8.1
Physical Properties of Group 13 Elements
Overview of Group 13
Group 13 (formerly Group IIIA or Group III) consists of: Boron (B), Aluminium (Al), Gallium (Ga), Indium (In), Thallium (Tl) (and Nihonium Nh, synthetic). This group bridges the p-block non-metals and metals — boron is a metalloid (semiconductor); aluminium and the rest are metals. All have the outer electron configuration ns²np¹ (3 valence electrons).
| Element | Z | Config | M.p. (°C) | Density (g/cm³) | Character |
|---|---|---|---|---|---|
| Boron (B) | 5 | [He]2s²2p¹ | 2076 | 2.34 | Metalloid / semi-conductor |
| Aluminium (Al) | 13 | [Ne]3s²3p¹ | 660 | 2.70 | Reactive metal |
| Gallium (Ga) | 31 | [Ar]3d¹⁰4s²4p¹ | 30 | 5.91 | Metal (m.p. just above room temp) |
| Indium (In) | 49 | [Kr]4d¹⁰5s²5p¹ | 157 | 7.31 | Soft metal |
| Thallium (Tl) | 81 | [Xe]4f¹⁴5d¹⁰6s²6p¹ | 304 | 11.85 | Soft heavy metal; toxic |
Trends Down Group 13
- Atomic radius: increases (new shells added)
- Ionic radius (M³⁺): increases
- First ionisation energy: decreases (outer electron further, more shielded)
- Electronegativity: decreases
- Metallic character: increases (B = metalloid; Tl = heavy metal)
- Melting point: No simple trend — Ga has anomalously low m.p. (30°C, melts in the palm of your hand)
- Oxidation state: +3 dominates for lighter members; +1 (inert pair effect) becomes more common for Tl
Inert Pair EffectGoing down Group 13 (and Group 14, 15), the ns² pair of electrons becomes increasingly reluctant to participate in bonding. This is because the ns² electrons penetrate the inner electron cores more and experience greater effective nuclear charge, making them harder to ionise. Result: heavier elements (Tl, Pb, Bi) prefer lower oxidation states (+1, +2, +3 rather than +3, +4, +5). For Group 13: Tl⁺ is more stable than Tl³⁺.
8.2
Reaction of Aluminium
Aluminium — Seemingly Unreactive but Actually Reactive
Pure aluminium is actually highly reactive (standard electrode potential E° = −1.66 V). However, in air it rapidly forms a thin, adherent, impervious layer of aluminium oxide (Al₂O₃) that protects the underlying metal from further reaction. This is called passivation.
This is why aluminium appears unreactive in everyday life despite being a very reactive metal. Anodising thickens this protective oxide layer deliberately.
Reactions with Non-metals
With oxygen (burns in air when oxide layer is removed):
4Al + 3O₂ → 2Al₂O₃ (exothermic; aluminium powder burns vigorously)
With halogens:
2Al + 3Cl₂ → 2AlCl₃ (aluminium chloride)
2Al + 3Br₂ → 2AlBr₃
With nitrogen (at high temperature):
2Al + N₂ → 2AlN (aluminium nitride; reacts with water to give NH₃)
With sulfur:
2Al + 3S → Al₂S₃
Reaction with Dilute Acids
2Al + 6HCl → 2AlCl₃ + 3H₂ (dissolves in HCl after oxide layer removed)
2Al + 3H₂SO → Al₂(SO)₃ + 3H₂
Note: Al is PASSIVATED by concentrated HNO₃ (conc. H₂SO also passivates)
— oxide layer formed and protects the metal (similar to Fe with conc. HNO₃)
Reaction with Alkalis (NaOH)
This is the key distinguishing reaction of aluminium. Unlike most metals, aluminium dissolves in hot concentrated NaOH solution (amphoteric behaviour):
2Al + 2NaOH + 2H₂O → 2NaAlO₂ + 3H₂ (cold/dilute NaOH)
Or in excess NaOH:
2Al + 2NaOH + 6H₂O → 2Na[Al(OH)₄] + 3H₂
(sodium tetrahydroxoaluminate — complex ion [Al(OH)₄]⁻)
Overall ionic equation:
2Al + 2OH⁻ + 6H₂O → 2[Al(OH)₄]⁻ + 3H₂
This happens because Al₂O₃ is amphoteric: it reacts with NaOH to give aluminate ion, which allows the metal underneath to continue reacting.
Thermite Reaction
2Al + Fe₂O₃ → Al₂O₃ + 2Fe (ΔH very negative → temp ~3000°C)
(thermite reaction — welding railway tracks; incendiary)
Aluminium's high affinity for oxygen (very negative ΔHf for Al₂O₃) allows it to reduce many metal oxides. Used industrially to produce Cr, Mn from their oxides (aluminothermic reduction).
8.3
Oxides and Hydroxides of Group 13 Elements
Aluminium Oxide (Al₂O₃)
Al₂O₃ (alumina) is the main oxide of aluminium. It is amphoteric — reacts with both acids and bases:
With acid (acts as base):
Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
Al₂O₃ + 3H₂SO → Al₂(SO)₃ + 3H₂O
With alkali (acts as acid):
Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O (simple form)
Al₂O₃ + 2NaOH + 3H₂O → 2Na[Al(OH)₄] (excess water)
Aluminium Hydroxide (Al(OH)₃)
Al(OH)₃ is also amphoteric:
With acid:
Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O
With excess alkali:
Al(OH)₃ + NaOH → Na[Al(OH)₄] (sodium tetrahydroxoaluminate)
Al(OH)₃ is a white gelatinous precipitate formed when NaOH is added to an Al³⁺ solution. It dissolves in excess NaOH (distinguishing Al³⁺ from other metal ions like Mg²⁺ whose hydroxide is insoluble in excess NaOH).
