Doc Brown's Chemistry  - Advanced Level Inorganic Chemistry Periodic Table Revision Notes

Part 8. The p–block elements: 8.1 Group 3/13 Boron and Aluminium in particular

(1) Group 3/13 Position in the periodic table - introduction, data, trends and electron configurations

Elect. potential = standard electrode potential at 298 K, 1M solution concentration

Some general comments and trends for group 3/13 elements of the periodic table

• Generally speaking down a p block group the element becomes more metallic, but boron is the only true non–metal, the rest are basically metals with a some non–metallic chemical character.

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(2) BORON – brief summary of a few points of its chemistry

• The structure of the element: boron

  • Non–metal existing as a giant covalent lattice, B_n, where n is an extremely large number.

• Physical properties of boron: 

  • Hard high melting solid; mpt 2300^oC; bpt 3659^oC;  poor conductor of heat/electricity.

• Group, electron configuration (and oxidation states) of boron: 

  • Gp3; e.c. 2,3 or 1s²2s²2p¹;  (+3 only) e.g. B₂O₃ and BCl₃ etc.

• Reaction of element boron with oxygen: 

  • Reacts when heated strongly in air to form boron oxide which has a giant covalent structure.

    • 4B_(s) + 3O_2(g) ==> 2B₂O_3(s)  

• Reaction of boron oxide with water:

  • Insoluble, no reaction but it is a weakly acidic oxide.

• Reaction of boron oxide with acids:

  • None, only acidic in acid–base behaviour.

• Reaction of boron oxide with strong bases/alkalis:

  • Presumably dissolves to give a solution of sodium borate.

• Reaction of boron with chlorine: 

  • Forms covalent liquid boron trichloride on heating in chlorine gas.

    • 2B_(s) + 3Cl_2(g) ==> 2BCl_3(l) 

• Reaction of boron chloride with water:

  • It hydrolyses to form boric acid and hydrochloric acid.

    • BCl_3(l) + 3H₂O_(l) ==> B(OH)_3(aq)^* + 3HCl_(aq) 

    • ^* can also be, but less accurately, written as H₃BO₃ 

• Reaction of boron with water:

  • None.

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(3) Some molecule/ion shapes and bond angles involving boron

Three bond pairs of electrons gives TRIGONAL PLANAR shape. The Q–X–Q bond angle is exactly 120o e.g. for gaseous boron hydride BH3 (X = B, Q = H). The borohydride (tetrahydroborate(III)) ion, BH4-, will be tetrahedral with a H-B-H bond angle of 109.5o.

Three bond pairs of electrons gives TRIGONAL PLANAR shape. The Q–X–Q bond angle is exactly 120o e.g. for gaseous boron trifluoride BF3 (Q = F, Cl and X = B)

H3N:=>BF3 Boron trifluoride (3 bonding pairs, 6 outer electrons) acts as a lone pair acceptor (Lewis acid) and ammonia (3 bond pairs) and lone pair which enables it to act as a Lewis base – a an electron pair donor. It donates the lone pair to the 4th 'vacant' boron orbital to form a sort of 'adduct' compound. Its shape is essentially the same as ethane, a sort of double tetrahedral with H–N–H, N–B–F and F–B–F bond angles of ~109o.

(4) Boron compound reducing agents in organic chemistry

• Derivatives of boron hydride are useful reducing agents in organic chemistry.

  • All the reduction reactions are shown as simplified equations.

• Sodium tetrahydrioborate(III), NaBH₄ (sodium borohydride) reduces aldehydes to primary alcohols and ketones to secondary alcohols.

• These reactions are essentially the reduction of the carbony1 group >C=O to >CHOH.

• The reaction can be carried out in water. The reduction mechanism is very complicated, but can be considered in a simplistic way as involving the donation of a hydride ion to the aldehyde/ketone.

