Periodicity

 How is the modern Periodic Table arranged?

By increasing atomic number (proton number).

What is a period?

A row of the Periodic Table. The period number tells you the number of occupied electron shells.

What is a group?

A column of the Periodic Table. Elements in the same group have the same number of outer-shell electrons, so they have similar chemical properties.

Why do elements in the same group have similar chemical properties?

They have the same number of outer-shell/valence electrons, so they have similar electron configurations involved in chemical reactions.

What are the four blocks of the Periodic Table?

s-block, p-block, d-block and f-block.

What determines which block an element belongs to?

The subshell being filled in its electron configuration.

Which elements are in the s-block?

Groups 1 and 2, with helium also classified as s-block.

Which elements are in the p-block?

Groups 13–18.

Which elements are in the d-block?

The transition metals in the central section of the Periodic Table.

Where is the f-block found?

The lanthanides and actinides, shown separately underneath the main Periodic Table.

Give an example of an s-block element and its electron configuration.

N

Give an example of a p-block element and its electron configuration.

O: 1s² 2s² 2p⁴. The p subshell is being filled, so O is p-block.

Give an example of a d-block element and its electron configuration.

Fe: [Ar] 4s² 3d⁶. The d subshell is being filled, so Fe is d-block.

Give an example of an f-block element.

Uranium (U) is an example of an f-block element.

How can you determine the electron configuration of an element from its position in the Periodic Table?

1. The period tells you the number of occupied electron shells.
2. The block tells you which type of subshell is being filled.
3. Build the electron configuration using the correct filling order.

What is the electron configuration of vanadium (V)?

V is in Period 4 and the d-block:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d³
or [Ar] 4s² 3d³.

What is atomic radius?

A measure of the distance between the nucleus and the outer occupied electron shell.

What happens to atomic radius across Period 3 from Na → Ar?

It decreases.

Why does atomic radius decrease across Period 3?

Proton number increases → nuclear charge increases. Electrons are added to the same principal energy level, so there is little additional shielding. The stronger attraction between the nucleus and outer electrons pulls them closer to the nucleus, so atomic radius decreases.

What is the full-mark explanation for the decrease in atomic radius across Period 3?

Increasing nuclear charge + similar shielding → stronger electrostatic attraction between the nucleus and outer electrons → outer electrons are pulled closer → smaller atomic radius.

Why doesn't shielding increase significantly across Period 3?

The added electrons are placed in the same outer shell, rather than a new inner shell.

What happens to nuclear charge across Period 3?

It increases because each successive element has an additional proton.

What is the overall trend in melting points across Period 3?

Melting point increases from Na → Si, reaches a maximum at Si, then decreases dramatically from Si → Ar.

What type of structure do Na, Mg and Al have?

A giant metallic lattice.

Why does melting point increase from Na → Mg → Al?

Metallic bonding becomes stronger because:
positive ion charge increases: Na⁺ → Mg²⁺ → Al³⁺
ionic radius decreases
the number of delocalised electrons per atom increases: 1 → 2 → 3
therefore electrostatic attraction between positive metal ions and delocalised electrons becomes stronger.
More energy is needed to overcome the metallic bonding.

Why does Na have a lower melting point than Mg and Al?

Na forms Na⁺ ions, has one delocalised electron per atom and has a larger ionic radius. Mg and Al have greater positive charge, smaller ions and more delocalised electrons, giving stronger metallic bonding.

Why does silicon have such a high melting point?

Silicon has a giant covalent structure. Each Si atom is covalently bonded to four other Si atoms in a giant 3D lattice. Many strong covalent bonds must be overcome, requiring a very large amount of energy.

Why is there a large decrease in melting point from Si → P?

Si has a giant covalent structure, whereas phosphorus exists as P₄ simple molecular molecules. Melting P₄ only requires overcoming weak intermolecular forces, rather than strong covalent bonds throughout a giant lattice.

What structure does phosphorus have at room temperature?

P₄ simple molecular molecules.

What structure does sulfur have at room temperature?

S₈ simple molecular molecules.

What structure does chlorine have at room temperature?