Comparison of Group 13 Oxide Properties
| Oxide | Character | Notes |
|---|---|---|
| B₂O₃ | Acidic | Dissolves in water to form boric acid H₃BO₃; reacts with bases |
| Al₂O₃ | Amphoteric | Reacts with both acids and NaOH; high m.p. (2072°C) |
| Ga₂O₃ | Amphoteric | Similar to Al₂O₃ |
| In₂O₃ | Weakly basic | More metallic than Al; oxide is more basic |
| Tl₂O₃ / Tl₂O | Basic | Tl⁺ is stable (inert pair); Tl₂O is the dominant oxide for Tl |
Trend in oxide characterAcross Group 13, oxide character changes from acidic (B₂O₃) → amphoteric (Al₂O₃, Ga₂O₃) → basic (In₂O₃, Tl₂O). This reflects increasing metallic character down the group. The same trend is observed across a period: from acidic oxides (right) to basic oxides (left), with amphoteric oxides in the middle.
8.4
Anomalous Properties of Boron
Why Boron is Anomalous
Boron (Period 2, Group 13) is a metalloid/semiconductor — not a metal. It has a very high melting point (2076°C), forms covalent compounds almost exclusively, and its chemistry differs greatly from aluminium. Boron's anomalies arise from its very small atomic radius and high charge density (like all Period 2 first members). Boron resembles silicon (Si) diagonally.
| Property | Typical Group 13 (Al) | Boron (anomalous) | Resembles Si (diagonal) |
|---|---|---|---|
| Physical state | Soft, metallic solid | Hard, black solid; semiconductor | Si is a semiconductor |
| Bonding in compounds | Ionic (Al³⁺ ion) | Predominantly covalent (B has high charge density) | Si forms covalent compounds |
| Oxide character | Al₂O₃ amphoteric | B₂O₃ is acidic (forms boric acid with H₂O) | SiO₂ is acidic |
| Chloride | AlCl₃ (covalent, Lewis acid) | BCl₃ (covalent, Lewis acid) | SiCl (covalent) |
| Hydride | AlH₃ (polymeric) | B₂H⁶ (diborane — bridged H, unique structure) | Silanes (SiₙH₂ₙ⁺₂) |
| Forms complex anions | [Al(OH)₄]⁻ | [BF₄]⁻, [B(OH)₄]⁻ | [SiF⁶]²⁻ |
| Reaction with NaOH | Al reacts: 2Al + 2NaOH + 6H₂O → 2Na[Al(OH)₄] + 3H₂ | B₂O₃ + 2NaOH → 2NaBO₂ + H₂O | SiO₂ + 2NaOH → Na₂SiO₃ + H₂O |
Electron Deficiency of Boron
Boron has only 3 valence electrons. In BCl₃ and BF₃, boron forms 3 bonds → only 6 electrons around B (not 8 — electron deficient). This makes boron compounds strong Lewis acids (electron pair acceptors):
BF₃ + F⁻ → [BF₄]⁻ (fluoride donates lone pair to empty orbital on B)
BF₃ + NH₃ → F₃B←NH₃ (dative bond: N donates lone pair to B)
Diborane (B₂H⁶) has a unique 3-centre 2-electron (3c-2e) bond involving bridging H atoms — different from any other element's hydride.
Boron in glass and ceramicsBoron is used in borosilicate glass (Pyrex): B₂O₃ + SiO₂ network → very low coefficient of thermal expansion → resists thermal shock (laboratory glassware, cooking dishes). Also used in borax (Na₂B₄O⁷·10H₂O) for washing powder and glass.
8.5
Identification of Al³⁺ Ion in Solution
Test for Al³⁺
The identification of Al³⁺ relies on the amphoteric nature of Al(OH)₃:
- Add NaOH(aq) dropwise to the test solution.
- A white gelatinous precipitate of Al(OH)₃ forms:
Al³⁺(aq) + 3OH⁻(aq) → Al(OH)₃(s) (white gelatinous ppt) - Continue adding NaOH — the precipitate dissolves in excess NaOH:
Al(OH)₃(s) + OH⁻(aq) → [Al(OH)₄]⁻(aq) (sodium tetrahydroxoaluminate, colourless)
Key distinction: The dissolution in excess NaOH distinguishes Al³⁺ from Mg²⁺ (Mg(OH)₂ does NOT dissolve in excess NaOH — not amphoteric).
| Ion | With NaOH (drop by drop) | Excess NaOH |
|---|---|---|
| Al³⁺ | White gelatinous precipitate Al(OH)₃ | Dissolves → [Al(OH)₄]⁻ (colourless) |
| Mg²⁺ | White gelatinous precipitate Mg(OH)₂ | Does NOT dissolve (insoluble in excess NaOH) |
| Fe²⁺ | Green precipitate Fe(OH)₂ | Does NOT dissolve (turns brown in air → Fe(OH)) |
| Fe³⁺ | Red-brown precipitate Fe(OH) | Does NOT dissolve |
| Cu²⁺ | Blue precipitate Cu(OH)₂ | Does NOT dissolve |
| Zn²⁺ | White precipitate Zn(OH)₂ | Dissolves → [Zn(OH)₄]²⁻ (Zn also amphoteric) |
Zn²⁺ also dissolves in excess NaOHBoth Al³⁺ and Zn²⁺ give white precipitates that dissolve in excess NaOH. To distinguish them: use a flame test (Al gives no characteristic colour; Zn also none), or add NH₃(aq) — Al(OH) does NOT dissolve in NH(aq), but Zn(OH)₂ does.
Aluminon Test
A more specific test for Al³⁺ in qualitative analysis uses aluminon (aurintricarboxylic acid) dye. Adding aluminon to a neutral or slightly acidic solution containing Al³⁺ produces a characteristic red lake (precipitate). This is more specific than the NaOH test for distinguishing Al³⁺ from Zn²⁺.