  • aldehyde: RCHO + 2[H] ==> RCH₂OH (R = H, alkyl or aryl)

  • ketone: R₂C=O + 2[H] ==> R₂CHOH (R = alkyl or aryl)

• NaBH₄, is not a powerful enough reducing agent to reduce carboxylic acids to a primary aliphatic alcohol.

• NaBH₄, is not a powerful enough reducing agent to reduce nitro–aromatic compounds to primary aromatic amines.

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(5) ALUMINIUM – Summary of some important points about its chemistry

• The structure of the element aluminium:

  • Giant lattice metallic structure of immobile positive metal ions surrounded by a 'sea' of freely moving mobile electrons (so–called delocalised electrons).

• Physical properties of aluminium: 

  • Moderately hard high melting solid; mpt 661^oC; bpt 2467^oC;  good conductor heat/electricity.

• Group, electron configuration (and oxidation states) of aluminium: 

  • Gp3; e.c. 2,8,3 or 1s²2s²2p⁶3s²3p¹;  (+3 only) e.g. Al₂O₃ and AlCl₃.

• Reaction of aluminium with oxygen: 

  • Reacts when heated strongly in air to form a white powder of aluminium oxide which has a giant ionic structure, (Al³⁺)₂(O^2–)₃.

    • 4Al_(s) + 3O_2(g) ==> 2Al₂O_3(s)  

    • The above reaction occurs very rapidly on a freshly cut aluminium surface, but the microscopic oxide layer inhibits any further reaction, giving aluminium a 'lower reactivity' than expected, and its excellent anti–corrosion properties.

• Reaction of aluminium oxide with water:

  • Insoluble, no reaction but it is an amphoteric oxide and forms salts with both acids and alkali (see below).

• Reaction of aluminium oxide with acids:

  • It behaves as a basic oxide dissolving to form the chloride, sulphate and nitrate salt in the relevant dilute acid.

  • Al₂O_3(s) + 6HCl_(aq) ==> 2AlCl_3(aq) + 3H₂O_(l) 

  • Al₂O_3(s) + 3H₂SO_4(aq) ==> Al₂(SO₄)_3(aq) + 3H₂O_(l) 

  • Al₂O_3(s) + 6HNO_3(aq) ==> 2Al(NO₃)_3(aq) + 3H₂O_(l) 

  • ionic equation: Al₂O_3(s) + 6H⁺_(aq) ==> 2Al³⁺_(aq) + 3H₂O_(l) 

• Reaction of aluminium oxide with strong bases/alkalis:

  • The oxide also behaves as an acidic oxide by dissolving in strong soluble bases to form aluminate(III) salts.

  • e.g. Al₂O_3(s) + 2NaOH_(aq) + 3H₂O_(l) ==> 2Na[Al(OH)₄]_(aq) 

  • forming sodium aluminate(III) with sodium hydroxide.

  • ionic equation: Al₂O_3(s) + 2OH^–_(aq) + 3H₂O_(l) ==> 2[Al(OH)₄]^–_(aq) 

  • Therefore aluminium oxide is an amphoteric oxide, because of this dual acid–base behaviour.

• Reaction of aluminium with chlorine: 

  • Burns when heated strongly in chlorine gas to form the white^* solid aluminium chloride on heating in chlorine gas.

    • 2Al_(s) + 3Cl_2(g) ==> 2AlCl_3(s)

    • ^* It is often a faint yellow in colour, due to traces of iron forming iron(III) chloride.

    • 

    • Aluminium chloride is a curious substance in its behaviour. The solid, AlCl₃, consists of an ionic lattice of Al³⁺ ions, each surrounded by six Cl^– ions, BUT on heating, at about 180^oC, the thermal kinetic energy of vibration of the ions in the lattice is sufficient to cause it break down and sublimation takes place (s ==> g). In doing so the co–ordination number of the aluminium changes from six to four to form the readily vapourised covalent dimer molecule, Al₂Cl₆, shown above.

    • The reaction with bromine is similar and forms the covalent dimer molecule Al₂Br₆ directly

    • For more see reactions of halogens with aluminium and iron

• Reaction of aluminium with water:

  • None due to protective oxide layer.

