Physical properties of Period 3 elements

The trends in atomic radius, first ionisation energy and melting/boiling points of the elements Na–Ar

Students should be able to:
• explain the trends in atomic radius and first ionisation energy
• explain the melting point of the elements in terms of their structure and bonding.

Key unifying theory : Effective nuclear charge density

Key definitions

Ionisation Energy: ΔH (energy, measured in Joules/mol) needed to remove 1 mole of electrons from 1 mol of gaseous atoms to form 1 mole of mono-positive gaseous ions

Al(g) → Al+ (g) + e− ΔH= 1st ionisation energy of Al

Al+(g) → Al2+(g) + e− ΔH= 2nd ionisation energy of Al

Electronegativity: a measure of the tendency of an atom to attract electron density within a covalent bonding pair of electrons.

The common theme that accounts for ALL the Trends across Period 3 is an increase in the effective nuclear charge density across Period 3.

The overall attraction between the nucleus and outer shell (highest energy level) valence electron(s) is called the effective nuclear charge density.

Effective nuclear charge density depends on:

• the number of nuclear protons (attractive force): increased proton number across Period 3
• the extent of shielding by the inner electrons: there is similar electron shielding across Period 3, as electrons are being added to the same main energy shell
• the distance between the nucleus and outer shell valence electrons (atomic radii): decreased atomic radii across Period 3

Consequently, increased effective nuclear charge density across Period 3 results in:
2)Higher electronegativity
3)Higher first ionisation energy (with two notable exceptions: Al, S)
Aluminium has a lower than expected increase in first ionisation energy because the valence e− of Aluminium is located in a higher energy 3p orbital which is further away from the nucleus and so less energy is needed to remove the electron.

Sulfur has a lower than expected increase in first ionisation energy because there is a repulsion between the electron pair occupying the 3p orbital subshell in Sulfur, so less energy is needed to remove the sulfur’s valence electron.

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Trends in melting points across Period 3

Na, Mg, Al, Si are all GIANT LATTICE structures involving either metallic (Na, Mg, Al) or giant covalent macromolecule (Si) bonding.

Whereas, P, S, Cl are COVALENT MOLECULES, with Ar being a monatomic non-metal gas. Melting is achieved by disrupting the considerably weaker van der Waals intermolecular forces that exist between the molecules.

Melting point steadily increases between Na and Al due to: increased metal cation charge (more protons), more delocalised e− per atom, smaller sized metal ions, hence stronger electrostatic attraction between cations and delocalised e− therefore stronger metallic bonding.

Silicon has a significantly higher melting point because Silicon is a giant covalent macromolecule. In order to melt Silicon there is a need to break down many strong covalent bonds between Si atoms, that can only be achieved by applying considerable energy.

Melting point of S8 > P4 > Cl2 > Ar because there is decreasing strength of vdW intermolecular forces of attraction as there is a decreasing size of molecule. This explanation is based on the known correlation between molecular size (molecular mass) and magnitude of vdW force of attraction between molecules.

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Worked examples of questions

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