- their structure and bonding
- their acid-base properties.
We shall consider the oxides of Group 1, Group 2, B and Al from group 3, C, Si and Pb from group 4, P, S, Cl, and the first d block.
Ionic Oxides and hydroxides are ionic if they have high melting and boiling temperatures and a crystal structure which is composed of separate ions.
Covalent Oxides and oxo-acids are molecular if they have low melting and boiling temperatures and they are made up of discrete molecules. A giant covalent compound has a high melting and boiling temperature with a continuous structure.
Layer lattice. Many hydroxides have a layer lattice structure with metal ions sandwiched between layers of hydroxide ions. These are ionic/covalent in character. On heating to a high temperature, they give ionic oxides.
Basic. Metals with electronegativity less than 1.3 form basic oxides and hydroxides. These oxides and hydroxides dissolve easily in dilute acids. If they are soluble in water they give rise to alkaline solutions.
Acidic. Non-metals with electronegativity greater than 2.0 form acidic oxides. If these are soluble in water they give oxo-acids which are acidic. The giant covalent oxides B2O3 and SiO2 do not dissolve in water but are classified as acidic as they react with fused NaOH; in addition their oxo-acids, prepared by hydrolysis of the chlorides, dissolve in aqueous NaOH.
Amphoteric. Metals with electronegativity between 1.3 and 2.0 should form amphoteric oxides but many of them have very insoluble hydroxides which dissolve only slightly in aqueous NaOH. These oxides will be considered to be amphoteric only if their hydroxides dissolve readily in aqueous NaOH; otherwise they will be considered basic.
Formula. If the hydroxy compound is basic or amphoteric it will be written in the form M(OH)x, even though many tri-positive metals form hydroxides with the formula MO.OH.xH2O.. If the hydroxy compound is acidic, it will be written in the form HxXOy.
Group 1 and Group 2.
All the Group 1 oxides are ionic and alkaline:
M2O (s) + H2O(l) à 2MOH(aq)
The same is true for Group 2 oxides:
MO(s) + H2O(l) à M(OH)2(aq)
although beryllium and magnesium hydroxides are nearly insoluble (solubility of hydroxides increases down the group).
Beryllium oxide dissolves in both dilute acid and alkali and so is amphoteric (compare with aluminium):
BeO(s) + 2H+(aq) à Be2+(aq) + H2O(l)
BeO(s) + 2OH- + H2O(l) à Be(OH)42-(aq)
Boron oxide is giant covalent and is acidic. It is insoluble in water but dissolves in fused NaOH. The weak acid, H3BO3, made by hydrolysis of BCl3, dissolves in aqueous NaOH (compare silicon).
B2O3(s) + NaOH(l) à 2NaBO2(l) + H2O(g)
H3BO3(s) + OH-(aq) à B(OH)4-(aq)
Aluminium oxide is ionic and the hydroxide forms a layer lattice. Its oxide shows similar behaviour to that of boron (substitute Al for B in the equations above), and as it also dissolves in acids, it is amphoteric.
Al2O3(s) + 6H+( aq) à 2Al3+(aq) + 3H2O(l)
Al(OH)3(s) + 3H+(aq) à Al3+(aq) + 3H2O(l)
The resulting aluminium ion, Al(H2O)63+(aq), has a high charge density and acts as a weak acid with pKa 4.8, the same as ethanoic acid.
[Al(H2O)6]3+(aq) + H2O(l) à [Al(H2O)5OH]2+ + H3O+(aq)
On the addition of aqueous NaOH, the equilibrium moves to the right until Al(OH)3(s) is precipitated; this easily reacts with excess NaOH to give the aluminate ion, Al(OH)4-(aq).
Carbon monoxide is molecular and is considered to be neutral. Formally it is the anhydride of methanoic acid HCOOH, but it does not react with water although the reverse process, the dehydration of methanoic acid with concentrated sulphuric acid, is a standard method for preparing carbon monoxide.
