3.1.1 (g) Melting Points across Periods 2 and 3

Syllabus

(g) explanation of the variation in melting points across Periods 2 and 3 in terms of structure and bonding (see also 2.2.2 o).

{Trend in structure from giant metallic to giant covalent to simple molecular lattice.}

What does this mean?

As you can see there's a fairly common pattern in both Periods.

The melting points rise from Grp 1 to Grp 4.

Then there is a huge drop off.

Why?

Increase from Group 1 to 3

Boron is a weird element and forms a giant covalent structure.

But we'll ignore that for now.

If we assume that all Grp 1 to Grp 3 elements have predictable metallic bonding then you should be able to account for this increase.

Remember that the atoms get smaller as we go across a Period (same shielding, increasing nuclear charge pulling outer shell inwards)

Also, Group 1 elements like Li/Na will form 1+ ions.

Group 2 elements like Be/Mg will form 2+ ions.

Group 3 elements like Al will form 3+ ions.

So, moving from Group 1 to Group 3 sees ions becoming smaller and more charged.

In other words, the ions have a higher charge-density as we move across the period.

And the metallic lattice will contain more electrons.

So the attractions are getting stronger and the melting point should become higher.

Increase from Group 3 to 4

At least in Period 3 this is accounted for by moving from a metallic structure to a giant covalent lattice.

These are particularly difficult to melt as they involve breaking a great many very strong covalent bonds.

You don't really need to know that Boron is also Giant Covalent.

But it helps account for Boron's melting point being much higher than Beryllium's when Aluminium's and Magnesium's does not have such a difference.

Decrease from Group 4 to Group 5

This is entirely due to changing from Giant covalent structures to simple covalent molecules.

Melting Giant structures is difficult due to all the strong covalent bonds that have to be overcome before atoms can begin to move apart.

This isn't necessary when melting simple molecular substances.

The covalent bonds are just as strong.

But the intermolecular forces are all that we have to overcome.

And these are perhaps 1% as strong.

So very little energy is needed.

What's going on from P to Ar?

Phosphorus, Sulphur, Chlorine are all elements.

So their molecules must be non-polar - no difference in electronegativity.

The only intermolecular forces that hold these elements together as solids must be van der Waals.

And the strength of van der Waals depends on numbers of electrons, so increases with molecular size and mass.

You need to learn that Phosphorus makes P₄ molecules.

These are relatively small so the van der Waals are not very strong and the melting point is low.

Certainly, P₄ molecules are much smaller than the S₈ molecules that Sulphur forms.

And so it can be no surprise that Sulphur has stronger van der Waals and a higher melting point.

Chlorine is smaller than either.

Its van der Waals are weaker.

Its melting point is lower.

But not as low as Argon's.

Argon is a Noble gas, it has a full shell.

It doesn't form molecules at all.

It exists as individual atoms so its van der Waals are very weak.

In Period 2 there is less difference in melting point from Group 5 to Group 7 than in Period 3.

But there's not so much difference between N₂, O₂ and F₂ molecules.

Even the small drop off from diatomic F₂ molecules to individual Ne atoms isn't huge.

Examiners rarely ask about Period 2.

There's more to ask about in Period 3.

Just don't rely on them never asking!

Exam-style Questions

1. (a) Explain why electrical conductivity decreases across Period 3 from Sodium to Phosphorus.

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(3)

(b) The table below shows the melting temperatures, Tm, of the Period 3 elements.

Explain the following in terms of structure and bonding.

(i) Magnesium has a higher melting temperature than Sodium.

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(2)

(ii) Silicon has a very high melting temperature.

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(2)

(iii) Sulphur has a lower melting temperature than Magnesium.

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(2)

(iv) Argon has a lower melting temperature than Chlorine.

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(3)

(Total 12 marks)

2. Why are the elements sodium to argon placed in Period 3 of the Periodic Table? Describe and explain the trends in electronegativity and atomic radius across Period 3 from Sodium to Sulphur. (7)

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Answers

1. (a) Na to Al are metals (1)

delocalised electrons (1)

Si and P are non metals (with covalent bonds)(1)

(b) (i) stronger metallic bonds in Mg (1)

smaller atoms/ions (1) (or more delocalised e’s)

(ii) macromolecular (1)

covalent (1)

OR Giant Covalent Lattice (2)

(iii) Sulphur is molecular (1)

weak forces between molecules (1)

(iv) argon exists as free atoms (1)

smaller than Cl2 molecules (1)

weaker van der Waals’ forces (1)

2

outer electrons are in third shell (1) (allow 3s, 3p etc NOT 3d)
electronegativity is the power to attract electrons in a covalent bond (or shared pair) (1)
electronegativity increases from Na to S (1)
because number of protons in the nucleus (or nuclear charge) increases (1)
and the electrons are in the same shell (or experience the same shielding) (1)
atomic radius decreases from Na to S (1)
because number of protons in the nucleus (or nuclear charge) increases (1)
and the electrons are in the same shell (or experience the same shielding) (1)

(for 'the same reasons as electronegativity increases scores one only) max 7

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