The Periodic
Table

H2 Chemistry (9476), JC1

By the end of this topic, you can answer

1

Why is a sodium ion smaller than a sodium atom, while a chloride ion is larger than a chlorine atom?

2

Silicon and phosphorus are neighbours, so why does silicon melt at about 1410 °C and white phosphorus at only 44 °C?

3

Why does barium react with cold water much faster than magnesium, and why is barium carbonate harder to decompose by heating?

4

Why does sodium oxide make water alkaline, while phosphorus(V) oxide makes it acidic? And why does PCl5 fume in water while NaCl simply dissolves?

5

A hot wire plunged into hydrogen iodide gives purple fumes, but nothing happens with hydrogen chloride. Why? And how would astatine behave?

Recall: what you know already

1From Atomic Structure: why does first IE generally increase across a period?

Nuclear charge increases, shielding is about constant, so the valence electron is more strongly attracted.[1]

2From Chemical Bonding: name the structure of Na, Si and Cl₂.

Giant metallic, giant molecular, simple molecular.[1]

3From O levels: what colour is iodine dissolved in hexane?

Purple.[1]
1

Organisation of the Periodic Table

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

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

Down a group, atomic radius ↑

  1. Number of filled principal quantum shells ↑
  2. → valence electrons further from nucleus and more shielded → weaker attraction
2

Atomic and Ionic Radius

2.2Ionic radius

Cation is smaller than its atom

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

Anion is larger than its atom

  1. Electrons are added to the outermost principal quantum shell to form anions
  2. → greater electron–electron repulsion in the outermost shell
2

Atomic and Ionic Radius

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

  1. Isoelectronic ions have the same number of electrons
  2. Atomic number ↑ → nuclear charge ↑ → ionic radius ↓
NoteP³⁻ ≫ Si⁴⁺: P³⁻ has one more filled principal quantum shell.

Your turnArrange in order of increasing radius: Mg²⁺, Na⁺, F⁻, O²⁻. Explain.

Mg²⁺ < Na⁺ < F⁻ < O²⁻. All have 10 electrons (isoelectronic).[1]
Nuclear charge decreases from Mg (12) to O (8), so the electrons are less strongly attracted and the radius increases.[1]
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

3.1Diagonal relationship

Group 1Group 2Group 13
Period 2Li 0.98Be 1.57B 2.04
Period 3Na 0.93Mg 1.31Al 1.61

Similar electronegativity diagonally → similar properties: Li₂CO₃ decomposes like MgCO₃, Be(OH)₂ is amphoteric like Al(OH)₃.

4

First Ionisation Energy

First IE generally ↑ across Period 3

  1. Nuclear charge ↑, shielding approx. constant (same shell)
  2. → more energy to remove the more strongly attracted valence electron
  • 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
  • Down a group: IE ↓ (more shells, more shielding)
5

Melting Point across Period 3

ElementStructurem.p.
Na, Mg, Algiant metallichigh, ↑ across
Sigiant molecularhighest
P₄, S₈, Cl₂simple molecularlow
Armonatomiclowest
5

Melting Point across Period 3

  • Na < Mg < Al: more valence e⁻ delocalised per atom and smaller, higher-charged cation → stronger metallic bonds
  • Si: a very large amount of energy is needed to break the extensive network of strong covalent bonds
  • S₈ > P₄ > Cl₂ > Ar: no. 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
6

Electrical Conductivity across Period 3

  1. Na → Al: giant metallic lattice, high conductivity from mobile delocalised electrons
  2. Conductivity ↑ Na → Al, as more valence electrons per atom are delocalised
  3. Si: sharp drop, giant molecular semiconductor, conducts more as T ↑
  4. P₄ → Ar: simple molecular or atoms, essentially zero, as there are no mobile charge carriers
6

Electrical Conductivity across Period 3

Quick checkWhich Period 3 element has the highest melting point?

