Periodicity melting points and ionisation energies

Graphical Data and Periodic Properties

A periodic property is any physical property whose value rises and falls in a repeating pattern when elements are plotted in order of increasing atomic number.

Plotting properties such as melting temperature, atomic radius, or first ionisation energy against atomic number for elements 1 to 36 reveals clear repeating peaks and troughs across each period. This repetition is what defines a property as “periodic”, because the pattern recurs every time a new period begins.

First ionisation energies for elements 1 to 36 span a very large numerical range, from a few hundred to several thousand kJ mol⁻¹. Plotting raw values compresses the lower numbers and obscures fine detail in the early periods.

Taking the logarithm of the first ionisation energy before plotting reduces the spread, so the periodic pattern becomes visually clearer and small dips within each period are easier to identify.

Examiner InsightWhen asked to define a periodic property, state explicitly that the property repeats at regular intervals across the periodic table. A vague answer such as “changes across a period” does not earn the mark.

Exam TipUse the word “repeats” or “repeating pattern” to secure the mark.

Trends in Melting Points and Ionisation Energies

Across Periods 2 and 3, melting temperatures, boiling temperatures, and ionisation energies all change in patterns that depend directly on structure, bonding, and electron configuration.

Melting and boiling temperatures across Period 3 rise from sodium to silicon, then fall sharply, because the type of structure and bonding changes across the period. Period 2 follows the same logic with lithium to carbon rising, then a sharp fall to nitrogen, oxygen, fluorine, and neon.

The table below summarises the structure and bonding for each Period 3 element and explains why its melting temperature is high or low.

Period 2 follows the same structural logic: Li, Be metallic; B, C giant covalent (with C giving the highest melting point); N₂, O₂, F₂, Ne simple molecular or monatomic with low melting points.

The first ionisation energy is the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.

The defining equation is shown below:

$$X(g)\to {X}^{+}(g)+{e}^{-}$$

Across Periods 2 and 3, first ionisation energy generally increases because nuclear charge increases while electrons are added to the same main shell. Shielding stays roughly the same, so the outer electron is held more strongly and harder to remove.

Two specific dips break this general increase in each period.

The first dip occurs between Group 2 and Group 3 (Be → B in Period 2; Mg → Al in Period 3). The outer electron in Group 3 is in a 2p (or 3p) orbital, which is at a slightly higher energy than the 2s (or 3s) orbital. This electron is therefore easier to remove despite the higher nuclear charge.

The second dip occurs between Group 5 and Group 6 (N → O in Period 2; P → S in Period 3). In Group 5 every p orbital holds one electron, but in Group 6 one p orbital holds a paired electron. The two paired electrons repel each other, so one is more easily removed.

Down a group, first ionisation energy decreases because three effects combine. The outer electron occupies a shell further from the nucleus. More inner shells provide greater shielding from the nuclear charge. These two effects outweigh the increase in nuclear charge, so the outer electron is held less strongly.

MisconceptionThe dip from N to O is not because oxygen has “more electrons” or “a higher nuclear charge”. It is because pairing two electrons in the same 2p orbital introduces repulsion, lowering the energy needed to remove one of those paired electrons.

Exam TipState “paired electrons in the same p orbital repel” to secure the mark.