Notes · ChemistryUK · A-Levels

Periodicity

The Periodic Table arranges the elements so that their properties repeat in a regular, predictable pattern. This chapter classifies the elements into s, p and d blocks by their electron configuration, and explains the trends across Period 3 in atomic radius, first ionisation energy and melting point in terms of structure and bonding - the clearest demonstration that periodic properties follow from electron arrangement.

3 sections~9 min reading time3 competenciesLevel Foundation 1 · Standard 2

T·0888 / 18

Exam profile

AO1 · Describe the classification of elements into s, p and d blocks and the Period 3 trendsAO2 · Explain the trends in atomic radius, ionisation energy and melting point from structure and bondingAO3 · Interpret periodic-property data and account for anomalies such as the aluminium and sulfur ionisation energies

Operators:statedescribeexplainpredictinterpret

basic level

AS-Level requires the block classification and the trends across Period 3 in atomic radius, first ionisation energy and melting point, with their explanations.

higher level

The full A-Level uses these trends synoptically to underpin the group chemistry and transition-metal chemistry that follow, and to interpret unfamiliar data.

Depth

Text

Contents · 3 sections

Periodicity

The Periodic Table and the s, p and d blocks#

FoundationAQA 7405 3.2.1.1DfE GCE Chemistry - periodicity

The Periodic Table lists the elements in order of increasing atomic number, arranged so that elements with similar properties fall in the same vertical group. A horizontal row is a period, and the period number equals the highest occupied principal quantum shell. Elements in the same group have the same number of outer-shell electrons, which is why they have similar chemical properties - chemistry is controlled by the outer electrons.

The table divides into blocks according to which type of sub-shell is being filled by the outer electrons. The s block (Groups 1 and 2) is filling an s sub-shell; the p block (Groups 3 to 0/8) is filling a p sub-shell; the d block (the transition elements) is filling a d sub-shell; and the f block (the lanthanides and actinides) an f sub-shell. Sodium (

1s^2\,2s^2\,2p^6\,3s^1

) is in the s block, chlorine (

\dots 3s^2\,3p^5

) in the p block.

Periodicity is the name for the repeating pattern of properties as you move across successive periods. A property such as first ionisation energy rises and falls in a similar way across each period, because the same pattern of filling sub-shells recurs. The regularity is a direct consequence of electron configuration: elements with analogous outer configurations sit in the same group and behave alike.

Being able to place an element in its block and group from its electron configuration, and vice versa, is the foundation of the whole topic. The last sub-shell to receive an electron identifies the block; the outer-shell electron count identifies the group. This link between configuration and position is what makes the trends across Period 3 explicable rather than a set of facts to be memorised.

Worked example

Placing an element from its configuration

An element has the configuration 1s2 2s2 2p6 3s2 3p3. State its period, block and group.

Period
The highest occupied shell is n = 3, so the element is in Period 3.
Block
The outer electrons are entering the 3p sub-shell, so it is in the p block.
Group
There are 5 outer-shell electrons (3s2 3p3), so it is in Group 5 - the element is phosphorus.

Result: Period 3, p block, Group 5: phosphorus.

Typical mistakes

Confusing the block (which sub-shell is filling) with the group number.
Forgetting that period number equals the highest occupied shell.

Active revision

State the block and group of the element with electron configuration $1s^2\,2s^2\,2p^6\,3s^2\,3p^4$ , and name it.

Active recall

Recall the key points — then reveal.

Sources: GCE AS and A level subject content for the sciences (Department for Education) · AQA A-level Chemistry 7405 specification (AQA)

Trends in atomic radius and ionisation energy across Period 3#

StandardAQA 7405 3.2.1.2DfE GCE Chemistry - Period 3 trends

Atomic radius across Period 3

Fig. 1Atomic radius decreases across the period as the rising nuclear charge pulls in the same outer shell.

Atomic radius decreases across Period 3 from sodium to chlorine. Moving across the period, each successive element has one more proton in the nucleus and one more electron, but the extra electron is added to the same outer shell, so the shielding by inner shells stays essentially the same. The increasing nuclear charge therefore pulls the same outer shell in more strongly, and the atoms get smaller. This shrinking radius underlies several of the other trends.

First ionisation energy generally increases across Period 3, because the rising nuclear charge (with similar shielding and a decreasing radius) holds the outer electrons more tightly, so more energy is needed to remove one. The trend is not perfectly smooth: there are two dips that reveal the sub-shell structure. Between magnesium and aluminium the value falls, because aluminium's outer electron is in a

3p

orbital, which is higher in energy and slightly shielded by the filled

3s

sub-shell, so it is easier to remove.

The second dip, between phosphorus and sulfur, arises from electron pairing. Phosphorus has three

3p

electrons, one in each

3p

orbital; sulfur has four, so one orbital now holds a pair. The repulsion between the two paired electrons in that orbital makes one of them easier to remove, so sulfur's first ionisation energy is slightly lower than phosphorus's. These two anomalies are direct experimental evidence for the existence and ordering of the

s

and

p

sub-shells.