8.6
Uses of Group 13 Elements
| Substance | Uses |
|---|---|
| Aluminium (Al) | Lightweight alloys (aircraft, cars, bicycles — alloys with Cu, Mg, Mn, Si); electrical cables (low density, good conductor); food packaging (foil); building materials; drinks cans; heat sinks |
| Al₂O₃ (alumina) | Refractory linings (very high m.p. ~2050°C); abrasive (emery, corundum); manufacture of Al metal (Hall-Héroult electrolysis); Al₂O₃ in rubies (Cr-doped) and sapphires (Ti-doped) |
| Al(OH)₃ | Antacid (neutralises stomach acid); water purification (flocculation — gelatinous Al(OH) traps impurities); mordant in dyeing textiles |
| AlCl₃ | Friedel-Crafts catalyst in organic synthesis (Lewis acid catalyst); anhydrous AlCl in acylation and alkylation reactions |
| Al₂(SO)₃ | Water treatment (added to water, hydrolyses to give Al(OH) which flocculates suspended particles); paper sizing; fire retardant |
| Aluminothermic reduction | Thermite reaction: Al reduces metal oxides at very high temperature; used to weld railway tracks (produces molten Fe) and to produce Cr, Mn metals |
| Boron (B) | Boron steel (neutron absorber in nuclear reactors); semiconductors; boron carbide (B₄C) armour; boron nitride (BN) lubricant; borosilicate glass (Pyrex) |
| Borax (Na₂B₄O⁷·10H₂O) | Washing powder; glass and ceramic manufacture; flux in soldering; fire retardant; antiseptic |
| Boric acid (H₃BO₃) | Antiseptic and eye wash; wood preservative (insecticide); mild buffer in laboratory |
| Gallium (Ga) | GaAs semiconductors (LEDs, solar cells, laser diodes); GaN (bright blue/white LEDs); gallium alloys (low m.p. → thermometers) |
| Indium (In) | ITO (indium tin oxide) for touchscreens and flat-panel displays; soldering alloys; thin-film solar cells |
| Thallium (Tl) | Thallium-doped NaI crystal detectors (gamma ray detection); formerly in rat poison (now banned — very toxic); optical lenses (infra-red) |
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Section 8.1 — Group 13 Properties & Aluminium
10 QuestionsQ1 of 10
Aluminium appears unreactive in air due to:
Passivation: Al reacts instantly with O₂ → thin Al₂O₃ layer → prevents further oxidation. This protective layer makes Al useful outdoors despite being reactive.
Q2 of 10
Al₂O₃ is described as amphoteric because:
Amphoteric: Al₂O₃ + 3H₂SO₄ → Al₂(SO₄)₃ + 3H₂O (acts as base). Al₂O₃ + 2NaOH + 3H₂O → 2Na[Al(OH)₄] (acts as acid). Reacts with both.
Q3 of 10
Aluminium is extracted by:
Al₂O₃ melts at 2072°C — too high. Dissolved in molten cryolite (Na₃AlF₆, lower m.p.) → electrolysis. Carbon anodes, steel cathode.
Q4 of 10
Al(OH)₃ reacts with NaOH to give:
Al(OH)₃ + OH⁻ → [Al(OH)₄]⁻ (aluminate). Amphoteric behaviour. Also: Al(OH)₃ + 3H⁺ → Al³⁺ + 3H₂O.
Q5 of 10
The electron deficiency of BF₃ (Group 13) means it acts as:
BF₃: B has only 6 electrons (no lone pair). Empty p orbital → accepts lone pairs → Lewis acid. Forms adducts with Lewis bases: BF₃ + NH₃ → F₃B←NH₃.
Q6 of 10
Which property distinguishes Al from transition metals?
Al always +3 (lost all 3 outer electrons). Transition metals: variable oxidation states (d electrons involved). Al is in Group 13 — not a transition metal.
Q7 of 10
The thermite reaction (Al + Fe₂O₃) is used for:
2Al + Fe₂O₃ → Al₂O₃ + 2Fe. Very exothermic (ΔH ≈ −850 kJ/mol). Al is a better reducing agent than Fe. Used to weld railroad tracks in situ.
Q8 of 10
Anodising aluminium involves:
Anodising: Al as anode in H₂SO₄ → O₂ produced at anode oxidises Al → thicker Al₂O₃ layer. Can be dyed. More durable protection.
Q9 of 10
Al reacts with dilute HCl to give:
2Al + 6HCl → 2AlCl₃ + 3H₂. Al is oxidised (+3), H is reduced (0). First, oxide layer must be dissolved/abraded.
Q10 of 10
Why is pure aluminium soft but alloyed aluminium (duralumin) hard?
Alloys: foreign atoms disrupt regular lattice → harder for planes to slip → harder, stronger material. Duralumin (Al + Cu + Mg): used in aircraft (strong, lightweight).
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✏️
Exercises
- Write balanced equations to show that Al₂O₃ is amphoteric, reacting with both dilute H₂SO and NaOH solution. Explain what "amphoteric" means.
Amphoteric = a substance that can react with both acids AND bases.
With acid (acts as base): Al₂O₃ + 3H₂SO → Al₂(SO)₃ + 3H₂O
With alkali (acts as acid): Al₂O₃ + 2NaOH + 3H₂O → 2Na[Al(OH)₄]
In the first reaction, Al₂O₃ accepts H⁺ from the acid (base behaviour). In the second, it donates a proton (indirectly) to OH⁻ (acid behaviour). - Describe how you would distinguish between solutions containing Al³⁺ and Mg²⁺ using NaOH solution. Write ionic equations for both precipitate formation and the dissolution of Al(OH)₃.
Add NaOH(aq) dropwise to each solution. Both give a white gelatinous precipitate. Continue adding excess NaOH:
• Al³⁺ solution: precipitate dissolves (Al(OH) is amphoteric). Mg²⁺ solution: precipitate remains (Mg(OH)₂ not amphoteric — does not dissolve in excess NaOH).
Formation: Al³⁺(aq) + 3OH⁻(aq) → Al(OH)₃(s)
Dissolution: Al(OH)₃(s) + OH⁻(aq) → [Al(OH)₄]⁻(aq) - Explain why aluminium appears unreactive in everyday use despite being a very reactive metal. Why does it dissolve in NaOH solution?
