Periodicity — Period 3

1. Trend in Atomic Radius

Trend: atomic radius decreases across the period (Na → Cl). (Board heading says "down the group" — slip; the 3 reasons and graph are for across the period. Down a group radius increases.)

1. Nuclear charge increases (more protons)
2. Shielding remains constant (same number of shells)
3. Valence electrons feel a stronger pull → pulled inwards → radius decreases

Same 3-point reasoning as ionisation energy in Atomic Structure — copy it every time.

2. Trend in Ionic Radius

Ions formed (remember!)

Na	Mg	Al	Si	P	S	Cl
Na⁺	Mg²⁺	Al³⁺	Si⁴⁺	P³⁻	S²⁻	Cl⁻

Metals (and Si) lose electrons → cations. P needs 3 e⁻, S needs 2, Cl needs 1 to complete the octet (like Ar) → anions. Argon is not discussed — inert noble gas, no reactions.

• Na⁺ → Si⁴⁺ : decreases, and P³⁻ → Cl⁻ : decreases — nuclear charge ↑, shielding constant, valence electrons feel stronger pull (same 3 reasons)
• Big increase between Si and P: the cations have 2 shells (lost the outer shell), the anions have 3 shells (gained electrons in the same shell).

3. Trend in First Ionisation Energy

General: increases across the period (nuclear charge ↑, shielding constant, radius ↓ → more energy to remove an e⁻). Down a group it decreases. (Board: "visit chapter Atomic Structure".)

• Mg → Al dips: Mg outer e⁻ is in 3s (lower energy); Al's is in 3p¹ (higher energy, easier to remove).
• P → S dips: P is 3p³ (one e⁻ per orbital); S is 3p⁴ — two e⁻ share one orbital → spin-pair repulsion → easier to remove.

4. Structures & Bonding of the Elements

Element	Bonding	Structure	m.p./b.p.	Conductivity
Na, Mg, Al	metallic	giant metallic lattice	high	good
Si	covalent	giant covalent	high (highest)	semiconductor
P₄, S₈, Cl₂	covalent	simple molecular	low	poor

• Giant structures: metallic, ionic and giant covalent are all giant (e.g. diamond, graphite, graphene, nanotubes, SiO₂). Buckyball (C₆₀) is simple molecular.
• Metals conduct in all states (solid, liquid, even gas) — free-moving electrons.
• P exists as P₄, S as S₈, Cl as Cl₂ — know these!

Si is a semiconductor: conducts a little, depending on conditions — draw it a small bit above zero on the graph.

5. Trend in Melting / Boiling Point

• Na → Al increases: (1) cationic charge increases (+1, +2, +3), (2) number of delocalised electrons per atom increases (1, 2, 3) → stronger metallic bonding.
• Si is highest: giant covalent structure — so many strong covalent bonds to break that it is above the metals.
• Sharp decrease after Si (between Si and P): P₄, S₈, Cl₂ are simple molecular — only temporary dipoles (intermolecular forces) are broken, not covalent bonds. All three are homonuclear → no permanent dipole; no H bonding (H not bonded to F/O/N).
• S₈ > P₄ > Cl₂: temporary-dipole strength depends on the number of electrons. S₈ has the most e⁻ (highest), Cl₂ the fewest (lowest).

6. Trend in Electrical Conductivity

• Na → Al increases — metals; number of delocalised electrons per atom increases (1 → 2 → 3). Al has the greatest conductivity in Period 3.
• Si is a semiconductor (small conductivity).
• P → Cl: zero — simple covalent, no free moving ions or electrons.

7. Reactions of the Elements with Oxygen

Most products can be guessed; learn the starred ones (esp. P₄O₁₀).