Interpreting the data requires linking each feature to structure. When you explain the general rise, cite increasing nuclear charge with similar shielding; when you explain a dip, name the specific sub-shell effect (the

3p

electron for Al, electron pairing for S). This is a classic AO2/AO3 skill: reading a graph of ionisation energy and accounting for both the overall trend and the anomalies.

Worked example

Explaining the magnesium-to-aluminium dip

The first ionisation energy of aluminium (578 kJ mol^-1) is lower than that of magnesium (738 kJ mol^-1). Explain why.

Compare the outer electrons
Magnesium's outer electron is removed from a 3s orbital; aluminium's is removed from a 3p orbital.
Relate the energies
The 3p sub-shell is higher in energy than 3s, and the 3p electron is slightly shielded by the filled 3s sub-shell.
Conclude
So aluminium's outer electron is held less tightly and is easier to remove, giving a lower first ionisation energy despite the greater nuclear charge.

Result: Aluminium's outer electron is in a higher-energy, shielded 3p orbital, so it is easier to remove than magnesium's 3s electron.

Typical mistakes

Explaining the aluminium dip by nuclear charge instead of the 3p electron being higher in energy and shielded by 3s.
Saying the shielding increases significantly across a period - it stays roughly constant as electrons enter the same shell.

Active revision

Explain why the first ionisation energy of sulfur is lower than that of phosphorus, even though sulfur has a greater nuclear charge.

Active recall

Recall the key points — then reveal.

Sources: AQA A-level Chemistry 7405 specification (AQA)

Trends in melting point across Period 3#

StandardAQA 7405 3.2.1.3DfE GCE Chemistry - melting points

Melting point across Period 3

Fig. 2The sharp peak at silicon (giant covalent) and the collapse to the molecular non-metals reflect the change in structure.

The melting point across Period 3 does not follow a single trend, because the type of structure and bonding changes across the period. Sodium, magnesium and aluminium are metals with giant metallic lattices; silicon is a giant covalent (macromolecular) solid; and phosphorus, sulfur, chlorine and argon are simple molecular (or, for argon, atomic). What melts, and therefore how much energy is needed, is completely different for each type.

Across the three metals the melting point rises from sodium to aluminium. Each has metallic bonding, but the charge on the ion increases from

\text{Na}^+

\text{Mg}^{2+}

\text{Al}^{3+}

, so more delocalised electrons are released per atom and the ions are smaller. The stronger electrostatic attraction between the more highly charged ions and the denser electron sea means more energy is needed to melt them, so the melting point climbs.

Silicon has by far the highest melting point in the period, because it is a giant covalent structure like diamond: melting it requires breaking a huge number of strong covalent bonds throughout the lattice. After silicon the melting point falls sharply, because phosphorus, sulfur and chlorine are simple molecular substances (

\text{P}_4

\text{S}_8

\text{Cl}_2

) held together only by weak van der Waals forces between the molecules. Melting these overcomes only the weak intermolecular forces, not the covalent bonds inside the molecules.

Among the molecular non-metals the melting points depend on molecular size: sulfur (

\text{S}_8

) is higher-melting than phosphorus (

\text{P}_4

) because its larger molecules have more electrons and stronger van der Waals forces, while chlorine (

\text{Cl}_2

) and argon (single atoms) are the lowest. When you explain any melting point, state the structure first, then name the force being overcome - covalent bonds for silicon, metallic bonding for the metals, van der Waals forces for the molecular elements.

Worked example

Explaining the drop from silicon to phosphorus

Explain why the melting point falls sharply from silicon (1687 K) to phosphorus (317 K).

Silicon's structure
Silicon is a giant covalent solid; melting it breaks many strong covalent bonds, needing a great deal of energy.
Phosphorus's structure
Phosphorus is simple molecular (P4); its molecules are held together only by weak van der Waals forces.
Conclude
Melting phosphorus overcomes only these weak intermolecular forces, not covalent bonds, so far less energy is needed and the melting point is much lower.

Result: The change from a giant covalent lattice to weak van der Waals forces between P4 molecules causes the sharp fall in melting point.

Typical mistakes

Saying covalent bonds break when phosphorus or sulfur melts - only the van der Waals forces between molecules do.
Treating the whole period as one trend instead of dividing it by structure type.

Active revision

Explain why silicon has a much higher melting point than both aluminium (a metal) and phosphorus (a molecular solid).

Active recall

Recall the key points — then reveal.

Sources: AQA A-level Chemistry 7405 specification (AQA)

References & sources

Sources

Department for Education

GCE AS and A level subject content for the sciences

AQA

AQA A-level Chemistry 7405 specification

Periodicity

3 sections~9 min reading time3 competenciesLevel Foundation 1 · Standard 2

Periodicity

Key points

Placing an element from its configuration

Key points

Explaining the magnesium-to-aluminium dip

Key points

Explaining the drop from silicon to phosphorus

Sources

Periodicity

Key points

Placing an element from its configuration

Key points

Explaining the magnesium-to-aluminium dip

Key points

Explaining the drop from silicon to phosphorus

Sources