Aluminium rapidly forms a thin, tough, impervious layer of Al₂O₃ on its surface when exposed to air. This oxide layer adheres strongly and acts as a physical barrier, preventing O₂, H₂O, and dilute acids from reaching the underlying metal (passivation). The standard electrode potential of Al is −1.66 V, confirming it is very reactive.
In NaOH solution: Al₂O₃ reacts with NaOH (amphoteric oxide): Al₂O₃ + 2NaOH + 3H₂O → 2Na[Al(OH)₄]. Once the protective oxide layer is dissolved, the bare Al metal reacts with water and NaOH: 2Al + 2NaOH + 6H₂O → 2Na[Al(OH)₄] + 3H₂. The continuous dissolution of the oxide layer allows the reaction to proceed. - State THREE ways in which boron behaves differently from aluminium. What element does boron resemble via the diagonal relationship, and why?
(1) B is a metalloid/semiconductor; Al is a reactive metal. (2) B forms covalent compounds (BCl, BF); Al forms predominantly ionic compounds (AlCl is covalent but Al³⁺ exists in solution). (3) B₂O is acidic (forms boric acid with water); Al₂O is amphoteric. (Also acceptable: B forms unique structures like diborane B₂H⁶ with 3-centre 2-electron bonds; B has electron-deficient compounds that are strong Lewis acids.)
B resembles silicon (Si) (diagonal relationship: Period 2, Group 13 vs Period 3, Group 14). Both B and Si: have covalent bonding, semiconducting properties, acidic oxides (B₂O and SiO₂ both react with alkali), form complex anions ([BO₃]⁻, [SiO₄]⁻), and form covalent chlorides (BCl, SiCl₄). The diagonal relationship exists because B and Si have similar charge densities despite being in different groups. - Write a balanced equation for the thermite reaction and explain why it is used to weld railway tracks.
2Al + Fe₂O₃ → Al₂O₃ + 2Fe
The reaction is highly exothermic (ΔH very negative; temperature reaches ~3000°C). This is hot enough to produce molten iron, which flows into a mould placed around the gap between two rail sections. On cooling, the iron solidifies, joining the rails. Advantages: portable, no external power needed, produces strong welds at the site. The driving force is aluminium's greater affinity for oxygen (more negative ΔHf for Al₂O compared to Fe₂O) → Al reduces Fe₂O. - Explain the inert pair effect and give ONE example of how it affects Group 13 chemistry (especially thallium).
The inert pair effect: going down the group, the ns² pair of valence electrons becomes increasingly stable and reluctant to participate in bonding. This is because the ns electrons penetrate the (n−1)d core and are more tightly held due to poor shielding by d electrons → higher effective nuclear charge on ns electrons → more energy required to ionise them. Result: heavier elements prefer lower oxidation states. Example: In Group 13, Al is typically +3; but thallium (Tl) forms stable +1 compounds (Tl⁺ salts) in addition to +3. Tl⁺ is more stable than Tl³⁺ because the 6s² pair in Tl is inert — not used in bonding. Similarly, Pb²⁺ is more stable than Pb⁺ in Group 14.
🧪
Multiple Choice Quiz — 25 Questions
Unit 8 Quiz
Select one answer per questionQ1
All Group 13 elements have the outer electron configuration:
Group 13: ns²np¹ (3 valence electrons). e.g. Al = [Ne]3s²3p¹; B = [He]2s²2p¹.
Q2
Al₂O₃ is described as amphoteric because it:
Amphoteric: reacts with both acids (Al₂O + H₂SO → Al₂(SO) + H₂O) and bases (Al₂O + NaOH → NaAlO₂ + H₂O).
Q3
When excess NaOH is added to a solution containing Al³⁺, the white precipitate:
Al(OH) + OH⁻ → [Al(OH)₄]⁻ (colourless, soluble). This distinguishes Al³⁺ from Mg²⁺ (Mg(OH)₂ insoluble in excess NaOH).
Q4
Why does aluminium appear unreactive in everyday life?
Al quickly forms a thin, adherent, impervious Al₂O layer (passivation) that protects the underlying metal from attack by air, water, and dilute acids.
Q5
The thermite reaction (2Al + Fe₂O₃ → Al₂O₃ + 2Fe) is used to:
The thermite reaction is extremely exothermic (~3000°C) producing molten Fe which solidifies into the gap between rail sections → permanent weld. No external power needed.
Q6
Boron forms B₂O₃ which is:
B₂O is acidic: dissolves in water to give boric acid HBO; reacts with alkalis. This reflects boron's non-metallic/metalloid character, unlike Al₂O (amphoteric).
Q7
The inert pair effect in Group 13 explains why:
Inert pair effect: going down Group 13, the ns² electrons become harder to ionise. For Tl (6s²6p¹), the 6s² pair is inert → Tl prefers +1 (loses only 6p¹) rather than +3 (loses 6s²6p¹).
Q8
AlCl₃ acts as a Lewis acid because:
Al in AlCl has only 6 electrons (3 bonds → 6e), leaving an empty orbital. It can accept a lone pair from a Lewis base (e.g. Cl⁻, NH) → Lewis acid. Used as catalyst in Friedel-Crafts reactions.
Q9
Boron resembles silicon via the diagonal relationship. Which property do they share?
B and Si: both semiconductors; both have acidic oxides (B₂O and SiO₂); both form covalent halides; both form network covalent structures; both exist as metalloids.
Q10
Al(OH)₃ is used in water treatment because:
Al₂(SO) is added to water → hydrolyses to form gelatinous Al(OH). This sticky gel flocculates (traps) suspended particles, bacteria, and colloids as it settles → water clarified. The floc is then filtered out.
Q11
Which equation correctly shows aluminium dissolving in NaOH(aq)?
2Al + 2NaOH + 6H₂O → 2Na[Al(OH)₄] + 3H₂. Al reacts with NaOH and water to form sodium tetrahydroxoaluminate and hydrogen gas.