AAl
BSi
CP
DS
NoteB: Si is giant molecular, so melting breaks an extensive network of strong covalent bonds.
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?

  1. Down the group, the number of principal quantum shells ↑ → atomic radius ↑
  2. Shielding effect ↑, which largely cancels out the increase in nuclear charge
  3. → valence e⁻ less strongly attracted by the nucleus → less energy is needed to remove them
  4. ∴ M loses its 2 valence e⁻ more readily → reducing power ↑
7

Group 2 Elements

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.

PredictHow would barium react with cold water compared with calcium?

More vigorously: Ba is a stronger reducing agent, so it loses its valence electrons more readily.[1]
Ba + 2H₂O → Ba(OH)₂ + H₂[1]
7

Group 2 Elements

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)

  1. Down the group, the radius of the metal cation ↑ and its charge density ↓
  2. → 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
  3. ∴ more energy is needed to decompose MCO₃ → thermal stability ↑
8

Oxides of Period 3

8.1Structure, bonding and nature

OxideOx. no.Structure and bondingNaturepH in water
Na₂O+1giant ionicbasic13
MgO+2giant ionicbasic9
Al₂O₃+3giant ionic, covalent characteramphoteric7 (insoluble)
SiO₂+4giant molecularacidic7 (insoluble)
P₄O₁₀+5simple molecularacidic2
SO₃+6simple molecularacidic1
8

Oxides of Period 3

8.2Equations, in order of the elements

OxideWith waterWith acidWith alkali
Na₂ONa₂O + H₂O → 2NaOHNa₂O + 2HCl → 2NaCl + H₂O–
MgOMgO + H₂O ⇌ Mg(OH)₂MgO + 2HCl → MgCl₂ + H₂O–
Al₂O₃no reactionAl₂O₃ + 6HCl → 2AlCl₃ + 3H₂OAl₂O₃ + 2OH⁻ + 3H₂O → 2[Al(OH)₄]⁻
SiO₂no reaction–SiO₂ + 2OH⁻ → SiO₃²⁻ + H₂O
hot, concentrated NaOH
P₄O₁₀P₄O₁₀ + 6H₂O → 4H₃PO₄–P₄O₁₀ + 12OH⁻ → 4PO₄³⁻ + 6H₂O
SO₃SO₃ + 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

Oxides of Period 3

8.3Why is Al₂O₃ amphoteric?

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

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

ChlorideOx. no.StructureWith waterpH
NaCl+1giant ionichydration only7
MgCl₂+2giant ionichydration, slight hydrolysis6.5
AlCl₃ (Al₂Cl₆)+3simple molecularhydration, appreciable hydrolysis3
SiCl₄+4simple molecularcomplete hydrolysis2
PCl₅+5simple molecularcomplete hydrolysis2
9

Chlorides of Period 3

9.2Equations, in order of the elements

ChlorideExcess waterLimited water
NaClNaCl(s) → Na⁺(aq) + Cl⁻(aq)–
MgCl₂MgCl₂ + 6H₂O → [Mg(H₂O)₆]²⁺ + 2Cl⁻
[Mg(H₂O)₆]²⁺ ⇌ [Mg(H₂O)₅(OH)]⁺ + H⁺		–
AlCl₃AlCl₃ + 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
SiCl₄SiCl₄ + 2H₂O → SiO₂ + 4HCl–
PCl₅PCl₅ + 4H₂O → H₃PO₄ + 5HClPCl₅ + H₂O → POCl₃ + 2HCl
9

Chlorides of Period 3

9.3Why is AlCl₃(aq) acidic?

  1. Al³⁺ has a high charge density
  2. → strongly polarises the O–H bonds of the water ligands
  3. → 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
10

Group 17: The Halogens

10.1Physical properties

At r.t.p.In hexaneIn waterX–X / kJ mol⁻¹
F₂pale yellow gasvery pale yellow–158
Cl₂greenish-yellow gaspale yellowpale yellow244
Br₂red-brown liquidorange-redorange193
I₂black solidpurplebrown151
NoteF–F is weaker than Cl–Cl: the lone pairs on the two small, close F atoms repel each other.
10

Group 17: The Halogens

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?

