12ChF321 Periodicity

Periodicity

Candidates should be able to:(a) describe the Periodic Table in terms of the arrangement of elements:

(i) by increasing atomic (proton) number,(ii) in periods showing repeating trends in physical and chemical properties,(iii) in groups having similar physical and chemical properties;

(b) describe periodicity in terms of a repeating pattern across different periods;(c) explain that atoms of elements in a group have similar outer shell electron configurations, resulting in similar properties;(d) describe and explain the variation of the first ionisation energies of elements shown by:

(i) a general increase across a period, in terms of increasing nuclear charge,(ii) a decrease down a group in terms of increasing atomic radius and increasing electron shielding outweighing increasing nuclear charge;

(e) for the elements of Periods 2 and 3:(i) describe the variation in electron configurations, atomic radii, melting points and boiling points,(ii) explain variations in melting and boiling points in terms of structure and bonding;

(f) interpret data on electron configurations, atomic radii, first ionisation energies, melting points and boiling points to demonstrate periodicity.

The Periodic TableEarly attempts to create a periodic table were based on atomic mass, but this was unsatisfactory as similar elements were not collected together (into groups).

The modern periodic table is arranged in order of increasing atomic number (proton number), with a new row (period) started every time electrons occupy a new shell. This results in the elements in each group having similar physical and chemical properties, and in repeating patterns or trends in physical and chemical properties across each period.

Periodicity – a trend which is repeated across different periods.

The periods in the Periodic Table show repeating trends in both physical and chemical properties. These repeats occur because elements in the same Group as each other have similar physical and chemical properties. Elements in the same group have similar properties because they have similar outer shell electron configurations.

We've seen periodicity already when we examined how Ionization Energies changed both down each Group and across a Period (see notes on Ionization Energy). Using Period 3 as an example, we can start to examine some other properties which show periodicity.

* bonding becomes increasingly covalent in character across the period

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Period 3 Na Mg Al Si P S Cl ArElectron arrangement

[Ne] 3s1

[Ne] 3s2

[Ne] 3s3 3p1

[Ne] 3s2 3p2

[Ne] 3s2 3p3

[Ne] 3s2 3p4

[Ne] 3s2 3p5

[Ne] 3s2 3p6

Structure Giant metallic lattice

Giant metallic lattice

Giant metallic lattice

Giant covalent lattice

Simple molecularP4

Simple molecularS8

Simple molecularCl2

Single atoms

Bonding* metallic metallic metallic covalent covalent Covalent Covalent N/A

12ChF321 Periodicity

BondingBonding in the Period 3 elements and their compounds becomes increasingly covalent in character across the period. Aluminium shows ionic bonding in some compounds and polar covalent bonding in others. This is because Aluminium is more electronegative than Group 1 or 2 metals, so there is less of a difference in electronegativity between aluminium and the non-metals it bonds with.

Atomic radiusProblem: electron clouds are fuzzy ! How can we measure to the edge of them ? Best solution is to find where two atoms of the element are bonded together and measure from nucleus to nucleus. This is why we can’t measure it for noble gases.

Period 3 Na Mg Al Si P S Cl ArAtomic radius (nm)

0.190 0.160 0.130 0.118 0.110 0.102 0.099 N/A

Note: 1nm = 1 x 10-9m

Trend: The atomic radius decreases across a periodWhy: Increasing atomic number causes increasing proton number and hence greater nuclear charge, so the outer electrons are pulled closer to the nucleus. Remember all the elements across a period have the same shielding from filled inner shells, so this is not a factor.

Electrical conductivity

Period 3: Na Mg Al Si P S Cl ArConductivity Good Good Good Poor Poor Poor Poor Poor

… and getting better … and getting worse

Electrical conductivity is related to the type of bonding. Na, Mg and Al have metallic bonding. They are good conductors because they possess mobile delocalized electrons which can carry the charge. Al is a better conductor than Na because it has 3 outer shell electrons rather than 1. These outer shell electron are the ones which are donated into the “sea” of mobile electrons, being able to move and thereby conduct electricity (and heat).

The remaining elements in the period have covalent bonding, and do not contain any mobile free electrons, so they are poor conductors. They get increasingly poor because an electron has to be removed from the outer shell in order for conduction to take place, and the ionization energy increases across the period.

Melting point and boiling point

Period 3 Na Mg Al Si P S Cl Armp °C 98 649 660 1410 44 113 -101 -189bp °C 883 1107 2467 2355 281 445 -35 -186

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12ChF321 Periodicity

Melting points relate to the type of structure. The more energy needed to break down the structure so that the particles can move around, the higher the melting point will be. Elements with a giant lattice (giant metallic lattice Na, Mg and Al, and the giant covalent lattice of Si) all require a lot of energy to break the strong bonds holding the atoms together in the lattice, so the melting points are high.

From sodium to aluminium the melting and boiling points increase because:- the ions have greater charge and so are more strongly attracted to the delocalized electrons

- there are an increasing number of delocalized electrons for the metal ions to be attracted to (Al has given up thee per atom, whereas Na only one per atom)

- so the metallic bonding gets stronger from Na to Al

Elements with simple molecular structures have only weak intermolecular forces holding the molecules to one another. To melt these so that the molecules can move around we only have to overcome these weak intermolecular forces, which requires little energy, so the melting points are low. We DON’T have to break the strong covalent bonds holding atoms together inside each molecule. Trends in boiling points follow the same pattern as the melting points, at a higher temperature.

For P4, S8 and Cl2 the size of the intermolecular forces (and therefore how high the melting and boiling points are) depends on the number of electrons in the molecule. The more electrons there are, the stronger the intermolecular forces (Van der Waals forces – we'll meet these in more detail later).

S8 has a higher melting and boiling point than P4, which in turn are higher than Cl2 because:- S8 has stronger intermolecular forces- because S8 has more electrons in the molecule

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