Element	Reactivity / observation	Equation
Na	vigorous, yellow flame (like sodium-vapour street lamps)	4Na(s) + O₂(g) → 2Na₂O(s)
Mg	vigorous, white flame	2Mg(s) + O₂(g) → 2MgO(s)
Al	not vigorous normally — protective Al₂O₃ layer. Powdered Al (more surface area) reacts vigorously	4Al(s) + 3O₂(g) → 2Al₂O₃(s)
Si	reacts slowly	Si(s) + O₂(g) → SiO₂(s)
P ★	reacts vigorously; product must be memorised: P₄O₁₀	P₄(s) + 5O₂(g) → P₄O₁₀(s)
S	once ignited, burns gently with a blue flame	S₈(s) + 8O₂(g) → 8SO₂(g)
Cl	does not react with oxygen

• Why does Al resist? Like rust on iron, a thin Al₂O₃ layer forms on the surface and stops O₂ reaching more Al. Powdering increases surface area.
• Na₂O, MgO, Al₂O₃ are solids (ionic, high m.p.); SO₂ is a gas.
• Contact process: SO₂ further oxidised using V₂O₅ catalyst (vanadium(V) oxide, "vanadium pentoxide"): 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) — reversible (see Equilibrium chapter).
• Why P₄O₁₀ and SO₂? "Internal chemistry" — they are the most stable oxides (P and S form others too: P₄O₆, SO₃).

8. Properties of the Oxides

Period 3 element	Na	Mg	Al	Si	P	S
Formula	Na₂O	MgO	Al₂O₃	SiO₂	P₄O₆ P₄O₁₀	SO₂ SO₃
Oxidation no. of element	+1	+2	+3	+4	+3 +5	+4 +6
Solubility in water	soluble	insoluble	insoluble	insoluble	soluble	soluble
Acid / base nature	basic	basic	amphoteric	acidic	acidic	acidic
Relative m.p./b.p.	high	high	high	high	low	low
Electrical conductivity	good (molten or aq)	good (molten or aq)	good (molten or aq)	poor	poor	poor
Chemical bonding	ionic	ionic	ionic	covalent	covalent	covalent
Structure	giant	giant	giant	giant	simple	simple
Used in ceramics?	No	Yes	Yes	Yes	No	No

P₄O₆ is only for completeness — P₄O₁₀ is more than enough to remember.

Reasons behind the table

• Oxidation numbers: +1…+6 = group number (valence electrons given to O).
• Solubility — two ways to dissolve: (1) already ionic → ions separate; (2) reacts with water to give a product that is ionic/soluble (Na₂O → Na⁺ + OH⁻ not Na⁺ + O²⁻; P₄O₁₀ → H₃PO₄ → H⁺ + PO₄³⁻; SO₃ → H₂SO₄). Don't only check "is it ionic?".
• MgO is the least soluble Group 2 oxide (solubility grows down the group). Al₂O₃ insoluble too. SiO₂ = sand → insoluble.
• m.p./b.p.: ionic (giant) high; SiO₂ giant covalent high — easy to forget in a rush!; P₄O₁₀, SO₂ simple molecular → low. Simple/giant: ionic, metallic, giant covalent → giant; only small covalent molecules are simple.
• Conductivity of ionic oxides: NOT in solid state. Conduct when molten or aqueous because ions become free to move.

Why not Na₂O? Na₂O is very reactive/basic — could decompose at much lower temperature (+1 charge → weaker ionic bonds than Mg²⁺/Al³⁺, lower m.p.), react (displace metals), give explosive reactions with water. So MgO, Al₂O₃, SiO₂ are used. P and S oxides are simple molecular (low m.p.).

Trends across the period (oxides)

• Oxidation number always positive: oxygen is more electronegative than any Period 3 element (O is 2nd most EN after F: F-O-N-Cl) → it pulls electrons away.
• Oxidation numbers increase: across the period valence electrons increase, hence more electrons are available for bonding. Ox. no. = group number.
• Bonding changes ionic → covalent: the difference in electronegativity between O and the Period 3 element gets less.

9. Acid / Base Nature of Oxides

Sodium oxide — strongly alkaline (pH 13)

Na₂O + 2HCl → 2NaCl + H₂O also shows its basic nature (neutralisation).

Magnesium oxide — weakly alkaline (pH 9)

Not very soluble; reaction with water is very slow; only slightly raises pH (8 or 9 fine).

Aluminium oxide — AMPHOTERIC ★ (behaves as base AND acid)

As a base (with any acid — deduce the salt; Al³⁺):

As an acid (with alkali) — must memorise:

Board writes it unbalanced (Al₂O₃ + NaOH + 3H₂O → NaAl(OH)₄); balanced version above needs 2NaOH and 2 NaAl(OH)₄ — sir said "add a 3" for the water.