  1. All halogens are simple molecular, non-polar X₂ molecules
  2. Down the group, Mr ↑ → no. of electrons ↑ → electron cloud more polarisable → stronger id–id
  3. → 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⁻.
  1. Down the group, the number of principal quantum shells ↑ → atomic radius ↑, and shielding ↑, which largely cancels out the increase in nuclear charge
  2. → the incoming e⁻ is less strongly attracted by the nucleus → X₂ gains electrons less readily
  3. ∴ oxidising power ↓: F₂ > Cl₂ > Br₂ > I₂
10

Group 17: The Halogens

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.

HalogenHalideReactionObservation
F₂Cl⁻, Br⁻, I⁻oxidises all of themF₂ also oxidises water, so it is not used in aqueous solution
Cl₂Br⁻Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂colourless solution turns orange, orange-red in hexane
Cl₂I⁻Cl₂ + 2I⁻ → 2Cl⁻ + I₂colourless solution turns brown, purple in hexane
Br₂Cl⁻no reactionsolution stays orange (from Br₂)
Br₂I⁻Br₂ + 2I⁻ → 2Br⁻ + I₂orange solution turns brown, purple in hexane
I₂Cl⁻, Br⁻no reactionsolution stays brown (from I₂)
10

Group 17: The Halogens

Quick checkChlorine water is added to aqueous KI, then the mixture is shaken with hexane. What colour is the hexane layer?

Acolourless
Borange-red
Cpurple
Dpale green
NoteC: Cl₂ oxidises I⁻ to I₂, which is purple in hexane.
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

HXH–X / kJ mol⁻¹On heating
HF562stable
HCl431stable
HBr366some brown Br₂ fumes
HI299purple I₂ fumes with a hot rod
  1. Valence orbital of X is larger and more diffuse → less effective overlap with H 1s
  2. → 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.

InformationDeduction
solid, very low conductivitysemiconductor
chloride is a liquid non-conductorsimple molecular chloride
white solid + strongly acidic solutionSiCl₄ + 2H₂O → SiO₂ + 4HCl

∴ X is silicon

12

Deducing and Predicting

12.2Predict: astatine, below iodine

PredictionReason
less electronegative than Imore shells, more shielding
dark solid, less volatilemore electrons, stronger id–id
HAt decomposes more easilyH–At bond weaker than H–I
cannot oxidise Cl⁻ to Cl₂weaker oxidising agent than I₂
13

Uses of Group 2 and Group 17 Compounds

  • MgO: refractory lining of furnaces (high m.p., low reactivity)
  • CaCO₃ (limestone): construction materials
  • CaO (quicklime): liming acidic soil, as most crops grow best at pH 6 to 7
  • BaSO₄: the barium meal in X-ray diagnosis
  • F: fluoride in toothpaste, PTFE on non-stick cookware, CFCs (these deplete the ozone layer)
  • Cl: disinfecting water, PVC, ClO⁻ in bleach
  • Br: silver bromide in photographic film
  • I: tincture of iodine, an antiseptic

Exit ticket

1Why is Na⁺ smaller than Na?

Na loses its valence electron, so Na⁺ has one less principal quantum shell.[1]

2Write the equation for SiCl₄ with water. What is the pH?

SiCl₄ + 2H₂O → SiO₂ + 4HCl[1]
pH 2[1]

3Why is BaCO₃ more thermally stable than MgCO₃?

Ba²⁺ is larger, with lower charge density[1]
so it polarises the CO₃²⁻ anion less: C–O bonds are weakened less.[1]

4Will Br₂ displace Cl⁻ from aqueous NaCl?

No: Br₂ is a weaker oxidising agent than Cl₂.[1]