Periodicity

Trends on Period – Properties of Atoms

Trends on Period – Properties of Atoms

  • Atomic size: as you move from left to right the nuclear charge of the atoms becomes larger but the shielding does not change. The outer electrons are, therefore, pulled in closer due to their attraction towards the nucleus. On crossing a period, the size of the atoms decreases.
  • Ionisation energy: in general, the ionisation energy increases as you move across period because the nuclear charge increases but the shielding does not change. However, there is a decrease between groups II and III and V and VI for different reasons. The decrease occurs between groups II and III because because Al consists of a 3p orbital for its outer shell while Mg is in a 3s orbital. It is easier to remove the 3p orbital as it is more efficiently shielded by its nucleus. Between groups V and VI, S consists of a paired 3p outermost electron. The repulsion in the orbital makes the electron easier to remove. P, on the other hand, has an unpaired 3p outermost electron which is harder to remove as there is less repulsion.
  • Electronegativity: across period 3 there is an increase in electronegativity. This is because there is no change in shielding but an increase in nuclear charge meaning the electrons have a stronger attraction to their atom and consequently more covalent bonds.

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Structure and bonding

In terms of structure and bonding, the elements widely vary. Across the period the metallic character decreases at a gradual rate.

Moving from left to right covalent bonding becomes more prevalent due to the increase in ionisation energy which means it is harder for electrons to be removed and metallic structures to form.

The noble gases have very high ionisation energies making metallic bonding impossible. Covalent bonding is not possible either as they do not contain any paired electrons. Without bonds they are able to exist as gaseous atoms.

  • Sodium, Magnesium and Aluminium: these elements are metals and, therefore, have quite high melting and boiling points. These points increase across a period as the charge of the elements increases and the size decreases. Electricity conductivity also increases as there is an increase in the number of delocalised elements per atom: sodium has one per atom, magnesium two per atom, and aluminium three per atom.
  • Silicon: this is a giant covalent macromolecule and, therefore, has very high melting and boiling points and does not conduct electricity well.
  • Phosphorus, sulphur, chlorine and argon: these form simple molecular structures and, therefore, have relatively low melting and boiling points. The points increase with the size of the molecule as the bigger the molecule the higher the magnitude of temporary and induced dipoles. As it exists as S8 molecules sulphur has the highest melting point. S8 molecules are relatively large meaning that the molecules contains a lot of electrons and quite strong Van der Waal’s forces. Phosphorus, on the other hand, contains less electrons and weaker Van der Waal’s forces which means that it has a lower melting point than sulphur. Chlorine has fewer electrons still and lower Van der Waal’s forces as it exists molecules while Argon, with the lowest melting and boiling points, exits as single atoms. These elements also do not conduct electricity very well.

Properties summary

Property type Trend
Atomic size Decreases across the period
First ionisation energy Increases across the period (except between Mg and Al, and P and S)
Electronegativity Increases across the period
Melting and boiling points Increases from Na to Al, Al to Si, and P to SDecreases from Si to P, and S to Ar
Electrical conductivity Increases from Na to AlZero from Si to Ar