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The Periodic Table

1

Organisation of the Periodic Table

s blockp blockd blockf blockPeriod 3Group 17

Periodicity: the repeating pattern of physical and chemical properties at regular intervals, with elements in order of increasing proton number
Blocks are named after the subshell holding the valence electrons

Group (column): same outer electronic configuration → similar chemical properties. Main group: group number = no. of valence electrons
Period (row): period number = principal quantum number, n, of the valence shell

2

Atomic and Ionic Radius

2.1Atomic radius

Atomic radius: half the distance between the nuclei of two adjacent atoms of the same element.

Across Period 3, atomic radius ↓

Nuclear charge ↑, electrons added to the same shell → shielding approx. constant
→ stronger attraction between nucleus and valence electrons

Down a group, atomic radius ↑

Number of filled principal quantum shells ↑
→ valence electrons further from nucleus and more shielded → weaker attraction

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2

Atomic and Ionic Radius (continued)

2.2Ionic radius

Cation is smaller than its atom

Atoms lose some or all of their valence electrons to form cations
→ cations usually have one less principal quantum shell of electrons

Anion is larger than its atom

Electrons are added to the outermost principal quantum shell to form anions
→ greater electron–electron repulsion in the outermost shell

Isoelectronic ions: Na⁺ → Si⁴⁺ and P³⁻ → Cl⁻

Isoelectronic ions have the same number of electrons
Atomic number ↑ → nuclear charge ↑ → ionic radius ↓

NoteP³⁻ ≫ Si⁴⁺: P³⁻ has one more filled principal quantum shell.

3

Electronegativity

DefinitionElectronegativity: the ability of an atom to attract the shared pair of electrons in a covalent bond.

Across a period ↑: nuclear charge ↑, shielding approx. constant
Down a group ↓: more shells, bonding pair further and more shielded
Across Period 3, the electronegativity difference with O or Cl ↓ → oxides and chlorides change from ionic to covalent

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4

First Ionisation Energy

General trend: first IE ↑ across Period 3

Nuclear charge ↑, shielding approx. constant (same shell)
→ more energy to remove the more strongly attracted valence electron

Down a group: IE ↓ (more shells, more shielding)

Exceptions: the two dips

Mg → Al: 3p e⁻ removed is higher in energy than 3s
P → S: Coulombic repulsion between the paired 3p e⁻ in S → one of them is easier to remove than an unpaired 3p e⁻ from P

5

Melting Point across Period 3

Element	Structure	m.p.
Na, Mg, Al	giant metallic	high, ↑ across
Si	giant molecular	highest
P₄, S₈, Cl₂	simple molecular	low
Ar	monatomic	lowest

Na < Mg < AlMore valence e⁻ delocalised per atom and smaller, higher-charged cation → stronger metallic bonds

SiA very large amount of energy is needed to break the extensive network of strong covalent bonds

S₈ > P₄ > Cl₂ > ArNo. of electrons in S₈ > P₄ > Cl₂ > Ar → more energy is needed to overcome the stronger id–id between S₈ molecules than between P₄, then Cl₂ molecules, then Ar atoms

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6

Electrical Conductivity across Period 3

Na → Al: giant metallic lattice, high conductivity from mobile delocalised electrons
Conductivity ↑ Na → Al, as more valence electrons per atom are delocalised
Si: sharp drop, giant molecular semiconductor, conducts more as T ↑
P₄ → Ar: simple molecular or atoms, essentially zero, as there are no mobile charge carriers

7

Group 2 Elements

7.1Reducing power

Reducing power: the ability of an element to lose electrons and reduce another species. Group 2 metals react as M → M²⁺ + 2e⁻.

Why does reducing power ↑ down the group?

Down the group, the number of principal quantum shells ↑ → atomic radius ↑
Shielding effect ↑, which largely cancels out the increase in nuclear charge
→ valence e⁻ less strongly attracted by the nucleus → less energy is needed to remove them
∴ M loses its 2 valence e⁻ more readily → reducing power ↑
E⦵(M²⁺/M) more negative down the group: Mg −2.38, Ca −2.87, Ba −2.90 V

Examples of differing reducing power: reaction with water

Mg: reacts very slowly with cold water, but readily with steam: Mg + H₂O(g) → MgO + H₂
Ca, Sr, Ba: react with cold water, increasingly vigorously down the group: M + 2H₂O → M(OH)₂ + H₂. Solutions become more alkaline

NoteGroup 2 elements only form the +2 state: removing a third e⁻ means removing an e⁻ from an inner shell.

