Periodicity (2) Melting and boiling points of the elements of Period 3

Here we go again with the concept of periodicity.

We’ve already said what periodicity is:

Periodicity is the regular recurrence of similar properties of the elements across the periodic table.

I think one of the ways we can get this circularity or regular recurrence of similar properties is to view the periodic table in a 3D circular form as here in this diagram:

As you can see from both views the first three rows are connected; there is no break in the flow of atomic numbers.  

Now as the elements with similar properties appear above each other we can see the regular recurrence of their properties. 

So this graph I made of the melting and boiling points ought to be circular–yes!!

What we see first about these values is that there is a pattern but unlike ionisation energy and electronegativity the trends in melting and boiling point across a period are not straightforward.

Let’s define boiling and melting first and then get to an explanation of what is going on here.

Then you can construct the chart for the Period 2 values and see how they compare with Period 3. 

What is meant by boiling point?

Boiling point is the temperature at which the vapour pressure of the element is equal to the current atmospheric pressure.  The energy being supplied to the element is used to break interatomic or intermolecular bonds in the liquid element to allow its particles to exist much further apart and move at great velocity in the gaseous state.  The stronger these forces of attraction are the higher the boiling point. 

What is meant by melting point?

Melting point is the temperature at which the particles of the element transition from the solid into the liquid state. To transition, energy is supplied to break bonds between particles in the solid state instead of raising the temperature of the element. The particles remain about as close as before but they now move randomly rather than vibrating on the spot. 

What these two definitions tell you is that the explanation for the variation in melting and boiling point is going to be to do with the bonding and structure of these elements. 

Explanation

Let’s now look at and try to explain the periodic pattern.

Sodium, Magnesium, Aluminium

The first thing we can see is that there are three metallic elements: sodium, magnesium and aluminium, with relatively low melting and boiling points.

Metallic bonds hold the atoms together in these three metals. 

Here is a pictorial description of metallic bonding for a group 1 and a group 2 metal.

So notice this as we go from sodium to magnesium to aluminium the charge on the metal ion increases from +1 to +3 that also means that the number of delocalised electrons per atom increases and that has the effect of increasing the strength of the metallic bonds. 

Silicon

So what happens with silicon?

Well, silicon is not a metal it is a metalloid i.e. it has some metallic characteristics (it looks like a metal grey and shiny) but crucially its structure and bonding is not metallic. 

The atoms of silicon are held together by strong directional covalent bonds in a huge atomic network. 

This is much like the structure of diamond that you might already be familiar with.

So each silicon atoms as we can see is covalently bonded to four other atoms in a tetrahedral arrangement. 

Each of these four covalent bonds is strong.

This tetrahedral arrangement extends to edge of a silicon crystal and is sometimes called a giant structure for obvious reasons. 

To merely melt this stuff, energy will need to be given to it to break each covalent bond if the atoms are to be set free to move randomly around each other. 

As you can see that energy value and hence the temperature at which silicon melts will be very high. 

Phosphorus, Sulfur, Chlorine and Argon

Again these are not metal but non–metals. 

They have molecular structures. 

They exist in small groups of atoms: P₄, S₈, Cl₂ and Ar.

There are strong forces of attraction between their atoms within each molecule; what are called intra–molecular bonds. 

But what matters for our explanation are the inter–molecular bonds.

These inter molecular bonds are very weak.

These are called van der Waals forces.

The strength of these van der Waals forces depends n the number of electrons in each molecule and so the melting and boiling points follow this pattern: S₈ > P₄ > Cl₂ > Ar.

The bigger the molecule, more electrons, the stronger the van der Waals force, the higher the melting and boiling point.

Periodicity (1) Ionisation energy and electronegativity of the elements

The Periodic Table is a wonderful thing.

It’s not just that there are these elements in groups that reflect their physical properties like you look at the Alkali Metals (Group 1) and see that they get more reactive with water as you go down the group but the properties are periodic.

The periodic nature of the physical properties is called periodicity.

Here’s a definition of periodicity:  Periodicity is the name given to the regular recurrence of similar physical and chemical properties of the elements and their compounds in the periodic table. 

It’s like this: your school timetable is periodic because the same chemistry lesson with the same teacher occurs at the same time each week. 

Your timetable is a periodic table of sorts!! (See, you just can’t get away from Chemistry its everywhere!!)

In life periodic things take place year on year: we go on holiday in the summer, we give presents to those we love at Christmas, we hope that people remember our birthday when it comes round every year. 

So how do the properties of the elements reveal their periodicity?

Let’s look at some of the obvious properties of the elements and see if we can see periodic patterns in them.

Now if you have a Data Book of elements’ properties you can draw up charts and tables to reveal this periodicity.

Here’s one of them:

1. First Ionisation Energy

Here is a plot of first ionisation energy (E_m1) against atomic number (Z).

You can see the regular repetition of similar (not identical) properties.

Every eighth element has the highest E_m1 starting with Helium. 

We notice the break for the first transition series at atomic number 21 Scandium.

But the pattern picks up again at atomic number 30 and peaks at Kr Krypton.

The alkali metals are also labelled because they have the lowest Em1 values.

Their values repeat at regular intervals: every eighth element. 

We can see from this 3D image the trends in the first ionisation values:

E_m1 decreases down a group, e.g. lithium (Li) to caesium (Cs), because the outer shell electrons get further from the positive electrostatic pull of the nucleus.

E_m1 values generally increase across a period, e.g. Lithium (Li) to Neon (Ne), because the size of the positive nucleus increases, increasing the positive pull of the nucleus on the outer electron shell.

2. Electronegativity

What is electronegativity?

This is the power of an atom of an element within a covalent bond to attract the bonding pair of electrons to itself. 

So here is a water molecule which we can represent like this to show how the oxygen atom tends to pull electrons in the two covalent bonds towards itself ( the arrows) and lead to an internal dipole in the water molecule.

The two hydrogen atoms end up slightly positively charged relative to the oxygen atom. 

This ability to attract the bonding electrons to itself is called an element’s electronegativity.

Electronegativity is measured on a scale from 0-4 first developed by one of the greatest ever chemists Linus Pauling. 

Here is how electronegativity varies across the periodic table and as you can see it is a periodic function of the elements.  

The values and trends are easier to see on the next diagram.

The value increases across the periodic table from left to right and it decreases down a group of the periodic table. 

It is highest at Fluorine (F).  (Why are there no values for the Noble Gases?)

It is at its lowest at Caesium and Francium. 

Metals tend to have low values and non–metals have high values, hence the colour densities in the diagram above. 

The reason why the values tend in these directions is due to the dependence of electronegativity on the atomic radii values: the higher the electronegativity the smaller the atomic radius. 

And the atomic radius tends to depend on the size of the nucleus, the number of electron shells and the tendency for the inner shells to shield the outer shell electrons from the influence of the nucleus.

So take the Fluorine atom that has the highest electronegativity, it has two shells with electron arrangement 1s², 2s², 2p⁵.  There is little chance of shielding of the outer shell electrons so it pulls bonding electrons easily to itself. 

But look at Caesium with a very low electronegativity.  It has 6 electron shells.  It is in Period 6.  Its outer shell electron is well shielded from the positive nucleus even though the nucleus is huge (Z=55) with 55 protons. So there is a great tendency of the outer electron to be lost and caesium atoms are easily oxidised.

Things to do:  

So why do the Noble gases not have an electronegativity value? 

Can you construct a chart to reveal the periodicity of the melting points and boiling points of the elements?

Can you calculate the atomic volumes (the volume of one mole of atoms of the element, hint: use the molar mass and the element’s density.) of the elements and show that this is also a periodic property.