Silicon dioxide — acidic (reacts with hot conc. NaOH) ★

Doesn't react with water (sand). Memorise the product Na₂SiO₃ (balancing: 2NaOH; board shows it unbalanced). Needs hot, concentrated NaOH.

Phosphorus(V) oxide — strong acid (pH 3)

pH 1–3 acceptable; board uses 3. Na₃PO₄ from PO₄³⁻.

Sulfur oxides — strongly acidic (pH 1)

10. Reactions of the Elements with Chlorine

Element	Reactivity	Equation
Na	react vigorously	2Na(s) + Cl₂(g) → 2NaCl(s)
Mg		Mg(s) + Cl₂(g) → MgCl₂(s)
Al		2Al(s) + 3Cl₂(g) → Al₂Cl₆(s) (more stable form; AlCl₃ also accepted)
Si	react slowly	Si(s) + 2Cl₂(g) → SiCl₄(l)
P		P₄(s) + 10Cl₂(g) → 4PCl₅(s) or P₄ + 6Cl₂ → 4PCl₃

Memorise SiCl₄ and PCl₅/PCl₃ (PCl₃ & PCl₅ appear again in Organic chemistry — either is accepted). Sulfur's chloride S₂Cl₂ is not well known — don't focus on it.

11. Properties of the Chlorides

Period 3 element	Na	Mg	Al	Si	P	S
Formula	NaCl	MgCl₂	Al₂Cl₆	SiCl₄	PCl₃ PCl₅	S₂Cl₂
Oxidation no. of element	+1	+2	+3	+4	+3 +5	+1
Chemical bonding	ionic	ionic	covalent & dative	covalent	covalent	covalent
Structure	giant	giant	simple	simple	simple	simple
Observation when added to water	solid dissolves; colourless solution		solid dissolves to give colourless acidic solution (+ steamy fumes of HCl(g) with SiCl₄ & PCl₅)
pH of solution	7	6.5	3	1	1	1

• Oxidation numbers: positive (Cl more electronegative) and increase with group number (valence electrons) — same reasons as oxides.
• Al₂Cl₆: each AlCl₃ has covalent + dative bonds; two AlCl₃ dimerise via a dative bond (Cl lone pair → Al). Not ionic → simple molecular (separate molecules, not a giant lattice). Important example!
• Al (and Si) chlorides are simple covalent — unlike Al₂O₃ which is ionic/giant.

12. Reactions of Chlorides with Water

• NaCl & MgCl₂: do not react — dissolve completely (dissociate into ions): Na⁺ + Cl⁻; Mg²⁺ + 2Cl⁻.

Why is MgCl₂ pH 6.5 and AlCl₃ pH 3? — polarisation of water

H₂O is already polar (O δ⁻, H δ⁺). A small, highly-charged cation attracts the O of water and polarises it → O–H bond weakens → H⁺ released → solution acidic.

• MgCl₂: Mg²⁺ polarises some water molecules and releases a little H⁺ → pH ≈ 6.5 (only slightly acidic). Na⁺ (+1) cannot polarise water enough → NaCl neutral, pH 7.
• Al₂Cl₆: breaks down in water and dissolves; Al³⁺ (charge +3) polarises water much more → more H⁺ → pH ≈ 3 (considerably acidic, not a strong acid).

Hydrolysis ★ (SiCl₄ and PCl₅ — memorise)

Board shows PCl₅ + H₂O → 5HCl + H₃PO₄ with the water coefficient unwritten; balanced needs 4H₂O (H: 8 = 5+3).

• Observations with SiCl₄: solid dissolves; off-white (white / yellowish-white) precipitate of SiO₂ (sand, insoluble); steamy fumes of HCl. Litmus turns red = a test (not observation) for HCl.
• SiCl₄ is not acidic itself — it forms HCl after hydrolysis (strong acid) → pH 1.
• PCl₅: gives two strong acids, H₃PO₄ + HCl → pH 1. (Strongly acidic; steamy HCl fumes.)

13. Past-Paper Questions (from the board)

14. Full Chapter Cheat-Sheet