Q12
Al₂O₃ is used as a refractory material because:
Refractories must withstand very high temperatures without melting or reacting. Al₂O m.p. ~2050°C; it is chemically inert and structurally stable → ideal for lining furnaces and kilns.
Q13
Which characteristic of boron compounds makes BF₃ a Lewis acid?
B forms 3 bonds in BF → 6 electrons on B (incomplete octet). Empty p orbital on B → can accept a lone pair (Lewis acid). e.g. BF + F⁻ → [BF₄]⁻.
Q14
GaAs is widely used in:
GaAs (gallium arsenide) is a III-V semiconductor with a direct band gap → highly efficient LEDs, laser diodes, and solar cells. GaN is used in bright blue/white LEDs.
Q15
The ionic equation for the reaction of Al(OH)₃ with excess NaOH is:
Al(OH) + OH⁻ → [Al(OH)₄]⁻. One additional OH⁻ is accepted by the amphoteric hydroxide, forming the tetrahydroxoaluminate complex ion.
Q16
Indium tin oxide (ITO) is used in:
ITO (indium tin oxide) is transparent to visible light and electrically conductive → ideal for touchscreens, LCD and OLED displays, and photovoltaic cells as a transparent electrode.
Q17
Aluminium reacts with dilute H₂SO to give:
2Al + 3H₂SO → Al₂(SO) + 3H₂. Al dissolves in dilute H₂SO (after oxide layer removed) to give aluminium sulfate and hydrogen gas.
Q18
Borosilicate glass (Pyrex) is valued because:
Borosilicate glass has <5% thermal expansion vs ordinary glass. It can be rapidly heated or cooled without cracking (thermal shock resistance) → ideal for lab glassware, oven dishes, and telescope mirrors.
Q19
Which product forms when aluminium is heated with nitrogen?
2Al + N₂ → 2AlN (at high temperature). AlN hydrolyses with water to give Al(OH) and NH: AlN + 3H₂O → Al(OH) + NH.
Q20
What distinguishes the white precipitate formed with Al³⁺ from that formed with Mg²⁺ when NaOH is added?
Both give white gelatinous precipitates initially. Key distinction: Al(OH) dissolves in excess NaOH (amphoteric) to give [Al(OH)₄]⁻. Mg(OH)₂ is not amphoteric → remains as precipitate in excess NaOH.
Q21
Aluminium is used for overhead electrical cables instead of copper because:
Al has about 60% the conductivity of Cu but is 3x less dense. For the same electrical resistance per unit length, Al cable weighs ~half of Cu cable → less strain on pylons → Al preferred for long-distance overhead transmission lines.
Q22
Borax (Na₂B₄O⁷·10H₂O) is used in washing powder because:
Borax in washing products: acts as a mild alkali (raises pH for better detergency); water softener (sequestrates Ca²⁺/Mg²⁺); perborate version releases H₂O₂ for bleaching at lower temperatures than chlorine bleach.
Q23
Which ion produces a red-brown precipitate with NaOH that is insoluble in excess NaOH?
Fe³⁺ + 3OH⁻ → Fe(OH)(s) — red-brown precipitate. Insoluble in excess NaOH (not amphoteric). Al³⁺ gives white ppt that dissolves. Mg²⁺ gives white ppt insoluble in excess. Zn²⁺ gives white ppt soluble in excess.
Q24
Al₂(SO)₃ is added to drinking water during treatment because:
Al₂(SO) hydrolyses at neutral pH: Al³⁺ + 3H₂O → Al(OH) + 3H⁺. Gelatinous Al(OH) flocs attract and bind suspended particles, clay, bacteria → settle out or are filtered.
Q25
The outer electron configuration ns²np¹ of Group 13 elements means they typically form:
Group 13 has 3 valence electrons (ns²np¹) → typically forms +3 ions (ionic) or 3 covalent bonds. Exception: inert pair effect in Tl makes +1 stable for the heaviest member.
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Unit Test — 50 Marks
Section A — Short Answer
20 marksQ1 [4 marks]
Describe how you would test for the presence of Al³⁺ ions in solution, clearly distinguishing it from Mg²⁺. Write relevant ionic equations. [4]
Add NaOH(aq) dropwise to the test solution. Both Al³⁺ and Mg²⁺ give a white gelatinous precipitate. [1]
Continue adding excess NaOH: Al(OH) dissolves (solution clears); Mg(OH)₂ remains as precipitate. [1]
Ionic equations:
Precipitation: Al³⁺(aq) + 3OH⁻(aq) → Al(OH)₃(s) [1]
Dissolution in excess: Al(OH)₃(s) + OH⁻(aq) → [Al(OH)₄]⁻(aq) [1]
(Mg(OH)₂ does not dissolve because it is not amphoteric — its Mg–O bonds are not weakened sufficiently by OH⁻.)
Continue adding excess NaOH: Al(OH) dissolves (solution clears); Mg(OH)₂ remains as precipitate. [1]
Ionic equations:
Precipitation: Al³⁺(aq) + 3OH⁻(aq) → Al(OH)₃(s) [1]
Dissolution in excess: Al(OH)₃(s) + OH⁻(aq) → [Al(OH)₄]⁻(aq) [1]
(Mg(OH)₂ does not dissolve because it is not amphoteric — its Mg–O bonds are not weakened sufficiently by OH⁻.)