7.2Thermal stability of carbonates

Thermal stability: the resistance of a compound to decomposition on heating. The thermal stability of Group 2 carbonates increases down the group.

MCO₃(s) → MO(s) + CO₂(g)

Why does thermal stability ↑ down the group?

Down the group, the radius of the metal cation ↑ and its charge density ↓
→ the ability of M²⁺ to polarise the electron cloud of the large CO₃²⁻ anion ↓, and the C–O bonds are weakened to a smaller extent
∴ more energy is needed to decompose MCO₃ → thermal stability ↑

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8

Oxides of Period 3

8.1Structure, bonding and nature

Oxide	Ox. no.	Structure and bonding	Nature	pH in water
Na₂O	+1	giant ionic	basic	13
MgO	+2	giant ionic	basic	9
Al₂O₃	+3	giant ionic, covalent character	amphoteric	7 (insoluble)
SiO₂	+4	giant molecular	acidic	7 (insoluble)
P₄O₁₀	+5	simple molecular	acidic	2
SO₃	+6	simple molecular	acidic	1

8.2Equations, in order of the elements

Oxide	With water	With acid	With alkali
Na₂O	Na₂O + H₂O → 2NaOH	Na₂O + 2HCl → 2NaCl + H₂O	–
MgO	MgO + H₂O ⇌ Mg(OH)₂	MgO + 2HCl → MgCl₂ + H₂O	–
Al₂O₃	no reaction	Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O	Al₂O₃ + 2OH⁻ + 3H₂O → 2[Al(OH)₄]⁻
SiO₂	no reaction	–	SiO₂ + 2OH⁻ → SiO₃²⁻ + H₂O hot, concentrated NaOH
P₄O₁₀	P₄O₁₀ + 6H₂O → 4H₃PO₄	–	P₄O₁₀ + 12OH⁻ → 4PO₄³⁻ + 6H₂O
SO₃	SO₃ + H₂O → H₂SO₄	–	SO₃ + 2OH⁻ → SO₄²⁻ + H₂O

Exam tipReactions with water and alkali are not redox: the oxidation number stays the same, so P₄O₁₀ (+5) gives H₃PO₄ (+5), not H₃PO₃.

8.3Why is Al₂O₃ amphoteric?

Al³⁺ is small and highly charged → high charge density
→ polarises the electron cloud of O²⁻ → partial covalent character
∴ reacts with both acids and bases

Al(OH)₃ is amphoteric too: Al(OH)₃ + 3H⁺ → Al³⁺ + 3H₂O and Al(OH)₃ + OH⁻ → [Al(OH)₄]⁻. NaOH and Mg(OH)₂ are basic only.

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8

Oxides of Period 3 (continued)

8.4Why are Al₂O₃ and SiO₂ insoluble?

Al₂O₃: very high lattice energy
SiO₂: extensive network of strong covalent bonds
Not enough energy is released on hydration to overcome these, so pH stays 7

9

Chlorides of Period 3

9.1Structure and behaviour in water

Chloride	Ox. no.	Structure	With water	pH
NaCl	+1	giant ionic	hydration only	7
MgCl₂	+2	giant ionic	hydration, slight hydrolysis	6.5
AlCl₃ (Al₂Cl₆)	+3	simple molecular	hydration, appreciable hydrolysis	3
SiCl₄	+4	simple molecular	complete hydrolysis	2
PCl₅	+5	simple molecular	complete hydrolysis	2

Highest oxidation number = no. of valence e⁻, as every valence e⁻ is used in bonding to the more electronegative O or Cl: +1 (Na) to +6 (S in SO₃).

9.2Equations, in order of the elements

Chloride	Excess water	Limited water
NaCl	NaCl(s) → Na⁺(aq) + Cl⁻(aq)	–
MgCl₂	MgCl₂ + 6H₂O → [Mg(H₂O)₆]²⁺ + 2Cl⁻ [Mg(H₂O)₆]²⁺ ⇌ [Mg(H₂O)₅(OH)]⁺ + H⁺	–
AlCl₃	AlCl₃ + 6H₂O → [Al(H₂O)₆]³⁺ + 3Cl⁻ [Al(H₂O)₆]³⁺ ⇌ [Al(H₂O)₅(OH)]²⁺ + H⁺	AlCl₃ + 3H₂O → Al(OH)₃ + 3HCl white solid, white fumes
SiCl₄	SiCl₄ + 2H₂O → SiO₂ + 4HCl	–
PCl₅	PCl₅ + 4H₂O → H₃PO₄ + 5HCl	PCl₅ + H₂O → POCl₃ + 2HCl

9.3Why is AlCl₃(aq) acidic?

