a level notes on gp3 boron and aluminium

7/27/2019 A Level Notes on Gp3 Boron and Aluminium

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Group 3/13 Introduction

down group 3/13 ===>

property\Zsymbol, name 5B Boron 13Al Aluminium 31Ga Gallium 49In Indium 81Tl Thallium

Period 2 3 4 5 6

Appearance (RTP) brown solid silvery solid silvery solid silvery solid silvery solid

melting pt./oC 2300 661 30 156 304

boiling pt./oC

3659

2467

2400

2080

1457

density/gcm-3 2.3 2.7 5.9 7.3 11.9

relative electrical conductivity


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Electrode potential E

M(s)/M3+

(aq) na -1.66V -0.53V -0.34V +0.72V

Electrode potential E

M+

(aq)/M3+

(aq)

na na na -0.44V +1.25V

Pauling electronegativity 2.04 1.61 1.81 1.76 1.80

simple electron config. 2,3 2,8,3 2,8,18,3 2,8,18,18,3 2,8,18,32,18,3

electron configuration [He]2s22p

1 [Ne]3s

23p

1 [Ar]3d

104s

24p

1 [Kr]4d

105s

25p

1 [Xe]4f

145d

106s

26p

1

principal oxidation states +3 +3 +1, +3 +1, +3 +1, +3

property\Zsymbol, name 5B Boron 13Al Aluminium 31Ga Gallium 49In Indium 81Tl Thallium

Some general group trends

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.

BORON - brief summary of a few points

The structure of the element:
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o Non-metal existing as a giant covalent lattice, Bn, where n is an extremely large number.

Physical properties:

o Hard high melting solid; mpt 2300oC; bpt 3659

oC; poor conductor heat/electricity.

Group, electron configuration (and oxidation states):

o Gp3; e.c. 2,3 or1s22s

22p

1; (+3 only) e.g. B2O3 and BCl3 etc.

Reaction of element with oxygen:

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

4B(s) + 3O2(g) ==> 2B2O3(s)

Reaction of oxide with water:

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

Reaction of oxide with acids:

o None, only acidic in acid-base behaviour.

Reaction of oxide with strong bases/alkalis:

o Presumably dissolves to give a solution of sodium borate.

Reaction of element with chlorine:

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

2B(s) + 3Cl2(g) ==> 2BCl3(l)

Reaction of chloride with water:

o It hydrolyses to form boric acid and hydrochloric acid.

BCl3(l) + 3H2O(l) ==> B(OH)3(aq)*+ 3HCl(aq)

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

Reaction of element with water:

o None.


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Some molecule shapes and bond angles

Three bond pairs of electrons gives TRIGONAL PLANAR shape. The Q-X-Q bond angle is exactly 120oe.g.

for gaseous boron hydride BH3 (X = B, Q = H).

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:=>BF3Boron 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.

-
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Boron compound reducing agents in organic chemistry

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

o 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.

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

o 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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ALUMINIUM - brief summary of a few points

The structure of the element:

o Giant lattice metallic structure of immobile positive metal ions surrounded by a 'sea' of freely moving mobile electrons (so-called

delocalised electrons).

Physical properties:

o Moderately hard high melting solid; mpt 661oC; bpt 2467

oC; good conductor heat/electricity.

Group, electron configuration (and oxidation states):

o Gp3; e.c. 2,8,3 or1s

2

2s

2

2p

6

3s

2

3p

1

; (+3 only) e.g. Al2O3 and AlCl3.

Reaction of element with oxygen:

o Reacts when heated strongly in air to form a white powder ofaluminium oxide which has a giant ionic structure, (Al3+

)2(O2-

)3.

4Al(s) + 3O2(g)==> 2Al2O3(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 oxide with water:

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

Reaction of oxide with acids:

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

o Al2O3(s) + 6HCl(aq)==> 2AlCl3(aq) + 3H2O(l)

o Al2O3(s) + 3H2SO4(aq)==> Al2(SO4)3(aq) + 3H2O(l)

o Al2O3(s) + 6HNO3(aq)==> 2Al(NO3)3(aq) + 3H2O(l)

o ionic equation: Al2O3(s) + 6H+

(aq)==> 2Al3+

(aq) + 3H2O(l)


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Reaction of oxide with strong bases/alkalis:

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

o e.g. Al2O3(s) + 2NaOH(aq) + 3H2O(l)==> 2Na[Al(OH)4](aq)

o forming sodium aluminate(III) with sodium hydroxide.

o ionic equation: Al2O3(s) + 2OH-(aq) + 3H2O(l)==> 2[Al(OH)4]

-(aq)

o Therefore aluminium oxide is an amphoteric oxide, because of this dual acid-base behaviour.

Reaction of element with chlorine:

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

gas.

2Al(s) + 3Cl2(g) ==> 2AlCl3(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 sublimationtakes 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, Al2Cl

6, shown above.