Q2 [4 marks]
Explain why aluminium is described as "apparently unreactive" yet reacts vigorously with NaOH solution. Write balanced equations for both reactions (with O₂ and with NaOH + H₂O). [4]
"Apparently unreactive": Al rapidly forms a thin, adherent, impervious Al₂O layer in air (passivation). This oxide layer protects the underlying metal from further oxidation, water, and dilute acids. So although Al has E°= −1.66 V (very reactive), it appears inert in everyday conditions. [1]
With O₂: 4Al + 3O₂ → 2Al₂O₃ [1]
In NaOH(aq): the NaOH dissolves the protective Al₂O layer (amphoteric oxide reacts with NaOH), exposing fresh Al which immediately reacts with water/NaOH:
2Al + 2NaOH + 6H₂O → 2Na[Al(OH)₄] + 3H₂ [1]
So in NaOH, both the oxide layer and the metal itself are dissolved → vigorous H₂ evolution. [1]
With O₂: 4Al + 3O₂ → 2Al₂O₃ [1]
In NaOH(aq): the NaOH dissolves the protective Al₂O layer (amphoteric oxide reacts with NaOH), exposing fresh Al which immediately reacts with water/NaOH:
2Al + 2NaOH + 6H₂O → 2Na[Al(OH)₄] + 3H₂ [1]
So in NaOH, both the oxide layer and the metal itself are dissolved → vigorous H₂ evolution. [1]
Q3 [4 marks]
State and explain the inert pair effect, and describe how it affects the stable oxidation state of thallium compared to aluminium. [4]
Inert pair effect: going down Group 13 (and Groups 14, 15), the ns² valence electron pair becomes increasingly stable (reluctant to participate in bonding/ionisation). [1]
Reason: The ns electrons in heavy elements penetrate the filled (n−1)d core. d electrons shield the ns electrons poorly → ns electrons experience higher effective nuclear charge → require more energy to ionise → tend to remain as a "lone pair". [1]
Al (Period 3): 3s²3p¹. Both 3s² and 3p¹ are readily ionised → Al is predominantly +3. [1]
Tl (Period 6): 6s²6p¹. The 6s² pair is inert (poorly shielded by the 4f and 5d core electrons) → Tl prefers to lose only the 6p¹ electron → stable as Tl⁺ (+1). Tl³⁺ (+3) exists but is a strong oxidising agent (unstable, tends to revert to Tl⁺). [1]
Reason: The ns electrons in heavy elements penetrate the filled (n−1)d core. d electrons shield the ns electrons poorly → ns electrons experience higher effective nuclear charge → require more energy to ionise → tend to remain as a "lone pair". [1]
Al (Period 3): 3s²3p¹. Both 3s² and 3p¹ are readily ionised → Al is predominantly +3. [1]
Tl (Period 6): 6s²6p¹. The 6s² pair is inert (poorly shielded by the 4f and 5d core electrons) → Tl prefers to lose only the 6p¹ electron → stable as Tl⁺ (+1). Tl³⁺ (+3) exists but is a strong oxidising agent (unstable, tends to revert to Tl⁺). [1]
Q4 [4 marks]
Write balanced equations for: (a) Al reacting with Cl₂ [1]; (b) Al₂O₃ reacting with NaOH [1]; (c) Al(OH)₃ reacting with HCl [1]; (d) thermite reaction [1]. [4]
(a) 2Al + 3Cl₂ → 2AlCl₃ [1]
(b) Al₂O₃ + 2NaOH + 3H₂O → 2Na[Al(OH)₄] (or Al₂O + 2NaOH → 2NaAlO₂ + H₂O in simplified form) [1]
(c) Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O [1]
(d) 2Al + Fe₂O₃ → Al₂O₃ + 2Fe [1]
(b) Al₂O₃ + 2NaOH + 3H₂O → 2Na[Al(OH)₄] (or Al₂O + 2NaOH → 2NaAlO₂ + H₂O in simplified form) [1]
(c) Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O [1]
(d) 2Al + Fe₂O₃ → Al₂O₃ + 2Fe [1]
Q5 [4 marks]
Describe THREE anomalous properties of boron compared to the rest of Group 13. Name the element boron resembles via the diagonal relationship and give one chemical similarity. [4]
Three anomalous properties of boron (any 3 of): [3]
(1) B is a metalloid/semiconductor; other Group 13 members are metals. Very high m.p. (2076°C) vs Al (660°C).
(2) B forms covalent compounds almost exclusively (B–Cl bond covalent in BCl; B₂H⁶ has unique bridging bonds). Other Group 13 elements form ionic compounds (Al³⁺).
(3) B₂O is acidic (dissolves in water to form HBO; reacts with NaOH). Al₂O is amphoteric; heavier Group 13 oxides are basic.
(4) B forms electron-deficient compounds (BCl, BF) that are strong Lewis acids; Al³⁺ is a Lewis acid in solution but weaker.
(5) Boron nitride (BN) has a graphite-like layered structure and a diamond-like structure (cubic BN).
Diagonal relationship with silicon (Si). One similarity: both B₂O and SiO₂ are acidic oxides that react with NaOH (B₂O + 2NaOH → 2NaBO₂ + H₂O; SiO₂ + 2NaOH → Na₂SiO₃ + H₂O). [1]
(1) B is a metalloid/semiconductor; other Group 13 members are metals. Very high m.p. (2076°C) vs Al (660°C).
(2) B forms covalent compounds almost exclusively (B–Cl bond covalent in BCl; B₂H⁶ has unique bridging bonds). Other Group 13 elements form ionic compounds (Al³⁺).
(3) B₂O is acidic (dissolves in water to form HBO; reacts with NaOH). Al₂O is amphoteric; heavier Group 13 oxides are basic.
(4) B forms electron-deficient compounds (BCl, BF) that are strong Lewis acids; Al³⁺ is a Lewis acid in solution but weaker.
(5) Boron nitride (BN) has a graphite-like layered structure and a diamond-like structure (cubic BN).