• Reactions of the hexa–aqua aluminium ion:

  • It gives a gelatinous white precipitate with sodium hydroxide or ammonia solution which displays amphoteric behaviour by dissolving in excess strong alkali (NaOH_(aq), NOT NH_3(aq)) and acids.

    • Al³⁺_(aq) + 3OH^–_(aq) ==> Al(OH)_3(s) 

    • or [Al(H₂O)₆]³⁺_(aq)  + 3OH^–_(aq) ==> [Al(OH)₃(H₂O)₃] + 3H₂O_(l) 

      • The hydroxide readily dissolves in acids to form salts:

      • Al(OH)_3(s) + 3H⁺_(aq) ==> Al³⁺_(aq) + 3H₂O_(l) 

        • or more elaborately: [Al(OH)₃(H₂O)₃] + 3H₃O⁺_(aq) [Al(H₂O)₆]³⁺_(aq)  + 3H₂O_(l) 

        • Thus showing amphoteric behaviour, since the hydroxide ppt. also dissolves in excess strong alkali (below).

    • [Al(H₂O)₆]³⁺_(aq) + 6OH^–_(aq) ==> [Al(OH)₆]^3–_(aq) + 6H₂O_(l)  (from original aqueous ion)

      • or [Al(OH)₃(H₂O)₃]_(s) + 3OH^–_(aq) ==> [Al(OH)₆]^3–_(aq) + 3H₂O_(l) (from hydroxide ppt.)

      • or more simply Al(OH)_3(s) + 3OH^–_(aq) ==> [Al(OH)₆]^3–_(aq)  (from hydroxide ppt.)

  • With aqueous sodium carbonate solution, the hydroxide ppt. is formed, and, because of its acidic nature, bubbles of carbon dioxide gas are evolved.

    • 2[Al(H₂O)₆]³⁺_(aq) + CO₃^2–_(aq) 2[Al(H₂O)₅(OH)]²⁺_(aq) + H₂O_(l) + CO_2(g)   

    • this process of proton donation continues until the gelatinous ppt. [Al(OH)₃(H₂O)₃]_(s) is formed, but will not dissolve in excess of the weak base/alkali.

    • Sodium carbonate is not a strong enough base–alkali to dissolve the aluminium hydroxide precipitate.

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(6) The extraction of aluminium

Aluminium is obtained from mining the mineral bauxite.

The purified bauxite ore of aluminium oxide is continuously fed in. Cryolite is added to lower the melting point and dissolve the ore.

Ions must be free to move to the electrode connections called the cathode (–, negative), attracting positive ions e.g. Al³⁺, and the anode (+, positive) which attracts negative ions e.g. O^2–.

When the d.c. current is passed through aluminium forms at the negative cathode (metal*) and sinks to the bottom of the tank.

At the positive anode, oxygen gas is formed (non–metal*). This is quite a problem. At the high temperature of the electrolysis cell it burns and oxidises away the carbon electrodes to form toxic carbon monoxide or carbon dioxide. So the electrode is regularly replaced and the waste gases dealt with! 

It is a costly process (6x more than Fe!) due to the large quantities of expensive electrical energy needed for the process.

* Two general rules:

• Metals and hydrogen (from positive ions), form at the negative cathode electrode.

• Non–metals (from negative ions), form at the positive anode electrode.

Raw materials for the electrolysis process:

• Bauxite ore of impure aluminium oxide [Al₂O₃ made up of Al³⁺ and O^2– ions]

• Carbon (graphite) for the electrodes.

• Cryolite reduces the melting point of the ore and saves energy, because the ions must be free to move to carry the current

• Electrolysis means using d.c. electrical energy to bring about chemical changes e.g. decomposition of a compound to form metal deposits or release gases. The electrical energy splits the compound!

• At the electrolyte connections called the anode electrode (+, attracts – ions) and the cathode electrode (–, attracts + ions). An electrolyte is a conducting melt or solution of freely moving ions which carry the charge of the electric current.