Al³⁺ has a high charge density
→ strongly polarises the O–H bonds of the water ligands
→ O–H weakens and breaks, releasing H⁺

9.4Why does SiCl₄ hydrolyse but not CCl₄?

Si has vacant, energetically accessible 3d orbitals to accept a lone pair from H₂O
C has no 3d orbitals, and the small C atom is shielded by four large Cl atoms

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10

Group 17: The Halogens

10.1Physical properties

	At r.t.p.	In hexane	In water	X–X / kJ mol⁻¹
F₂	pale yellow gas	very pale yellow	–	158
Cl₂	greenish-yellow gas	pale yellow	pale yellow	244
Br₂	red-brown liquid	orange-red	orange	193
I₂	black solid	purple	brown	151

NoteF–F is weaker than Cl–Cl: the lone pairs on the two small, close F atoms repel each other.

10.2Volatility

Volatility: the ease with which a substance changes into a gas. The lower the b.p., the more volatile the substance.

Why does volatility ↓ down the group?

All halogens are simple molecular, non-polar X₂ molecules
Down the group, M_r ↑ → no. of electrons ↑ → electron cloud more polarisable → stronger id–id
→ more energy is needed to overcome the id–id between X₂ molecules → b.p. ↑ → volatility ↓

10.3Oxidising power

Oxidising power: the ability of an element to gain electrons and oxidise another species. Halogens react as X₂ + 2e⁻ → 2X⁻.

Why does oxidising power ↓ down the group?

Down the group, the number of principal quantum shells ↑ → atomic radius ↑
Shielding effect ↑, which largely cancels out the increase in nuclear charge
→ the incoming e⁻ is less strongly attracted by the nucleus → X₂ gains electrons less readily
∴ oxidising power ↓: F₂ > Cl₂ > Br₂ > I₂
E⦵(X₂/X⁻) less positive down the group: F₂ +2.87, Cl₂ +1.36, Br₂ +1.07, I₂ +0.54 V

10.4Displacement reactions: each halogen with the other halide ions

A halogen oxidises the halide ions of the halogens below it. Observations are for aqueous solutions, shaking with hexane shows the halogen formed more clearly.

Halogen	Halide	Reaction	Observation
F₂	Cl⁻, Br⁻, I⁻	oxidises all of them	F₂ also oxidises water, so it is not used in aqueous solution
Cl₂	Br⁻	Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂	colourless solution turns orange, orange-red in hexane
Cl₂	I⁻	Cl₂ + 2I⁻ → 2Cl⁻ + I₂	colourless solution turns brown, purple in hexane
Br₂	Cl⁻	no reaction	solution stays orange (from Br₂)
Br₂	I⁻	Br₂ + 2I⁻ → 2Br⁻ + I₂	orange solution turns brown, purple in hexane
I₂	Cl⁻, Br⁻	no reaction	solution stays brown (from I₂)

76Topic 05 · The Periodic Table

11

Hydrogen Halides

11.1Boiling points

HCl → HI: more electrons → stronger id–id → b.p. ↑
HF: anomalously high, due to hydrogen bonding

11.2Thermal stability

HX	H–X / kJ mol⁻¹	On heating
HF	562	stable
HCl	431	stable
HBr	366	some brown Br₂ fumes
HI	299	purple I₂ fumes with a hot rod

Why does thermal stability decrease down the group?

Valence orbital of X is larger and more diffuse → less effective overlap with H 1s
→ H–X bond weaker → less energy needed to break it

12

Deducing and Predicting

12.1Identify element X

X is a solid with very low conductivity at r.t. Its only chloride is a liquid non-conductor, which hydrolyses in water to give a white solid and a strongly acidic solution.

Answer structureSet out each piece of information beside what you deduce from it, then conclude.

Information	Deduction
solid, very low conductivity	semiconductor
chloride is a liquid non-conductor	simple molecular chloride
white solid + strongly acidic solution	SiCl₄ + 2H₂O → SiO₂ + 4HCl

∴ X is silicon

12.2Predict: astatine, below iodine

Prediction	Reason
less electronegative than I	more shells, more shielding
dark solid, less volatile	more electrons, stronger id–id
HAt decomposes more easily	H–At bond weaker than H–I
cannot oxidise Cl⁻ to Cl₂	weaker oxidising agent than I₂

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