Cl₂ simple molecular molecules.

What structure does argon have?

Argon is monatomic — it exists as individual Ar atoms.

What forces act between P₄, S₈ and Cl₂ molecules?

Induced dipole–dipole forces (London dispersion forces).

Why does sulfur have a higher melting point than phosphorus?

S₈ molecules are larger and contain more electrons than P₄ molecules. They therefore have stronger induced dipole–dipole forces, requiring more energy to overcome.

Why does chlorine have a lower melting point than sulfur?

Cl₂ molecules are smaller and contain fewer electrons than S₈, so they have weaker induced dipole–dipole forces.

Why does argon have such a low melting point?

Argon is monatomic, so only very weak induced dipole–dipole forces act between individual Ar atoms. Very little energy is needed to overcome these forces.

What is the key difference between melting a giant covalent substance and melting a simple molecular substance?

Giant covalent: strong covalent bonds throughout the structure must be overcome.
Simple molecular: only intermolecular forces between molecules are overcome; covalent bonds within the molecules remain intact.

What is ionisation?

The removal of one or more electrons from an atom or ion.

What is 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.

Write the general equation for first ionisation energy.

X(g) → X⁺(g) + e⁻

Give the first ionisation equation for magnesium.

Mg(g) → Mg⁺(g) + e⁻

Why are ionisation energies always positive?

Energy must be supplied to overcome the electrostatic attraction between the nucleus and the electron being removed.

What three factors affect ionisation energy?

Nuclear charge, atomic radius and electron shielding.

How does increasing nuclear charge affect ionisation energy?

More protons → stronger electrostatic attraction between the nucleus and the electron → higher ionisation energy.

How does increasing atomic radius affect ionisation energy?

A larger radius means the outer electron is further from the nucleus → weaker attraction → lower ionisation energy.

How does increasing electron shielding affect ionisation energy?

Inner electrons shield the outer electron from the full nuclear charge → weaker attraction → lower ionisation energy.

What is the general trend in first ionisation energy across Period 3?

It generally increases from Na → Ar.

Why does first ionisation energy generally increase across Period 3?

Nuclear charge increases, electrons are added to the same principal energy level, shielding remains similar and atomic radius decreases. Therefore, attraction between the nucleus and outer electron increases, so more energy is required to remove the electron.

What is the trend in first ionisation energy down a group?

It decreases.

Why does first ionisation energy decrease down a group?

Nuclear charge increases, but atomic radius and electron shielding also increase. The increased radius and shielding outweigh the increased nuclear charge, so attraction to the outer electron decreases and less energy is needed to remove it.

What are the two important dips in first ionisation energy across Period 3?

Mg → Al and P → S.

Why is the first ionisation energy of Al lower than Mg?

Mg: [Ne] 3s²
Al: [Ne] 3s² 3p¹
The electron removed from Al is in a 3p subshell, whereas Mg loses a 3s electron. The 3p electron is higher in energy and experiences weaker attraction, so it is easier to remove.
Therefore, Al has a lower first ionisation energy than Mg.

What are the first ionisation energies of Mg and Al?

Mg ≈ 738 kJ mol⁻¹
Al ≈ 578 kJ mol⁻¹

Why is the first ionisation energy of S lower than P?

P: [Ne] 3s² 3p³
S: [Ne] 3s² 3p⁴
In P, the 3p electrons occupy separate orbitals. In S, one 3p orbital contains a pair of electrons, causing greater electron–electron repulsion. One of the paired electrons is therefore easier to remove.

What are the first ionisation energies of P and S?

P ≈ 1011 kJ mol⁻¹
S ≈ 999 kJ mol⁻¹

What is the full-mark explanation for the general increase in first ionisation energy across Period 3?

Nuclear charge increases while shielding remains similar → stronger electrostatic attraction between the nucleus and outer electron → atomic radius decreases → more energy is required to remove the outer electron.

What is a successive ionisation energy?

The energy required to remove an electron from a successively more positively charged gaseous ion.

Write the equation for second ionisation energy.

X⁺(g) → X²⁺(g) + e⁻

Write the second ionisation equation for magnesium.

