Periodicity

In a nutshell

Periodicity is the repeating pattern in the properties of elements as you move across a period of the Periodic Table.

This subtopic does two things. First, it sorts every element into an s, p, d or f block from where its highest energy electron sits. Second, it looks at how three physical properties of the Period 3 elements (atomic radius, first ionisation energy and melting point) change across the row from Na to Ar, and why.

The "why" is where the marks are: every trend has to be explained in terms of structure and bonding, not just described.

Assumed knowledge: Atomic structure, Bonding.

Core content

Classification: the s, p, d and f blocks

An element is classified as s, p, d or f block according to the sub-shell (sub-level) that holds its highest energy electron (the last electron added). An element's position in the table, and so its block, is fixed by its proton number.

• s block: highest energy electron in an s sub-shell. Groups 1 and 2 (plus H and He).
• p block: highest energy electron in a p sub-shell. Groups 13 to 18 (3 to 0), except He.
• d block: highest energy electron in a d sub-shell. The transition elements in the centre of the table.
• f block: highest energy electron in an f sub-shell. The lanthanides and actinides, shown as two detached rows below the main table.

Two quick reads:

• Na is 1s² 2s² 2p⁶ 3s¹: the last electron is in a 3s sub-shell, so Na is in the s block.
• Al is 1s² 2s² 2p⁶ 3s² 3p¹: the last electron is in a 3p sub-shell, so Al is in the p block.

Trend in atomic radius (Na to Cl)

Atomic radius decreases across the period.

The causal chain (this is an explain, so give every link):

1. Going across the period the number of protons increases, so the nuclear charge increases.
2. The outer electrons are being added to the same shell, so the shielding stays about the same (similar shielding).
3. So there is a greater attraction between the nucleus and the outer electrons, which are pulled in closer, and the atomic radius gets smaller.

Argon is normally left off the atomic-radius graph. It forms no bonds, so it has no comparable (metallic or covalent) radius: the only radius you can measure for it is a van der Waals radius, which is larger and not on the same scale.

Trend in first ionisation energy (Na to Ar)

The first ionisation energy is the energy needed to remove one electron from each atom in one mole of gaseous atoms, to form one mole of gaseous 1+ ions. The conditions in bold are exactly the words that earn the mark.

The general equation (state symbols are required):

Na(g) ⟶ Na⁺(g) + e^-

The general trend is an increase across the period. Explanation:

1. The number of protons increases, so the nuclear charge increases.
2. The outer electrons are in the same shell, so there is similar shielding, and the atomic radius decreases.
3. So there is a stronger attraction between the nucleus and the outer electron, and more energy is needed to remove it.

Two drops break this trend, and AQA expects you to explain both.

Drop 1: Mg to Al.

• Magnesium's outer electron is in a 3s sub-shell; aluminium's outer electron is in a 3p sub-shell.
• The 3p sub-shell is higher in energy than the 3s (and is slightly shielded by the 3s electrons), so aluminium's outer electron is easier to remove.
• So aluminium has a lower first ionisation energy than magnesium.

Drop 2: P to S.

• In phosphorus (3p³) each of the three 3p orbitals holds a single electron. In sulfur (3p⁴) one 3p orbital now holds a pair of electrons.
• The two paired electrons in the same 3p orbital repel each other, so one is easier to remove.
• So sulfur has a lower first ionisation energy than phosphorus.

Still don't get it? · why sulfur's first ionisation energy dips below phosphorus

Picture a row of empty two-seater benches: these are the three 3p orbitals. Phosphorus has three electrons to seat, so each one gets a bench to itself, nobody crowded. Sulfur has a fourth electron, and there are no empty benches left, so it has to sit down next to someone who is already there. Two people squeezed onto one bench press against each other, and it is easy to nudge one of them off the end.

Now the chemistry. In phosphorus (3p³) each 3p orbital has one electron in it. In sulfur (3p⁴) one orbital holds two electrons, a pair. Both electrons are negative, so the two paired electrons repel each other, and that repulsion makes one of them a little easier to pull away.

So the first electron removed from sulfur comes from a 3p orbital that holds a pair of electrons, those paired electrons repel, less energy is needed, and sulfur's first ionisation energy is lower than phosphorus's. The mark needs the words "paired electrons in a 3p orbital" and "repel". Just "electrons repel", or "sulfur has a stable half-filled shell", earns nothing.

