(1) Variation of 1st Ionisation enthalpy across Period 2

ΔH/kJmol^-1 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 Group 0/18 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 2, the Group1 Alkali Metal (lithium, lowest Z) has the lowest 1st ionisation energy and the Group 0/18 Noble Gas (neon, highest Z) has the highest 1st ionisation energy value and most values follow the general trend of increasing from left to right across period 2.

However there are two anomalies in the atomic number versus 1st ionisation energy graphs for period 2.

It should be first pointed out that these anomalies are due to the complex behaviour of the quantum levels in multi-electron systems - do not expect any perfect trends in chemistry, thanks to quantum physics!

(i) A decrease from Be [1s²2s²] to B [1s²2s²2p¹]

Box diagram of 2s2p orbitals ==>

The anomalously low value for boron is considered to be due to the first time a 2p electron is shielded by the full 2s sub-shell and, more importantly, the 2p electron is a  bit further away on average from the nucleus than the 2s electrons (so less strongly bound). The effect overrides the effect of increasing proton number i.e. the increase in positive nuclear charge from Be to B. However, after the kink, the continued increase in nuclear charge ensures the Period 2 trend for the 1st ionisation energy continues as expected until oxygen, the 2nd anomaly.

(ii) A decrease from N [1s²2s²2p³] to O [1s²2s²2p⁴]

Box diagram of 2s2p orbitals ==>

Prior to the 4th 2p 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 value for oxygen is considered to be due to the effect of the first pairing of electrons in the p orbitals producing a repulsion effect that overrides the effect of increasing proton number (increase in positive nuclear charge). From the 'kink', the Period 2 trend for the 1st ionisation energy continues as expected from oxygen to neon with increase in nuclear charge.

You see the same anomalous pattern in Period 3

See also 4.3 Period 2 element trends in bonding, structure, oxidation state, formulae & reactions and 6.4 Important element trends down a Group

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(2) Variation of atomic radius across period 2

It can be described as the radius of the 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, pulling 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 2, the Group1 Alkali Metal (lithium, lowest Z) has the largest atomic radius and the Group 7/17 Halogens & Group 0/18 Noble Gas (fluorine & neon, highest Z's) have the smallest atomic radii (there is some uncertainty in the noble gas radii).

See also 4.3 Period 2 element trends in bonding, structure, oxidation state, formulae & reactions and 6.4 Important element trends down a Group

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(3) Variation of electronegativity across Period 2 (Pauling scale)

Electronegativity is 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 2, the Group1 Alkali Metal (lithium, lowest Z) has the lowest electronegativity and the Group 7 Halogen & Group 0/18 Noble Gas (fluorine & neon, 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 acquires 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 C^δ+-F^δ- in covalent CF₄.

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. Li⁺ F^-

See also 4.3 Period 2 element trends in bonding, structure, oxidation state, formulae & reactions and 6.4 Important element trends down a Group

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(4) Variation of melting points and boiling points across Period 2

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 still quite high boiling, metallic lattice of lithium in Group1 of moderately strong bonding with one outer delocalised valence electron  ==> a much higher melting/boiling metallic lattice with 2 outer electrons for beryllium.

For groups 3/4 boron/carbon (B/C) you have a very high melting giant covalent lattices of a strong 3D or 2D network of strong covalent bonds. The mpts and bpts are even higher than the preceding metals because 3/4 outer valence electrons are involved in the bonding.

From Group 5 onwards there is a dramatic fall as the elements now consist of low melting small covalent molecules (N₂, O₂, F₂ and Ne) only held together by weak inter-molecular forces (transient dipole - induced dipole interactions).

See also 4.3 Period 2 element trends in bonding, structure, oxidation state, formulae & reactions and 6.4 Important element trends down a Group

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(5) Variation of relative electrical conductivity across Period 2

Not surprisingly, the highest values correspond to the metals at the start of the period with the greatest number of outer electrons that can be delocalised.

Increases dramatically from left to right for Groups 1-2 as the metallic lattice contains 1-2 mobile delocalised electrons involved in electrical conduction.

From Group 3 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 the graphite allotrope of carbon conducts electricity via the delocalised electrons in the linked hexagons of carbon atoms, it is still a very poor electrical conductor compared to metals (diamond is virtually an insulator).

See also 4.3 Period 2 element trends in bonding, structure, oxidation state, formulae & reactions and 6.4 Important element trends down a Group

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(6) Variation of density across Period 2

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

The density increases from lithium to beryllium as the atomic radii decrease and the bonding gets stronger with 1 ==> 2 bonding electrons (delocalised outer valency electrons in the metal lattice).

However, they are relatively low densities compared to most metals.

Boron and silicon have a low density, typical of non-metallic covalent solids.

Nitrogen, oxygen, fluorine and neon are small covalent molecules and have very low densities being gaseous at room temperature because only weak intermolecular forces act between them.

Even the densities of the liquid are quite low, again typical of low atomic number non-metals.

See also 4.3 Period 2 element trends in bonding, structure, oxidation state, formulae & reactions and 6.4 Important element trends down a Group

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