Diagonal relationship with silicon (Si). One similarity: both B₂O and SiO₂ are acidic oxides that react with NaOH (B₂O + 2NaOH → 2NaBO₂ + H₂O; SiO₂ + 2NaOH → Na₂SiO₃ + H₂O). [1]
Section B — Extended Answer
30 marksQ6 [8 marks]
Describe fully the chemistry of aluminium with: (a) oxygen [2]; (b) dilute hydrochloric acid [2]; (c) dilute sodium hydroxide solution [2]; (d) iron(III) oxide [2]. Include balanced equations and explain any trends or features. [8]
(a) With oxygen [2]: Al burns in O₂ or reacts slowly in air: 4Al + 3O₂ → 2Al₂O. Very exothermic (ΔHf Al₂O = −1676 kJ/mol). In air, a thin, impervious Al₂O layer forms immediately (passivation) → protects metal. Al powder burns vigorously in pure O₂ with a bright white flame. [2]
(b) With dilute HCl [2]: 2Al + 6HCl → 2AlCl + 3H₂. Al dissolves in dilute HCl once the oxide layer is dissolved. Al⁺ ions form in solution. Vigorous effervescence of H₂ gas. Note: Al is passivated by concentrated HNO (conc. H₂SO also) — oxide layer reformed rapidly → no further reaction with these concentrated acids. [2]
(c) With dilute NaOH(aq) [2]: Amphoteric behaviour. The NaOH dissolves the protective Al₂O layer: Al₂O + 2NaOH + 3H₂O → 2Na[Al(OH)₄]. Exposed metal then reacts: 2Al + 2NaOH + 6H₂O → 2Na[Al(OH)₄] + 3H₂. H₂ gas evolves vigorously; solution becomes colourless (tetrahydroxoaluminate). Most metals do NOT react with NaOH — this reaction is characteristic of amphoteric metals (Al, Zn, Pb, Be). [2]
(d) With Fe₂O (thermite) [2]: 2Al + Fe₂O → Al₂O + 2Fe. Highly exothermic — temperature exceeds 3000°C → produces molten iron. Driving force: Al has much more negative ΔHf for its oxide than Fe → Al displaces Fe from its oxide. Used to weld railway tracks (portable; no external power; produces strong welds) and to prepare reactive metals (Cr, Mn) from their oxides (aluminothermic reduction). [2]
(b) With dilute HCl [2]: 2Al + 6HCl → 2AlCl + 3H₂. Al dissolves in dilute HCl once the oxide layer is dissolved. Al⁺ ions form in solution. Vigorous effervescence of H₂ gas. Note: Al is passivated by concentrated HNO (conc. H₂SO also) — oxide layer reformed rapidly → no further reaction with these concentrated acids. [2]
(c) With dilute NaOH(aq) [2]: Amphoteric behaviour. The NaOH dissolves the protective Al₂O layer: Al₂O + 2NaOH + 3H₂O → 2Na[Al(OH)₄]. Exposed metal then reacts: 2Al + 2NaOH + 6H₂O → 2Na[Al(OH)₄] + 3H₂. H₂ gas evolves vigorously; solution becomes colourless (tetrahydroxoaluminate). Most metals do NOT react with NaOH — this reaction is characteristic of amphoteric metals (Al, Zn, Pb, Be). [2]
(d) With Fe₂O (thermite) [2]: 2Al + Fe₂O → Al₂O + 2Fe. Highly exothermic — temperature exceeds 3000°C → produces molten iron. Driving force: Al has much more negative ΔHf for its oxide than Fe → Al displaces Fe from its oxide. Used to weld railway tracks (portable; no external power; produces strong welds) and to prepare reactive metals (Cr, Mn) from their oxides (aluminothermic reduction). [2]
Q7 [8 marks]
Describe the amphoteric nature of Al₂O₃ and Al(OH)₃. Write all relevant equations. Compare the oxide character of Group 13 oxides from B to Tl, and explain how this reflects the trend in metallic character. Also compare with the trend in oxide character across Period 3. [8]
Amphoteric Al₂O [2]:
With acid: Al₂O + 6HCl → 2AlCl + 3H₂O (acts as base, accepts H⁺)
Al₂O + 3H₂SO → Al₂(SO) + 3H₂O
With base: Al₂O + 2NaOH + 3H₂O → 2Na[Al(OH)₄] (acts as acid, reacts with OH⁻)
Amphoteric Al(OH) [2]:
With acid: Al(OH) + 3HCl → AlCl + 3H₂O
With base (excess): Al(OH) + OH⁻ → [Al(OH)₄]⁻
Trend in Group 13 oxide character [2]: B₂O (acidic) → Al₂O, Ga₂O (amphoteric) → In₂O (weakly basic) → Tl₂O (basic). This reflects the increase in metallic character down the group: B is a non-metal/metalloid (non-metallic → acidic oxide); Al/Ga are metals with amphoteric behaviour (borderline metallic); In/Tl are fully metallic (metallic → basic oxide). A general rule: non-metallic oxides are acidic; metallic oxides are basic; transition zone gives amphoteric.
Comparison with Period 3 [2]: Across Period 3 from left to right: Na₂O (strongly basic) → MgO (basic) → Al₂O (amphoteric) → SiO₂ (weakly acidic) → P₂O₅ (acidic) → SO (strongly acidic). Metallic character decreases across the period → oxide character changes from basic to acidic. Al₂O sits at the intersection, being amphoteric. The pattern across a period mirrors the trend down Group 13 in reverse: both relate to the balance between metallic/non-metallic character.
With acid: Al₂O + 6HCl → 2AlCl + 3H₂O (acts as base, accepts H⁺)
Al₂O + 3H₂SO → Al₂(SO) + 3H₂O
With base: Al₂O + 2NaOH + 3H₂O → 2Na[Al(OH)₄] (acts as acid, reacts with OH⁻)
Amphoteric Al(OH) [2]:
With acid: Al(OH) + 3HCl → AlCl + 3H₂O
With base (excess): Al(OH) + OH⁻ → [Al(OH)₄]⁻
Trend in Group 13 oxide character [2]: B₂O (acidic) → Al₂O, Ga₂O (amphoteric) → In₂O (weakly basic) → Tl₂O (basic). This reflects the increase in metallic character down the group: B is a non-metal/metalloid (non-metallic → acidic oxide); Al/Ga are metals with amphoteric behaviour (borderline metallic); In/Tl are fully metallic (metallic → basic oxide). A general rule: non-metallic oxides are acidic; metallic oxides are basic; transition zone gives amphoteric.