Trend in melting point (Na to Ar)

The skill here is always the same: first decide the structure and bonding of the element, then decide what has to be broken to melt it.

Working across the row:

• Na, Mg, Al (metals): melting point rises. Going Na to Mg to Al the ions have a greater positive charge, are smaller, and release more electrons into the sea of delocalised electrons (1, then 2, then 3). So the attraction between the positive ions and the delocalised electrons is stronger: stronger metallic bonding, higher melting point.
• Si is macromolecular (giant covalent). Melting it means breaking many strong covalent bonds, which takes a great deal of energy, so silicon has the highest melting point in the period.
• P₄, S₈ and Cl₂ are simple molecular. Only the weak van der Waals forces between the molecules are broken; the strong covalent bonds inside each molecule stay intact. So their melting points are low.
• S₈ has a higher melting point than P₄ because the S₈ molecule is larger and has more electrons than P₄, so the van der Waals forces between S₈ molecules are stronger.
• Ar exists as single atoms with only very weak van der Waals forces between them, so it has the lowest melting point.

Still don't get it? · melting point: work out what you are actually breaking

Think of a bar of chocolate made of little squares. You can snap it into separate squares without damaging any single square, or you can crush a square itself. Those are two completely different amounts of effort. Melting is like separating things, and the question is always: what are the separate "things", and what holds them to each other?

To melt an element you pull its particles apart from one another. You do not break up the particles themselves.

• In a metal (Na, Mg, Al) the particles are positive ions sitting in a sea of delocalised electrons, so melting weakens the metallic bonding between them.
• In silicon the whole crystal is one giant network of atoms joined by covalent bonds. There are no separate molecules to pull apart, so the only thing you can break is the covalent bonds themselves, and there are many strong ones. That is why silicon's melting point is so high.
• In P₄, S₈ and Cl₂ the particles are small molecules. The covalent bonds inside each molecule are strong, but to melt you only separate whole molecules from each other, which means breaking the weak van der Waals forces between them, not the covalent bonds.

So for the molecular elements the mark is "weak van der Waals forces between molecules are broken", never "covalent bonds are broken", and for silicon it is "many strong covalent bonds are broken". Mixing these up, such as "van der Waals forces between the atoms of a metal" or "the covalent bonds in sulfur are broken", scores zero.

Worked examples

These are model explain answers. The lesson is how many separate points an answer needs, and in what order, so each numbered point below is worth one mark.

1. Define the first ionisation energy of magnesium and write an equation for it.

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

Mg(g) ⟶ Mg⁺(g) + e^- (1 for the equation, 1 for both state symbols)

The marks are won on the exact conditions: keep "one mole" in both halves and show both species as (g).

2. Explain why the first ionisation energy of aluminium is lower than that of magnesium. (3 marks)

1. Aluminium's outer electron is removed from a 3p sub-shell, whereas magnesium's is removed from a 3s sub-shell.
2. The 3p sub-shell is higher in energy than the 3s (and is slightly shielded by the 3s electrons).
3. So less energy is needed to remove aluminium's outer electron, giving aluminium the lower first ionisation energy.

Where the mark is lost: writing only that the electron is "easier to remove" without naming the 3p sub-shell, or without stating that aluminium's value is lower.

3. Explain why the first ionisation energy of sulfur is lower than that of phosphorus. (2 marks)

1. In sulfur the first electron is removed from a 3p orbital that contains a pair of electrons (in phosphorus every 3p orbital has just one electron).
2. The paired electrons repel each other, so less energy is needed to remove one.

Where the mark is lost: saying "electrons repel" without specifying the paired electrons in a 3p orbital.

4. Explain why aluminium has a higher melting point than sodium. (3 marks)

1. Both are giant metallic structures (positive ions in a sea of delocalised electrons).
2. Aluminium has a higher ionic charge, a smaller ion and more delocalised electrons (3 per atom, against 1 for sodium).
3. So the attraction between the positive ions and the delocalised electrons is stronger (stronger metallic bonding), needing more energy to melt.

Where the mark is lost: comparing ionisation energies instead of melting points, or naming the wrong force (ionic, intermolecular or van der Waals) for a metal.