Carbon dioxide is molecular and dissolves in water to give carbonic acid, which is a weak acid with pK1 6.4:
CO2(g) + H2O(l) à H2CO3(aq)
Only about 1 molecule in 600 of CO2(g) is converted into the acid and most of it remains in the form CO2(aq).
Silicon dioxide is giant covalent , containing single Si-O bonds. A molecular structure would contain Si=O bonds; these would be much weaker than two of the Si-O single bonds found in the silica lattice. SiO2 is insoluble in water but reacts with fused NaOH:
SiO2(s) + 2NaOH(l) à Na2SiO3(l) + H2O(g)
The very weak acid, H4SiO4(s), prepared by the hydrolysis of the chloride, also dissolves in aqueous NaOH.
H4SiO4(s) + 2OH-(aq) à SiO32-(aq) + 3H2O(l).
Lead monoxide is an ionic solid whose colour varies from yellow to orange. It is red when hot and turns back to yellow on cooling. It disolves in dilute nitric acid (the sulphate and chloride are insoluble in cold water) and in concentrated aqeuous sodium hydroxide..
PbO(s) + 2H+(aq) à Pb2+(aq) + H2O(l)
PbO(s) + H2O(l) + OH-(aq) à Pb(OH)3-(aq)
Addition of aqueous NaOH to a solution of a lead salt gives a white precipitate usually written as Pb(OH)2(s) (but possibly 2PbO.H2O)which reacts with excess NaOH to give Pb(OH)3-(aq), plumbate(II).
Lead dioxide, PbO2(s) is an ionic solid which is brown. It decomposes to PbO and oxygen on heating.. At low temperature it dissolves in concentrated HCl to give PbCl4 and the complex ion PbCl62-. On warming, these substances give off chlorine gas. PbO2(s) also dissolves in concentrated NaOH.
PbO2(s) + 2OH-(aq) + 2H2O(l) à Pb(OH)62-(aq).
Lead(IV) hydroxide is unknown; PbO2(s) is precipitated instead.
Phosphorus, sulphur amd chlorine.
All form molecular oxides which are acidic.
Phosphorus(III) oxide and phosphorus(V) oxide have molecular formulae P4O6(s) and P4O10(s) but are often seen written as P2O3(s) and P2O5(s). They react with water to give the weak acid H3PO3 with pK1 2.0, and the slightly stronger acid, H3PO4 with pK1 1.49. This illustrates the general trend that the oxide of the higher oxidation state is the more acidic.
P2O3(s) + 2H2O(l) à 2H3PO3(aq)
P2O5(s) + 2H2O(l) à 2H3PO4(aq)
Sulphur dioxide and sulphur trioxide dissolve in water to give the weak acid, H2SO3 with pK1 1.81), and the very strong acid, H2SO4. Note that even in this case the second ionisation is small enough for it to have very little effect on the pH; see the pH of sulphuric acid page.
Chlorine forms several oxides of which Cl2O and Cl2O7 are the most important. They react with water to form the weak acid HClO with pKa 7.53), and the very strong acid, HClO4.
Cl2O(g) + 2H2O(l) à 2HClO(aq)
Cl2O7(l) + 2H2O(l) à 2HClO4(aq)
3d block metals
Scandium oxide is very similar to aluminium oxide
The most important oxide of titanium is TiO2 which is acidic.
Vanadium has four principal oxidation states. VO and V2O3 are both ionic and basic and dissolve in acids to give the lilac [V(H2O)6]2+ ion and the blue-green [V(H2O)6]3+ ion. Both of these ions are readily oxidised to V(IV) or V(V). VO2 is ionic and amphoteric. It dissolves in acids to give the blue VO2+ ion (probably [VO(H2O)5]2+) and in concentrated alkali to give the brown VO44- ion. V2O5 has a borderline ionic/covalent structure. It reversibly loses oxygen on heating, and this may be why it is such an efficient catalyst in the Contact Process for the conversion of sulphur dioxide to sulphur trioxide. It is amphoteric.