Comparison with Period 3 [2]: Across Period 3 from left to right: Na₂O (strongly basic) → MgO (basic) → Al₂O (amphoteric) → SiO₂ (weakly acidic) → P₂O₅ (acidic) → SO (strongly acidic). Metallic character decreases across the period → oxide character changes from basic to acidic. Al₂O sits at the intersection, being amphoteric. The pattern across a period mirrors the trend down Group 13 in reverse: both relate to the balance between metallic/non-metallic character.
Q8 [6 marks]
The industrial extraction of aluminium uses the Hall-Héroult process. Although this involves electrolysis (covered in later units), explain why Al₂O₃ (not AlCl) is used as the electrolyte; describe why cryolite (NaAlF⁶) is added; and explain the advantage of the protective oxide layer on Al in terms of industrial applications. [6]
Why Al₂O not AlCl [2]: AlCl is covalent → it does not conduct electricity as a liquid (no free ions). Al₂O is ionic → provides Al³⁺ and O²⁻ ions when molten → conducts electricity and can be electrolysed. AlCl also decomposes on heating before melting (sublimes at ~180°C). Al₂O (m.p. 2050°C) must be molten for electrolysis.
Role of cryolite (NaAlF⁶) [2]: Pure Al₂O melts at 2050°C — impractically high for industrial electrolysis. Dissolving Al₂O in molten cryolite (NaAlF⁶) reduces the operating temperature to ~950°C, making the process economically and practically feasible. Cryolite also improves the electrical conductivity of the electrolyte. The cryolite does not get consumed — it acts as a solvent.
Industrial advantage of oxide layer [2]: In everyday use, the self-forming Al₂O layer makes Al resistant to corrosion without any additional coating → no painting or galvanising needed (unlike iron). This makes Al ideal for: aircraft bodies (corrosion-resistant, lightweight); food packaging (no contamination from corrosion); building materials (low maintenance); drink cans; marine applications. Anodising (electrochemical thickening of the oxide layer) further increases corrosion resistance and allows colouring → architectural applications.
Role of cryolite (NaAlF⁶) [2]: Pure Al₂O melts at 2050°C — impractically high for industrial electrolysis. Dissolving Al₂O in molten cryolite (NaAlF⁶) reduces the operating temperature to ~950°C, making the process economically and practically feasible. Cryolite also improves the electrical conductivity of the electrolyte. The cryolite does not get consumed — it acts as a solvent.
Industrial advantage of oxide layer [2]: In everyday use, the self-forming Al₂O layer makes Al resistant to corrosion without any additional coating → no painting or galvanising needed (unlike iron). This makes Al ideal for: aircraft bodies (corrosion-resistant, lightweight); food packaging (no contamination from corrosion); building materials (low maintenance); drink cans; marine applications. Anodising (electrochemical thickening of the oxide layer) further increases corrosion resistance and allows colouring → architectural applications.
Q9 [8 marks]
A student adds NaOH solution dropwise to four separate test solutions labelled A, B, C, D. Results: A gives green precipitate insoluble in excess NaOH; B gives white precipitate soluble in excess NaOH; C gives blue precipitate insoluble in excess NaOH; D gives white precipitate insoluble in excess NaOH. (a) Identify the ions in A, B, C, D. [4] (b) Write ionic equations for the formation of each precipitate. [4]
(a) Identifications [4]:
A: Green precipitate insoluble in excess NaOH → Fe²⁺. Fe(OH)₂ is green; insoluble in excess NaOH. (Note: Fe(OH)₂ may turn brown in air as it oxidises to Fe(OH).) [1]
B: White precipitate soluble in excess NaOH → Al³⁺ (or Zn²⁺). Al(OH) is amphoteric, dissolves in excess NaOH to give [Al(OH)₄]⁻. [1]
C: Blue precipitate insoluble in excess NaOH → Cu²⁺. Cu(OH)₂ is pale blue; insoluble in excess NaOH. [1]
D: White precipitate insoluble in excess NaOH → Mg²⁺. Mg(OH)₂ is white gelatinous; not amphoteric, does not dissolve in excess NaOH. [1]
(b) Ionic equations [4]:
A: Fe²⁺(aq) + 2OH⁻(aq) → Fe(OH)₂(s) [green] [1]
B: Al³⁺(aq) + 3OH⁻(aq) → Al(OH)₃(s) then Al(OH)₃(s) + OH⁻(aq) → [Al(OH)₄]⁻(aq) [1]
C: Cu²⁺(aq) + 2OH⁻(aq) → Cu(OH)₂(s) [blue] [1]
D: Mg²⁺(aq) + 2OH⁻(aq) → Mg(OH)₂(s) [white; does not dissolve in excess NaOH] [1]
A: Green precipitate insoluble in excess NaOH → Fe²⁺. Fe(OH)₂ is green; insoluble in excess NaOH. (Note: Fe(OH)₂ may turn brown in air as it oxidises to Fe(OH).) [1]
B: White precipitate soluble in excess NaOH → Al³⁺ (or Zn²⁺). Al(OH) is amphoteric, dissolves in excess NaOH to give [Al(OH)₄]⁻. [1]
C: Blue precipitate insoluble in excess NaOH → Cu²⁺. Cu(OH)₂ is pale blue; insoluble in excess NaOH. [1]
D: White precipitate insoluble in excess NaOH → Mg²⁺. Mg(OH)₂ is white gelatinous; not amphoteric, does not dissolve in excess NaOH. [1]
(b) Ionic equations [4]:
A: Fe²⁺(aq) + 2OH⁻(aq) → Fe(OH)₂(s) [green] [1]
B: Al³⁺(aq) + 3OH⁻(aq) → Al(OH)₃(s) then Al(OH)₃(s) + OH⁻(aq) → [Al(OH)₄]⁻(aq) [1]
C: Cu²⁺(aq) + 2OH⁻(aq) → Cu(OH)₂(s) [blue] [1]
D: Mg²⁺(aq) + 2OH⁻(aq) → Mg(OH)₂(s) [white; does not dissolve in excess NaOH] [1]