Periodicity three :Trend in the physical properties of the alkaline earth metals.

GCSE Science

Group II – The Alkaline Earth Metals. (Variation in Physical Properties)

Beryllium (Be) (Atomic number = 4)

1s2 2s2

Magnesium (Mg) (Atomic number =12)

1s2 2s2 2p6 3s2

Calcium (Ca) (atomic number = 20)

1s2 2s2 2p6 3s2 3p6 4s2

Strontium (Sr) (Atomic number = 38)

1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s2

Barium (Ba) (Atomic number = 56)

1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p6 6s2

Notice that each alkaline earth metal has two electrons in its outermost shell and that these electrons occupy s orbitals.

1) Trends in the physical properties of the alkaline earth metals.

The trend in atomic radius

The graph below shows the trend in atomic radius for the alkaline earth metals.

Notice that as the atomic number increases the size of the atoms also increase.

In other words as group two is descended, atomic radius increases

Why is this? Look closely at the electronic structures of each element provided for you at the start of this section.

You should see that as you go down the group the number of electron shells increases by one each time (for example the outer shell for beryllium is the second shell and for magnesium it is the third etc)

This means that progressively the outer electrons are further from the nucleus and so since atomic radius is the distance from the centre of the atom to the perimeter of the outermost electron shell, the atoms become progressively bigger as the group is descended.

2) First Ionisation energy

A graph showing the first ionisation energies of the group II elements is shown below.

Notice that first ionisation energies decrease as the group is descended (i.e. it becomes progressively easier to remove an electron from the atoms of the elements).

Why should it be easier to remove an electron from the outer shell of calcium than it is to remove one from the outer shell of magnesium, for example?

The outer electrons in the magnesium atom are in the third electron shell while the outer electrons in calcium are in the fourth electron shell. This means that calcium’s outer electrons are more distant from the nucleus and are more shielded than magnesium’s.

Therefore, calcium’s outer electrons are less under the control of the nucleus and so can be more easily lost.

The same argument can be applied all the way down the group.

(If you are unclear about the explanation of this you may need to look back at your earlier work on electronic configuration)

3) Electronegativity

A graph showing the trend in electronegativity down group II is shown below :

Notice that electronegativity decreases as the group is descended (i.e.as the group is decended the elements attract electrons less strongly). Why should this be?

We should remember that electronegativity gives an idea of the relative ability of an atom to attract the electrons ...

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The same argument can be applied all the way down the group.

(If you are unclear about the explanation of this you may need to look back at your earlier work on electronic configuration)

3) Electronegativity

A graph showing the trend in electronegativity down group II is shown below :

Notice that electronegativity decreases as the group is descended (i.e.as the group is decended the elements attract electrons less strongly). Why should this be?

We should remember that electronegativity gives an idea of the relative ability of an atom to attract the electrons in a covalent bond. If any atom forms a covalent bond it is reasonable to expect that the electrons in the bond are shared in the outer most electron shell of that atom.

So electronegativity is about the ability of the nucleus in the atom to attract electrons from the outermost shell. Why should this ability decrease down the group?

As already stated, the distance from the outer shell to the nucleus gradually increases down the group. This means that as the group is descended the outer electrons are more distant from the nucleus, there will be more shielded from the effects of the nucleus by inner shells and so progressively the outer electrons become less attracted to the nucleus.

Therefore, as the group is descended the relative ability of the atoms to attract electrons decreases.

The same argument can be applied all the way down the group.

Questions

What is the trend in atomic radius down group two? How can this be explained by consideration of the electronic structures of the atoms of the elements involved?

2) The following questions refers to the graph of change in first ionisation energy down group two.

Do you expect that, in general, these values are high or low in comparison to the first ionisation energies of other groups in the periodic table? Explain your answer.
Why does first ionisation energy decrease as the group is descended?

(a) What is meant by the term electronegativity ?

(b) What is the trend in electronegativity down group two? Explain why this is.
(c) The group one metals all have relatively low electronegativities in comparison to other elements in the periodic table. Explain why this is.

Melting point

The melting point of an element is determined by the forces that hold the atoms together.

If forces are strong then considerable energy must be supplied to break them and cause the atoms (at least in part) to move apart from one another (i.e. melt).

Therefore elements with strong forces between atoms will have high melting points

Alternatively, if the forces that exist between atoms are weak then little energy need be provided to cause the atoms to move separately from one another.