• At the negative  (–) cathode, reduction occurs (electron gain) when the positive aluminium ions are attracted to it. They gain three electrons to change to neutral Al atoms.

  • Al³⁺ + 3e^– ==> Al

• At the positive (+) anode, oxidation takes place (electron loss) when the negative oxide ions are attracted to it. They lose two electrons forming neutral oxygen molecules.

  • 2O^2– ==> O₂ + 4e^– 

  • or 2O^2– – 4e^– ==> O₂ 

• Note: Reduction and Oxidation always go together!

• The overall electrolytic decomposition is ...

  • aluminium oxide => aluminium + oxygen

  • 2Al₂O₃ ==> 4Al + 3O₂

  • and is a very endothermic process, lots of electrical energy input!

• GENERAL NOTE ON ELECTROLYSIS:

• Any molten or dissolved material in which the liquid contains free moving ions is called the electrolyte.

• Ions are charged particles e.g. Na⁺ sodium ion, or Cl^– chloride ion, and their movement or flow constitutes an electric current, because a current is moving charged particles.

• What does the complete electrical circuit consist of?

  • There are two ion currents in the electrolyte flowing in opposite directions:

    • positive cations e.g. Al³⁺ attracted to the negative cathode electrode,

    • and negative anions e.g. O^2– attracted to the positive anode electrode,

    • BUT remember no electrons flow in the electrolyte, only in the graphite or metal wiring!

  • The circuit of 'charge flow' is completed by the electrons moving around the external circuit e.g. copper wire or graphite electrode, from the positive to the negative electrode

  • This e^– flow from +ve to –ve electrode perhaps doesn't make sense until you look at the electrode reactions, electrons released at the +ve anode move round the external circuit to produce the electron rich negative cathode electrode.

• Electron balancing: In the above process it takes the removal of four electrons from two oxide ions to form one oxygen molecule and the gain of three electrons by each aluminium ion to form one aluminium atom. Therefore for every 12 electrons you get 3 oxygen molecules and 4 aluminium atoms formed.

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(7) The chemical properties and uses of aluminium

• Aluminium can be made more resistant to corrosion by a process called anodising. Iron can be made more useful by mixing it with other substances to make various types of steel. Many metals can be given a coating of a different metal to protect them or to improve their appearance.

  • Aluminium is a reactive metal but it is resistant to corrosion. This is because aluminium reacts in air to form a layer of aluminium oxide which then protects the aluminium from further attack.

    • This is why it appears to be less reactive than its position in the reactivity series of metals would predict.

  • For some uses of aluminium it is desirable to increase artificially the thickness of the protective oxide layer in a process is called anodising.

    • This involves removing the oxide layer by treating the aluminium sheet with sodium hydroxide solution.

    • The aluminium is then placed in dilute sulphuric acid and is made the positive electrode (anode) used in the electrolysis of the acid.

    • Oxygen forms on the surface of the aluminium and reacts with the aluminium metal to form a thicker protective oxide layer. 

  • Aluminium can be alloyed to make 'Duralumin' by adding copper (and smaller amounts of magnesium, silicon and iron), to make a stronger alloy used in aircraft components (low density = 'lighter'!), greenhouse and window frames (good anti–corrosion properties), overhead power lines (quite a good conductor and 'light'), but steel strands are included to make the 'line' stronger and poorly electrical conducting ceramic materials are used to insulate the wires from the pylons and the ground.

    • There is a note about structure of metal alloys on the metallic bonding page.

• More on the reactions of aluminium

  • Reaction with aluminium with chlorine

  • 

  • If dry chlorine gas Cl₂ is passed over heated iron or aluminium, the chloride is produced. The experiment (shown above) should be done very carefully by the teacher in a fume cupboard.