5. Explain why silicon has the highest melting point in Period 3. (2 marks)

1. Silicon is macromolecular (giant covalent).
2. Melting it requires many strong covalent bonds to be broken, which needs a great deal of energy.

Where the mark is lost: saying covalent bonds are "weakened" or "loosened" rather than broken, or describing silicon as molecular or ionic.

6. Explain why sulfur has a higher melting point than phosphorus. (2 marks)

1. Both are simple molecular, so melting breaks the van der Waals forces between molecules.
2. S₈ is a larger molecule with more electrons than P₄, so the van der Waals forces are stronger and more energy is needed.

Where the mark is lost: saying the covalent bonds in sulfur are broken, or that "the sulfur atom is bigger", instead of comparing the molecules and their van der Waals forces.

Common exam mistakes

• Thinking periodicity means a trend down a group. It is a repeating trend across a period (Q15, Q21).
• Naming a block from the number of outer electrons ("six outer electrons, so block 6") or calling it the "p-shell". The block is named after the sub-shell (sub-level) holding the highest energy electron, so use "sub-level" or "sub-shell", not "shell" (Q34, Q3).
• In the first ionisation energy definition, mixing up "mole" and "atom" (keep one mole in both the electron and the atom part, do not mix and match) and leaving out the gaseous state. Both lose the mark (Q17, Q18, Q22, Q26).
• In the equation, omitting state symbols (both species must be (g)), putting the electron on the wrong side, or showing two electrons removed. The state symbols carry their own mark (Q11, Q22).
• Explaining the ionisation-energy trend with "more neutrons" (the nuclear charge rises because of more protons), dropping the word "outer" from "outer electron", or leaving out "similar shielding" for the final mark. "A bigger nucleus" is not the same as "a greater nuclear charge" (Q18, Q19, Q22).
• For the Mg to Al dip, writing only that aluminium's electron is "easier to remove" without stating its first ionisation energy is lower and that the electron comes from a 3p sub-shell (Q2, Q18).
• For the P to S dip, saying "electrons repel" without specifying the paired electrons in a 3p orbital, confusing it with the lone-pair repulsion used for molecular shapes, or claiming a "stable half-filled shell". None of these score (Q1, Q6, Q22, Q26, Q33).
• For the melting point of the metals, blaming "ionic bonding", "intermolecular forces", "van der Waals forces" or even "hydrogen bonding" (all wrong for a metal), or answering with a comparison of ionisation energies instead of melting point. The mark needs a stronger attraction between the positive ions and the delocalised electrons (Q2, Q6, Q39).
• For silicon, saying covalent bonds are "loosened" or "weakened" rather than broken, or calling silicon ionic or molecular. You must say many strong covalent bonds are broken (Q9, Q18, Q20).
• For S₈ against P₄, attributing sulfur's higher melting point to breaking covalent bonds, or to "the sulfur atom being bigger". It is because the S₈ molecule is larger with more electrons, giving stronger van der Waals forces between molecules (Q9, Q20).
• Picking the wrong element for the extremes: the highest first ionisation energy in Period 3 is argon, not sulfur, and the highest melting point is silicon, not argon or aluminium (Q4, Q12, Q19).

Key definitions

• First ionisation energy - the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions. For element X: X(g) ⟶ X⁺(g) + e^-.
• Periodicity - a regular, repeating pattern (trend) in the properties of elements across a period as the proton number increases.
• s, p, d or f block - an element is in the block named after the sub-shell (sub-level) that holds its highest energy electron, which is fixed by its proton number.
• Metallic bonding - the electrostatic attraction between positive metal ions and the delocalised electrons (the force broken when a metal melts).

Specification

• I can classify an element as s, p, d or f block from its position or electron configuration, and relate this to its proton number.
• I can state and explain the trend in atomic radius across Period 3 in terms of nuclear charge and shielding.
• I can define first ionisation energy and write an equation for it with state symbols.
• I can state and explain the general trend in first ionisation energy across Period 3.
• I can explain the drop in first ionisation energy from Mg to Al, and from P to S.
• I can explain the trend in melting point across Period 3 in terms of the structure of, and bonding in, the elements.

Related notes

• Atomic structure
• Bonding
• Group 2, the alkaline earth metals
• Group 7, the halogens
• Properties of Period 3 elements and their oxides (A-level)