Mg⁺(g) → Mg²⁺(g) + e⁻

Why do successive ionisation energies increase?

After an electron is removed, the ion becomes more positively charged and there is less electron–electron repulsion. The remaining electrons experience greater electrostatic attraction to the nucleus, so more energy is needed to remove each successive electron.

Why is Mg's second ionisation energy much higher than its first?

Both electrons are removed from the 3s subshell, but after the first electron is removed, the Mg⁺ ion is more positively charged. The remaining electron is therefore more strongly attracted to the nucleus.

What does a large jump between successive ionisation energies indicate?

It indicates that an electron is being removed from an inner shell, closer to the nucleus and with less shielding.

Where is the large jump in Mg's successive ionisation energies?

Between IE₂ and IE₃.

Why is there a large jump between Mg's second and third ionisation energies?

Mg has electron configuration [Ne] 3s².
IE₁ removes the first 3s electron and IE₂ removes the second 3s electron. IE₃ must remove an electron from the inner 2p shell. This electron is much closer to the nucleus and experiences much stronger attraction, so much more energy is required.

How can successive ionisation energies provide evidence for electron shells?

A large jump occurs when electron removal moves from an outer shell to an inner shell. The position of the jump therefore indicates how many electrons were present in the outer shell.

What are the three Period 3 trends you need to know?

Atomic radius, first ionisation energy and melting point.

What happens to atomic radius from Na → Ar?

Decreases.

What happens to first ionisation energy from Na → Ar?

Generally increases, with dips at Al and S.

What happens to melting point from Na → Ar?

Increases Na → Si, then decreases sharply Si → Ar.

What are the three key structures across Period 3?

Na–Al: giant metallic
Si: giant covalent
P–Ar: simple molecular/monatomic

What is the key explanation for the Na → Mg → Al melting-point trend?

Increasing positive charge + decreasing ionic radius + more delocalised electrons → stronger metallic bonding.

What is the key explanation for the Si melting point?

Giant covalent structure → many strong covalent bonds → large amount of energy needed to break them.

What is the key explanation for P → Ar melting points?

Simple molecular/monatomic substances have weak induced dipole–dipole forces, so little energy is needed to overcome them.

What are the three key electron factors to remember for ionisation-energy explanations?

Nuclear charge, atomic radius and electron shielding.

What is the most important exam technique for trend questions?

Don't just state the trend. Give the full cause → effect chain.
For example:
Increasing nuclear charge → stronger attraction → smaller atomic radius → higher first ionisation energy.

Can you explain why atomic radius decreases across Period 3?

Increasing nuclear charge + similar shielding → stronger attraction → electrons pulled closer → smaller radius.

Can you explain why first ionisation energy generally increases across Period 3?

Increasing nuclear charge + similar shielding → stronger attraction + decreasing radius → more energy needed to remove an electron.

Can you explain the Mg → Al ionisation-energy dip?

Al loses a 3p electron, which is higher in energy than Mg's 3s electron → easier to remove → lower IE.

Can you explain the P → S ionisation-energy dip?

S has a paired 3p electron → greater electron–electron repulsion → easier to remove → lower IE.

Can you explain the Na → Mg → Al melting-point increase?

Increasing positive ion charge + smaller ionic radius + more delocalised electrons → stronger metallic bonding → more energy required to melt.

Can you explain why Si has the highest melting point?

Si has a giant covalent lattice with strong covalent bonds throughout → large amount of energy required to break bonds.

Can you explain why the melting point drops dramatically from Si → P?

Changes from giant covalent Si to simple molecular P₄ → only weak intermolecular forces need to be overcome.

Can you explain why S₈ has a higher melting point than P₄?

S₈ is larger and has more electrons → stronger induced dipole–dipole forces → more energy needed to overcome them.

Can you explain why Ar has such a low melting point?

Ar is monatomic and only has very weak induced dipole–dipole forces between atoms.

Can you explain a large jump in successive ionisation energies?

The next electron is being removed from an inner shell → closer to nucleus + less shielding → much stronger attraction → much more energy required