    • 2Al_(s) + 3Cl_2(g) ==> 2AlCl_3(s)

    • The aluminium can burn intensely with a violet flame, white fumes of aluminium chloride sublime from the hot reacted aluminium and the white solid forms on the cold surface of the flask (its often discoloured yellow from the trace chlorides of copper or iron that may be formed).

      • 2Fe_(s) + 3Cl_2(g) ==> 2FeCl_3(s)

    • Aluminium chloride reacts exothermically as it is hydrolysed by water to give the metal hydroxide and fumes of hydrogen chloride, and so dry conditions are needed.

    • Aluminium chloride cannot be made in an anhydrous form from aqueous solution neutralisation. This is because on evaporation the compounds contain 'water of crystallisation'. On heating the hydrated salt it hydrolyses and decomposes into water, the oxide or hydroxide and fumes of hydrogen chloride.

    • Reaction of chloride with water:

      • With a little water it rapidly, and exothermically hydrolyses to form aluminium hydroxide and nasty fumes of hydrogen chloride gas.

        • AlCl_3(s) + 3H₂O_(l) ==> Al(OH)_3(s) + 3HCl_(g) 

      • However, if a large excess of water is rapidly added, a weakly acidic solution of aluminium chloride is formed, with the minimum of nasty fumes!

        • AlCl_3(s) + aq ==> Al³⁺_(aq) + 3Cl^–_(aq) 

        • or more correctly: AlCl_3(s) + 6H₂O_(l) ==> [Al(H₂O)₆]³⁺_(aq) + 3Cl^–_(aq) 

      • The solution is slightly acidic, because the hexa–aqua aluminium ion can donate a proton to a water molecule forming the oxonium ion.

        • [Al(H₂O)₆]³⁺_(aq) + H₂O_(l) [Al(H₂O)₅OH]²⁺_(aq) + H₃O⁺_(aq) 

  • The surface of aluminium goes white when strongly heated in air/oxygen to form white solid aluminium oxide. Theoretically its quite a reactive metal but an oxide layer is readily formed even at room temperature and this has quite an inhibiting effect on its reactivity. Even when scratched, the oxide layer rapidly reforms, which is why it appears to be less reactive than its position in the reactivity series of metals would predict but the oxide layer is so thin it is transparent,  so aluminium surfaces look metallic and not a white matt surface.
    • aluminium + oxygen ==> aluminium oxide
    • 4Al_(s) + 3O_2(g) ==> 2Al₂O_3(s)
    • Under 'normal circumstances' in the school laboratory aluminium has virtually no reaction with water, not even when heated in steam due to a protective aluminium oxide layer of Al₂O₃. (see above) The metal chromium behaves chemically in the same way, forming a protective layer of chromium(III) oxide, Cr₂O₃, and hence its anti–corrosion properties when used in stainless steels and chromium plating. Although this again illustrates the 'under–reactivity' of aluminium, the Thermit Reaction shows its rightful place in the reactivity series of metals.
  • The Thermit reaction: However the true reactivity of aluminium can be spectacularly seen when its grey powder is mixed with brown iron(III) oxide powder. When the mixture is ignited with a magnesium fuse (needed because of the very high activation energy!), it burns very exothermically in a shower of sparks to leave a red hot blob of molten=>solid iron and white aluminium oxide powder. Note the high temperature of the magnesium fuse flame is so high, the oxide layer (to the delight of all pupils) fails to inhibit the displacement reaction! yippee!
    • aluminium + iron(III) oxide ==> iron + aluminium oxide
      • aluminium + iron(III) oxide ==>  aluminium oxide + iron
      • 2Al_(s) + Fe₂O_3(s) ==> Al₂O_3(s) + 2Fe_(s)
    • This is a typical displacement reaction by a more reactive metal displacing a less reactive metal from one of its compounds.
  • Slow reaction with dilute hydrochloric acid to form the colourless soluble salt aluminium chloride and hydrogen gas.
    • aluminium + hydrochloric acid ==> aluminium chloride + hydrogen
    • 2Al_(s) + 6HCl_(aq) ==> 2AlCl_3(aq) + 3H_2(g)
  • The reaction with dilute sulphuric acid is very slow to form colourless  aluminium sulphate and hydrogen.
    • aluminium + sulphuric acid ==> aluminium sulphate + hydrogen
    • 2Al_(s) + 3H₂SO_4(aq) ==> Al₂(SO₄)_3(aq) + 3H_2(g)
  • If the surface of aluminium is treated with less reactive metal salt, it is still possible to get a displacement reaction. Check this out by leaving a piece of aluminium foil in copper(II) sulphate solution and a patchy pink colour of copper metal slowly appears over many hours/days?
    • aluminium + copper(II) sulphate ==> aluminium sulphate + copper
    • 2Al(s) + 3CuSO₄(aq) ==> Al₂(SO₄)₃(aq) + 3Cu(s)
    • ionic redox equation: 2Al(s) + 3Cu²⁺(aq) ==> 2Al³⁺(aq) + 3Cu(s)