Elements with weak structure have low melting points.

In general, we can say that all of the group two elements must have pretty strong forces holding them together, We can assume this because all of the alkaline earth metals are solid at room temperature.

However, how do we explain the trend down the group?

Well first we must remember that all of the elements in group II belong to a band of substances that make up well over half of all the elements in the periodic table – they are all metals!!!

In order to explain the trend in melting point down goup II, we must investigate the type of bonding that holds metals together.

In the next section of this work we consider metallic bonding.

Metallic bonding

Metallic bonding bestows certain properties on metal that we are very familiar with. Man has been exploiting the properties of metals for thousands of years!

The main properties of metals are :

Metals have high melting points – All metals (with just one exception – mercury) are solids at room temperature. This suggests that metallic bonding is very strong.

All metals conduct electricity when solid.

Metals are malleable and ductile – this means that although metal atoms are strongly bonded to one another, the metals structure can be stretched and bent into different shapes.

Note : Metals are malleable and ductile but they still have a strong structure (we shape metals to form bridges, ships etc but rely on the strength of the structure thereafter).

High Melting Points

What do all metals have in common?

Electronic structure.

In terms of electronic structure we can say that all metal elements are distinctive in that they have relatively few electrons in their outermost shells.

We can see this by considering the alkaline earth metals, each member of the group has only two electrons in their outermost shells.

In other words, in metals the outermost shell are only just beginning to fill and so their nuclei are relatively weak (compared with elements further along the same period).

We can use this to explain the properties of metals and the particular properties of the alkaline earth metals (group II).

High Melting Points

As all metal atoms have relatively weak nuclei and the outer electrons are relatively few (and usually pretty well shielded by inner electron shells), the outer electrons of metal can be easily lost.

In fact it is most common that in the structure of a metal the “outer electrons” of any one atom are not even orbiting that atom!! In fact they are thought to be lost into a sea of similar electrons that flow around the structure (They are said to have become delocalised from their atoms. This “sea of delocalised electrons” acts like a very strong glue holding the structure of the metal together. This helps to explain why we have to supply a lot of heat energy to break the metal structure apart.

The atoms themselves will have lost electrons and so become positively charged and are referred to a positive centres.

The reason that these electrons can be lost is because the structure of metals as positive centres in a sea of electrons is more stable than single atoms. The stability gain is sufficient to overcome the ionisation energy required to delocalise the electrons in the outermost shell. (See diagram below)

This still leaves us with the question, why do all the group II metals have high melting points?

Well we should expect the group II metals to have rather strong metallic forces since each atom is capable of losing two electrons into the sea of electrons meaning that the sea of electrons will be very strong.

Notice also that the positive centres in the group II metal structures would normally carry a 2+ charge.

See diagram below

OK! But how do we explain the trend down the group? As we can see from the graph (reproduced again below) the general trend is that the metallic bond strength decreases down the group. How do we explain this when we know that each of the atoms in any group II element should release two electrons into the sea of electrons –surely the melting points should all be about the same?

Actually no, because we need to consider atomic size (or more correctly the size of the positive centres left behind when the outer electrons delocalise). The sea of electrons need to cover the positive centres like wrapping paper needs to cover Christmas presents. All that we have said above is that each of the presents provides the same amount of wrapping paper. But in the case of strontium (for example) the positive centres are large (big present) and the wrapping paper will barely be enough to cover the area required, while for beryllium the positive centres are very small (little present) and there is more than enough wrapping paper to go around!!

So in the case of barium the sea of electrons is thinly spread out and so is less effective, this makes it easier to disrupt the structure and the melting point is lower.

See diagram provided below

Only one last problem remains, as you can see from our chart the melting point of magnesium is uncharacteristically low compared to the other group II elements (ie it does not appear to follow the trend). Why is this?

Well the positive centres in the elements below magnesium are packed in a different way compared to the atoms in magnesium (ie it has a different crystal structure). It is easier to break apart the crystal structure of magnesium and so its overall melting point is lowered.

Conductance of Electricity.

Metals are Malleable.

Questions

Metallic structure can be regarded as “positive centres “ surrounded by a sea of electrons.

Explain, with the aid of a diagram, what this means.
Use this model to explain the following phenomena

Metal, as elements, are generally regarded as having high melting points.
Metals can conduct electricity when solid
The alkaline earth metals, in comparison to other metals, have rather low melting points
Melting point decreases as group two is descended