Reaction of element with water:

o None due to protective oxide layer.

Reactions of the hexa-aqua aluminium ion:

o 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 NH3(aq)) and acids.
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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.)

o With aqueous sodium carbonate solution, the hydroxide ppt. is formed, and, because of its acidic nature, bubbles ofcarbon

dioxide gas are evolved.

2[Al(H2O)6]3+

(aq) + CO32-

(aq) 2[Al(H2O)5(OH)]2+

(aq) + H2O(l) + CO2(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.

The extraction of aluminium
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Aluminium is obtained from mining the mineral bauxite.

The purified bauxite ore ofaluminium 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 connectionscalled the cathode (-, negative),

attracting positive ions e.g. Al3+

, and the anode (+, positive) which attracts negative ions e.g. O2-

.

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 O2-

ions]

Carbon(graphite) for the electrodes.


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

Electrolysismeans 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.

The redox details of the electrode processes:

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.

o

Al

3+

+ 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.

o 2O2-

==> O2 + 4e-

o or2O2-

- 4e-==> O2

Note: Reduction and Oxidation always go together!

The overall electrolytic decomposition is ...

o aluminium oxide => aluminium + oxygen

o 2Al2O3==> 4Al + 3O2

o 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.


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Ionsare 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?

o 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. O2-

attracted to the positive anode electrode,

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

o 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

o 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.

The properties and uses of aluminium


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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 ofsteel. Many metals can be given a coating of a different metal to protect them or to improve their

appearance.

o Aluminiumis 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.

o 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.

o 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 aboutstructure of metal alloys on the metallic bonding page.

Reactions of aluminium

o Reaction with chlorine
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o

o Ifdry 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) + 3Cl2(g)==> 2AlCl3(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) + 3Cl2(g)==> 2FeCl3(s)

Aluminium chloride reacts exothermically as it is hydrolysed by waterto 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.

AlCl3(s) + 3H2O(l)==> Al(OH)3(s) + 3HCl(g)

However, if a large excess of wateris rapidly added, a weakly acidic solution of aluminium chloride is formed,

with the minimum of nasty fumes!

AlCl3(s) + aq ==> Al3+

(aq) + 3Cl-(aq)


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or more correctly: AlCl3(s) + 6H2O(l)==> [Al(H2O)6]3+

(aq) + 3Cl-(aq)

The solution is slightly acidic, because the hexa-aqa 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)

o The surface ofaluminium 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) + 3O2(g)==> 2Al2O3(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, theThermit Reactionshows its rightful place in the reactivity series of metals.

o 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) + Fe2O3(s)==> Al2O3(s) + 2Fe(s)
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This is a typical displacement reaction by a more reactive metal displacing a less reactive metal from one of its

compounds.

o 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)==> 2AlCl3(aq) + 3H2(g)

o The reaction with dilute sulphuric acid is very slow to form colourless aluminium sulphate and hydrogen.

aluminium + sulphuric acid ==> aluminium sulphate + hydrogen

2Al(s) + 3H2SO4(aq)==> Al2(SO4)3(aq) + 3H2(g)

o 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)

Amphoteric nature of aluminium hydroxide and acidity of the hexaaquaaluminium ion

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

o [Al(H2O)6]3+

(aq) + 3OH-(aq) ==> [Al(H2O)3(OH)3](s) + 3H2O(aq)


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o 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.

o 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).

o [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.

o 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) ==> AlCl3(aq) + 3H2O(l)

Al(OH)3(s) + 3HNO3(aq) ==> Al(NO3)3(aq) + 3H2O(l)

2Al(OH)3(s) + 3H2SO4(aq) ==> Al2(SO4)3(aq) + 6H2O(l)

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

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

2[Al(H2O)6]3+

(aq) + 3CO32-

(aq) ==> 2[Al(H2O)3(OH)3](s) + 3CO2(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) + CO32-

(aq) ==> 2[Al(H2O)4(OH)2]+

(aq) + CO2(g) + 3H2O(aq)


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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.

o 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.

o [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.

o 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.

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

o [Al(H2O)6]3+

=> [Al(OH)(H2O)5]2+

=> [Al(OH)2(H2O)4]+

=> [Al(OH)3(H2O)3](s) precipitate

o dissolving => [Al(OH)4(H2O)3]-=> [Al(OH)5(H2O)]

2-=> [Al(OH)6]

3-


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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-).

1234From 1 to 7 happen as you add more alkali, increasing pH and the OH

-concentration, removing protons from the aluminium

complex.

567 *From 7 back to1 represents what happens when you add acid, decreasing pH, increasing H

+/H3O

+concentration and

protonating the aluminium complex.

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

o 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:

o 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.

o RC N + 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.

o C6H5NO2 + 6[H] ==> C6H5NH2 + 2H2O

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

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

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

a level notes on gp3 boron and aluminium

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