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(8) Amphoteric nature of aluminium hydroxide and the acidity of the hexaaquaaluminium ion

• The addition of limited amounts of the bases sodium hydroxide or ammonia solution to an aluminium salt solution.

  • [Al(H₂O)₆]³⁺_(aq) + 3OH^–_(aq) ==> [Al(H₂O)₃(OH)₃]_(s) + 3H₂O_(aq)

  • A white gelatinous precipitate of aluminium hydroxide is formed.

    • Simplified equation: Al³⁺_(aq) + 3OH^–_(aq) ==> Al(OH)_3(s)

• The further addition of excess sodium hydroxide or ammonia solution.

  • With excess ammonia there is no effect, but with excess sodium hydroxide the aluminium hydroxide dissolves to form a soluble aluminate complex anion – therefore exhibiting amphoteric behaviour. since the hydroxide will also dissolve in acids (paragraph below NaOH equation).

  • [Al(H₂O)₃(OH)₃]_(s) + 3OH^–_(aq) ==> *[Al(OH)₆]^3–_(aq) + 3H₂O_(aq)

    • Simplified equation: Al(OH)_3(s) + 3OH^–_(aq) ==> *[Al(OH)₆]^3–_(aq)

    • *The products will be an equilibrium mixture including [Al(H₂O)₂(OH)₄]^–_(aq) and [Al(H₂O)(OH)₅]^2–_(aq) too. You could write the equation in terms of forming these species too and any of the three possibilities should get you the marks.

  • To complete the 'amphoteric' picture of aluminium hydroxide we consider it dissolving in mineral acids to form typical salts e.g. aluminium chloride, aluminium nitrate and aluminium sulphate.

    • Al(OH)_3(s) + 3HCl_(aq) ==> AlCl_3(aq) + 3H₂O_(l)

    • Al(OH)_3(s) + 3HNO_3(aq) ==> Al(NO₃)_3(aq) + 3H₂O_(l)

    • 2Al(OH)_3(s) + 3H₂SO_4(aq) ==> Al₂(SO₄)_3(aq) + 6H₂O_(l)

• The addition of sodium carbonate solution to an aluminium salt solution.

  • Bubbles of carbon dioxide and a white gelatinous precipitate of aluminium hydroxide are formed.

    • 2[Al(H₂O)₆]³⁺_(aq) + 3CO₃^2–_(aq) ==> 2[Al(H₂O)₃(OH)₃]_(s) + 3CO_2(g) + 3H₂O_(aq)

    • There several equation 'permutations' to represent this quite complicated reaction, so I've just composed one that shows the formation of both observed products. Since sodium carbonate solution is alkaline you can legitimately write a hydroxide ppt. equation as for sodium hydroxide above but it doesn't show the formation of carbon dioxide.

      • You can write an equation to show the formation of carbon dioxide leaving a soluble cationic complex of aluminium in solution and this equation fits in well with the acid–base nature of this reaction.

