases gradually as we move down the group. It should be noted that these elements are smaller when compared to alkali metals. This is because of the presence of two electrons in the outermost shell. Hence, the effect of the nucleus on the outermost shell is comparatively more than the alkali metals which have a single electron.
    • Density:
Due to presence of two electrons in the outermost shell they can be more densely packed compared to alkali metals.
    • Melting and Boiling Points:
These elements have a higher boiling and melting points. Due to presence of two electrons in the valence shell they remain tightly packed in solid state.
    • Ionization Energy:
Alkaline earth metals have a smaller size and higher nuclear charge as a result of which the valence electrons are held strongly. Hence, more amount of energy is required to remove an electron from the valence electron resulting in high ionization energies.
    • Oxidation State:
Due to the presence of two electrons in the valence shell, elements of group IIA exhibit +2 oxidation i.e. they are bivalent.
    • Electropositivity:
As the size to charge ratio is very high, they are highly electropositive in nature.
    • Electronegativity:
As they are highly electropositive, they are less electronegative. Electronegativity decreases as we move down the group. Beryllium is highly electronegative due to small size.
    • Conductivity:
They are good conductors of heat and electricity. This is due to the presence of a two electrons that can very easily move within the crystal lattice of the elements.
    • Flame colorization:
When the elements are heated, electrons present in the valence shell are excited to higher energy levels. When the excited electrons return back after losing the energy they emit certain amount light. All Alkaline earth metals except beryllium and magnesium emit various colors depending on the degree of excitation. These two elements due to their smaller size have high ionization energies and high excitation energy. Hence, they are not excited to higher energy states and thus no flame colorization. .
    • Reducing property:
The two electrons in the valence shell can be very easily removed. Hence, all alkaline earth metals act as strong reducing agents. The reducing nature increases as we move down the group. However, the reducing nature of alkaline earth metals is less than their fellow s-block elements, the alkali metals.


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