V2O5(s) + 2H+(aq) à 2VO2+(aq) + H2O(l) yellow solution
V2O5(s) + 6OH-(aq) à 2VO43-(aq) + 3H2O(l) colourless solution
The two main oxides of chromium are the green Cr2O3 which is ionic, and the red CrO3 which is covalent. The former is amphoteric as the green-grey precipitate of Cr(OH)3 dissolves in acids to give the purple [Cr(OH)6]3+ ion, and in alkali to give a green solution which may possibly contain the Cr(OH)4- ion. With ammonia solution, the green-grey precipitate dissolves to give a pink-purple [Cr(NH3)6]3+ complex ion.
[Cr(H2O)6]3+(aq) + 3OH-(aq) à Cr(OH)3(s) + 6H2O(l)
This is a deprotonation or acid-base reaction.
Cr(OH)3(s) + 6NH3(aq) à [Cr(NH3)6]3+(aq) + 3OH-(aq)
This is a ligand exchange reaction.
Chromium trioxide is acidic and dissolves in water to give the fairly strong acid, H2CrO4 with pKa 0.74. On heating it loses oxygen to give Cr2O3. When treated with alkali the solution contains the yellow CrO42-(aq) ion. This reacts with strong acids to give an orange solution containing the dichromate (VI) ion:
|2CrO4 2-(aq)||+2H+(aq)||à||Cr2O72-(aq)||+ H2O(l)|
Manganese(II) compounds contain the very pale pink d5 ion which is resistant to oxidation. If alkali is added to a solution containing the Mn2+(aq) ion, the off white precipitate of Mn(OH)2 is readily oxidised by air to brown hydrated Mn2O3 and then possibly to 2MnO2.H2O. Reduction of the purple MnO4- ion in alkaline solution (for example by alkenes) yields brown hydrated MnO2.
Iron(II) oxide FeO is difficult to prepare as it rapidly oxidises in air to give Fe2O3, which contains the d5 ion. Both of these oxides are basic. Freshly precipitated Fe(OH)2 is green but it rapidly darkens on exposure to air and finally forms brown Fe(OH)3.
Pink cobalt compounds contain the [Co(H2O)6]2+ ion. On the addition of NaOH a precipitate of Co(OH)2 is formed which starts off blue but turns pink on warming. The blue form may have four, rather than six, OH- groups round each cobalt ion in the solid state. With ammonia Co(OH)2 dissolves giving a complex which undergoes air oxidation to give the brown [Co(NH3)6]3+(aq) ion.
Solutions of nickel(II) containing the green [Ni(H2O)6]2+(aq) ion give, with NaOH, a pale green precipitate of Ni(OH)2 which dissolves in ammonia to give the purple [Ni(NH3)6]2+(aq) ion.
When copper(I) oxide Cu2O(s) is treated with aqueous sulphuric acid, the Cu+(aq) ion initially formed disproportionates to a brown precipitate of Cu(s) and the blue Cu2+(aq) ion, The hydrated [Cu(H2O)6]2+(aq) ion has four water molecules more firmly bound than the other two, and it is these water molecules which are substituted when pale blue Cu(OH)2(s) dissolves in aqueous ammonia to give the deep blue [Cu(NH3)4(H2O)2]2+(aq) ion.
Zinc oxide goes yellow on heating and returns to white again on cooling. This is result of movement of ions within the crystal lattice which is reversible. It is amphoteric, giving the [Zn(H2O)4]2+(aq) ion with acids and the Zn(OH)3-(aq) and Zn(OH)42-(aq) ions when treated with NaOH. It also reacts with ammonia (contrast Al(OH)3(s), which does not) to give the [Zn(NH3)4]2+(aq) ion.
Peter Hughes was in the Chemistry Department at Westminster School, where he was formerly Head of Science.
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