        • [Al(H₂O)₆]³⁺_(aq) + CO₃^2–_(aq) ==> 2[Al(H₂O)₄(OH)₂]⁺_(aq) + CO_2(g) + 3H₂O_(aq)

        • This equation shows the hexaaquaaluminium ion acting as a Bronsted–Lowry acid donating two protons to the carbonate ion (B–L base) to form carbon dioxide and water.

    • This reaction shows why 'aluminium carbonate' 'Al₂(CO₃)₃' cannot exist. The hydrated highly charged central metal ion is too acidic to co–exist with a carbonate ion. The same situation applies to the chromium(III) Cr³⁺ and iron(III) Fe³⁺ ions i.e. no chromium(III) carbonate or iron(III) carbonate exists. However with a lesser charged, lesser acidic ion, carbonates can exist, so there is an iron(II) carbonate FeCO₃.

  • Aluminium salt solutions are slightly acidic for the same reasons as the carbonate reaction – namely the acidity of the hexaaquaaluminium ion i.e. a acting as a proton donor.

  • [Al(H₂O)₆]³⁺_(aq) + H₂O_(l) [Al(H₂O)₅(OH)]²⁺_(aq) + H₃O⁺_(aq) 

• The addition of excess sodium carbonate solution has no further effect. Sodium carbonate is too weak a base to effect the amphoteric nature of aluminium hydroxide and dissolve the aluminium hydroxide precipitate.

  • For strong alkalis like sodium hydroxide the whole sequence of each theoretical step of aluminium hydroxide precipitation and its subsequent dissolving in strong base–alkali is shown the series of diagrams below.

  • All are, for simplicity, treated as octahedral complexes of 6 ligands – either water H2O or hydroxide ion OH–.

  • [Al(H₂O)₆]³⁺ => [Al(OH)(H₂O)₅]²⁺ => [Al(OH)₂(H₂O)₄]⁺ => [Al(OH)₃(H₂O)₃]_(s) precipitate

  • dissolving => [Al(OH)₄(H₂O)₃]^– => [Al(OH)₅(H₂O)]^2– => [Al(OH)₆]^3–

			The sequence of aluminium hydroxide precipitate formation and its subsequent dissolving in excess strong alkali. Each step is essentially one of proton removal from each complex (from 3+ to 3–).

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(9) Aluminium compound reducing agents in organic chemistry

• Lithium tetrahydridoaluminate(III), LiAlH₄ (lithium tetrahydride) reduces aldehydes to primary alcohols and ketones to secondary alcohols.

• LiAlH₄ is a more powerful reducing agent than NaBH₄ and reacts violently with water, so the reaction must be carried out in an inert solvent like ethoxyethane ('ether'). The initial product is hydrolysed by dil. sulphuric acid.

  • aldehyde: RCHO + 2[H] ==> RCH₂OH (R = H, alkyl or aryl)

  • ketone: R₂C=O + 2[H] ==> R₂CHOH (R = alkyl or aryl)

• LiAlH₄ is a more powerful reducing agent than NaBH₄, and in ether solvent, readily reduces carboxylic acids to primary alcohols. The reaction can be summarised as:

  • RCOOH + 4[H] ==> RCH₂OH + H₂O (R = H, alkyl or aryl)

• LiAlH₄ is a more powerful reducing agent than NaBH₄ and in ether solvent will reduce nitriles to primary aliphatic amines.

  • RCN + 4[H] ==> RCH₂NH₂ (R = H, alkyl or aryl)

• LiAlH₄ is a more powerful reducing agent than NaBH₄ and in ether solvent readily reduces nitro–aromatics to primary aromatic amines.

  • C₆H₅NO₂ + 6[H] ==> C₆H₅NH₂ + 2H₂O

• methylnitrobenzenes would be reduced to methylphenylamine primary amines, i.e.

  • CH₃C₆H₄NO₂ + 6[H] ==> CH₃C₆H₄NH₂ + 2H₂O

• as will any aromatic compound with a nitro group (–NO₂) directly attached to a benzene ring.

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