Periodicity - Trends Along Period 3 (A-Level Chemistry)

Periodicity

Periodicity refers to the repeating pattern in physical and chemical properties of the elements across a period of the periodic table.

In order to exemplify this, we will look at elements in Period 3:

Atomic Radius Trend Across Period 3

• The atomic radius trend of elements decreases as you move across a period. The proton number of the element increases as you move across a period which means that the number of protons increases. The nucleus gains a more positive charge as the number of protons increases. This means there is a stronger attraction between the nucleus and the electrons in the orbitals so the electrons are pulled inwards towards the nucleus. This contracts the radius inwards and therefore the atomic radius decreases.

Trend in Melting Point Across Period 3

• The melting point of the elements across a period changes according to their structure. Both structure and bonding have an effect on the melting point of an element as we learned earlier in the book.

We will now discuss the trend in melting points across Period 3:

1. The change in melting point from Sodium to Aluminium (Na to Al).

• Sodium, magnesium and aluminium are all metals and therefore have metallic bonding.
• As well learned earlier, the more delocalised electrons present and the smaller the radius of the atom, the higher the melting point of the metal.
• As we move across period 3 the number of delocalised electrons per metal atom increases and the radius of the elements decreases. This means the melting point increases.
• This is because there is a greater electrostatic attraction between the positive ions and delocalised electrons and hence the metallic bond is stronger and requires more energy to break.

2. The melting point of Silicon (Si)

• Silicon has a macromolecular structure which consists of covalent bonding.
• As we learned earlier, macromolecular structures have strong covalent bonds that hold the atoms together.
• These strong covalent bonds require a large amount of energy to break and therefore silicon has a high melting point.

3. The change in melting point from Phosphorous to Chlorine (P to Cl)

• Phosphorous (P4), sulfur (S8) and chlorine (Cl2) are simple molecular substances which consist of van der Waals forces.
• As learned earlier, van der Waals forces are weak intermolecular forces which require a small amount of energy to break. For this reason, the melting points of these simple molecular substances are low.
• The melting point of these substance depends on the varying strength of van der Waals forces. The shape of a molecule and the distance between the molecules affects the strength of induced dipole-dipole forces.
• The stronger the induced dipole-dipole forces, the higher the melting point. A larger molecule contains more electrons, therefore it consists of larger electron clouds. The greater the number of electron clouds, the stronger the induced dipole-dipole forces. More energy is required to break stronger induced dipole-dipole forces, therefore the melting point is higher.
• As sulfur is the largest molecule out of the three, it contains the most number of electrons and the strongest van der Waals forces. Therefore sulfur has the highest melting point, compared to phosphorous and chlorine.

4. The melting point of Argon (Ar)

• Argon is a noble gas and has a very low melting point as it exists as a monoatomic element consisting of very weak van der Waals forces.
• As the van der Waals forces in argon are weak, a very small amount of energy is needed to break them and therefore the melting point is low.

Electrical Conductivity Trend Across Period 3

• Electrical conductivity increases across metals. Sodium, magnesium and aluminum all exist as giant metallic lattices and hence are able to conduct electricity. As you move along the period, the number of mobile electrons donated by each atom into the sea of delocalised electrons increases so that the metal becomes better at conducting electricity.
• Electrical conductivity drops when we reach non-metals. As we move from aluminium to silicon, electrical conductivity drops dramatically. This is because silicon exists as a giant covalent structure, with no free ions or electrons capable of carrying the current. Silicon is a semiconductor. Electrical conductivity continues to decrease even further along the period.

First Ionisation Energy Trend Across Period 3

The first ionisation energy is the amount of energy required to remove one electron from the outermost shell of each atom in one mole of atoms of an element in their gaseous state.

• The first ionisation energy of the element increases across a period. As you move across a period, the number of protons increases and therefore the attraction between the nucleus and the outermost electrons increases. More energy is therefore required to remove the outermost electron, meaning the first ionisation energy increases.

The are however some exceptions to this trend:

1. The change in first ionisation energy from Magnesium to Aluminium (Mg to Al).

• The first ionisation energy of aluminium is slightly lower than that of magnesium.
• The outermost electron in aluminium is in a p orbital whereas that for magnesium is in an s orbital.
• Therefore, the outermost electron for aluminium is slightly further from the nucleus and experiencing slightly more shielding which means it will be less strongly attracted to the nucleus.

2. The change in first ionisation energy from Phosphorous to Sulphur (P to S).

• The first ionisation energy of sulphur is slightly lower than that of phosphorous.
• The outermost electron in sulphur is in a complete atomic orbital whereas the outermost electron in phosphorus is unpaired.
• As a result of spin pair repulsion, the outermost electron in sulphur is easier to remove.

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