(1) Variation of 1st Ionisation enthalpy across Period 3

ΔH for the process X_(g) ==> X⁺_(g) + e^-

The energy required to remove the most loosely bound electron (kJmol^-1) from the gaseous atoms at 298K/1atm.

The peaks correspond with the Noble Gases at the end of a period and the troughs with the Group 1 Alkali Metals at the start of a period.

As you go across the period from one element to the next, the positive nuclear charge is increasing by one unit as the atomic/proton number increases by one unit and the charge is acting on electrons in the same principal quantum level. The effective nuclear charge can be considered to be equal to the number of outer electrons (this is very approximate and NOT a rule) and this is increasing from left to right as no new quantum shell is added i.e. no extra shielding. Therefore the outer electron is increasingly more strongly held by the increasing positive charge of the nucleus and so, increasingly, more energy is needed remove it.

So, for Period 3, the Group1 Alkali Metal (sodium, lowest Z) has the lowest 1st ionisation energy and the Group 0/18 Noble Gas (argon, highest Z) has the highest 1st ionisation energy value.

However there are two anomalies in the atomic number versus 1st ionisation energy graphs for period 3.

A decrease from Mg [1s²2s²2p⁶3s²] to Al [1s²2s²2p⁶3s²3p¹]

3s3p orbitals ==>

The anomalously low values for aluminium is considered to be due to the first time a 3p electron is shielded by the full 3s sub-shell and probably being a bit further away from the nucleus than the 3s electrons. The effect overrides the effect of increasing proton number i.e. positive nuclear charge.

A decrease from P [1s²2s²2p⁶3s²3p³]  to S [1s²2s²2p⁶3s²3p⁴]

3s3p orbitals ==>

Prior to the 4th 3p electron, the other three p electrons occupy separate p sub-orbitals (Hund's Rule of maximum multiplicity to minimise repulsion between adjacent orbitals. The anomalously low values for sulphur is considered to be due to the effect of the first pairing of electrons in the 3p orbitals producing a small repulsion effect that overrides the effect of increasing proton number (positive nuclear charge).

See also 5.3 Period 3 element trends in bonding, structure, oxidation state, formulae & reactions and 6.4 Important element trends down a Group

(2) Variation of atomic radius across Period 3

Can be defined as volume within which 95% of the electron charge exists on a time averaged basis.

The peaks correspond with the Group 1 Alkali Metals at the start of a period and the troughs with the Group 0 Noble Gases at the end of a period.

It generally decreases from left to right across a period, as the actual and effective nuclear charge increases within the same principal quantum level with increase in proton number, pulls the electron cloud closer to the nucleus without any increase in shielding. The argument is almost identical to that for increasing ionisation energy.

So, for Period 3, the Group1 Alkali Metal (sodium, lowest Z) has the largest atomic radius and the Group 7/17 Halogens & Group 0/18 Noble Gas (chlorine & argon, highest Z's) have the smallest atomic radii (there is some uncertainty in the noble gas radii).

See also 5.3 Period 3 element trends in bonding, structure, oxidation state, formulae & reactions and 6.4 Important element trends down a Group

(3) Variation of electronegativity across Period 3 (Pauling scale)

The power of an atom, in terms of an electric field effect, to attract electron charge towards it, in the context of a pair of electrons of a covalent bond linking it to another different atom.

The peaks correspond to the Group 7 Halogens/Group 0 Noble Gases at the end of a period and the troughs' correspond to the most electropositive Group 1 Alkali Metals at the start of a period.

It generally increases from left to right across a period, as the actual and effective nuclear charge increases within the same principal quantum level, pulling the bonding electron cloud (bonding pair of electrons) closer to the nucleus (see 1st IE arguments) i.e. increase in proton charge without increase in shielding. The argument is almost identical to that for increasing ionisation energy.

So, for Period 3, the Group1 Alkali Metal (sodium, lowest Z) has the lowest electronegativity and the Group 7 Halogen & Group 0/18 Noble Gas (chlorine & argon, highest Z's) have the highest electronegativities (there is some uncertainty in the noble gas electronegativities).

In the context of a bond between two different elements, the element with the greater electron pulling power aquires a partial negative charge and the other less electronegative element a partial positive charge.

So, in the covalent bond M^δ+-X^δ-, X has the greater electronegativity e.g. the polar bond Si^δ+-Cl^δ- in covalent SiCl₄.

This has major consequence on the type of bonding from ionic oxides and chlorides to non-metallic covalent oxides and chlorides. If the difference is large an ionic bond results. e.g. Na⁺ Cl^-

See also 5.3 Period 3 element trends in bonding, structure, oxidation state, formulae & reactions and 6.4 Important element trends down a Group

(4) Variation of melting points and boiling points across Period 3

Trends in melting/boiling point can be complicated because of significant differences in the structure of the element.

The melting points and boiling points tend to peak in the middle of  Periods 2 and 3 (Groups 3/13 and 4/14) and the lowest values at the end of the period - the Noble Gases.

Generally you are moving from a low melting, but the still quite high boiling, metallic lattice of sodium in Group1 of moderately strong bonding with one outer delocalised valence electron  ==> a much higher melting/boiling metallic lattice with much stronger bonding due to 2/3 outer electrons for Mg/Al respectively.

At Group 4 you have a very high melting giant covalent lattice of a strong 3D network of covalent bonds.

From Group 5 onwards there is a dramatic fall as the elements now consist of low melting small covalent molecules (P₄, S₈, Cl₂ and Ar) only held together by weak inter-molecular forces (transient dipole - induced dipole interactions).

See also 5.3 Period 3 element trends in bonding, structure, oxidation state, formulae & reactions and 6.4 Important element trends down a Group

(5) Variation of relative electrical conductivity across Period 3

The peaks correspond to the metals in the middle of the period with the greatest number of outer electrons that can be delocalised.

Increases dramatically from left to right for Groups 1 ==> 3 as the metallic lattice contains 1 ==> 3 mobile delocalised electrons involved in electrical conduction.

From Group 4 to 0 the element structure changes to giant covalent lattice or simple molecular structures with no free delocalised electrons within the structure to convey an electric current. Although silicon can be described as a semi-metal (wrongly in my view) it is virtually an insulator unless doped with other elements.

See also 5.3 Period 3 element trends in bonding, structure, oxidation state, formulae & reactions and 6.4 Important element trends down a Group

(6) Variation of density across Period 3

The peaks correspond to the metals in the middle of the period with the strongest bonding in the solid.

The density increases from sodium to aluminium as the atomic radii decrease and the bonding gets stronger with 1 ==> 3 bonding electrons (delocalised outer valency electrons in the metal lattice).  However, they are relatively low densities compared to most metals.

Silicon, phosphorus and sulphur have a low densities, typical of non-metallic covalent solids.

Chlorine and argon are small covalent molecules and have very low densities being gaseous at room temperature because only weak intermolecular forces act between them.

See also 5.3 Period 3 element trends in bonding, structure, oxidation state, formulae & reactions and 6.4 Important element trends down a Group