Reactions of Inorganic Compounds in Aqueous Solution Metal-aqua ions Metal aqua ions are formed in aqueous solution. Each complex has six water ligands arranged in a octahedral shape. The complex has the same charge as the metal ion as water is a neutral ligand 2 O 2 O O 2 2+ Cu O 2 O 2 O 2 [Cu blue [Co [Fe pink green [Fe [Cr ]3+ [Al Ruby Violet colourless (appears (appears green) Yellow) In solution Cr(III) often appears green and Fe(III) appears yellow/brown due to hydrolysis reactions. The ruby and violet colour is only really seen in solid hydrated salts that contain these complexes. Cr (ruby) Fe (violet) [Zn colourless [Mn Very pale pink [Ni green N Goalby chemrevise.org 1
Acidity or hydrolysis reactions The following equilibria happen in aqueous solutions of metal ions. [M + 2 O [M O) 5 (O)] + + 3 [M + 2 O [M O) 5 (O) + 3 For example, [Cu + 2 O [Cu O) 5 (O)] + + 3 The equilibria lead to generation of acidic solutions with M 3+ ions, and very weakly acidic solutions with M 2+ ions. The 3 + ions are noticeably more acidic. 3+ 2+ O O 2O O 2O Al + + Al 2O 2O The acidity of [M is greater than that of [M in terms of the greater polarising power (charge/size ratio) of the 3+ metal ion. The greater the polarising power, the more strongly it attracts the water molecule. This weakens the O- bond so it breaks more easily. Lewis acids and bases Definitions: Lewis acid: electron pair acceptor Lewis base: electron pair donator In the formation of complex ions the ligand is the Lewis base because it is donating a pair of electrons in the dative covalent bond and the metal ion is the Lewis acid. Metal aqua ions can also do acidity type reactions with the bases (sodium hydroxide) O -, (ammonia) N 3 and (sodium carbonate) CO 3. Reaction with limited O - and limited N 3 Aqueous complex ions react with limited amounts of sodium hydroxide and ammonia to form coloured precipitates. The colours of the precipitates formed can be used to identify the metal ion Examples Cu blue ppt Co blue Fe Ni Green ppt Mn pale Zn White ppt N Goalby chemrevise.org 2
Al (O) 3 (s) white ppt Cr(O) 3 (s). Fe (O) 3 (s) The bases O - and ammonia when in limited amounts form the same hydroxide precipitates. They form in deprotonation acid base reactions [Cu (aq) + 2O - (aq) Cu + 2 2 [Co (aq) + 2N 3 (aq) Co + 2N 4 + (aq) [Ni (aq) + 2O - (aq) Ni + 2 2 [Zn (aq) + 2N 3 (aq) Zn + 2N 4 + (aq) [Cr (aq) + 3O - (aq) Cr (O) 3 (s) + 3 2 [Fe (aq) + 3N 3 (aq) Fe (O) 3 (s) + 3N 4 + (aq) ere the N 3 and O - ions are acting as Bronsted-Lowry bases accepting a proton This process can happen step wise removing one proton at a time. One proton is removed at each step until the hydroxide precipitate is formed. e.g. [Cr (aq) + O - (aq) [Cr O) 5 (O) (aq) + 2 [Cr O) 5 (O) (aq) + O - (aq) [Cr ( ] + (aq) + 2 [Cr ( ] + (aq) + O - (aq) Cr (O) 3 (s) + 2 In some cases if sodium hydroxide is continued to be added until in excess then further deprotonation takes place, and the hydroxide precipitate dissolves. Reaction with excess O - With excess NaO, the Cr, Zn and Al hydroxides dissolve. Cr becomes [Cr(O) 6 ] 3- (aq) ution. Al becomes [Al(O ] - (aq) colourless solution. Zn becomes [Zn(O ] (aq) colourless solution Cr (O) 3 (s) + 3O - (aq ) [Cr(O) 6 ] 3- (aq) + 3 2 Al (O) 3 (s) + O - (aq ) [Al(O ] - (aq) + 3 2 Zn + 2O - (aq ) [Zn(O ] (aq) + 4 2 The hydroxides are classed as amphoteric because they react and dissolve in both acids and bases. + O - [Al(O) 3 ](s) [Al(O ] [Al( - 2 acidic neutral alkaline Cr (O) 3 (s) + 3 + (aq ) [Cr (aq) Al (O) 3 (s) + 3 + (aq ) [Al (aq) N Goalby chemrevise.org 3
Reaction with excess N 3 With excess N 3 ligand substitution reactions occur with some of the hydroxides and their precipitates dissolve Cu, Co, Zn, Ni and Cr hydroxides all dissolve in excess ammonia Cr becomes [Cr(N 3 ) 6 purple solution Zn becomes [Zn(N 3 colourless solution Ni becomes [Ni(N 3 ) 6 blue solution Co becomes [Co(N 3 ) 6 pale yellow solution Cu becomes [Cu(N 3 deep blue solution [Cu(N 3 [Co(N 3 ) 6 [Ni(N 3 ) 6 The ligands N 3 and 2 O are similar in size and are uncharged. Ligand exchange occurs without a change of co-ordination number for Co, Ni and Cr. Cr(O) 3 (s) + 6N 3 (aq) [Cr(N 3 ) 6 (aq) + 3 2 + 3O - (aq) Co( (s) + 6N 3 (aq) [Co(N 3 ) 6 (aq) + 4 2 + 2O - (aq) This substitution may, however, be incomplete as in the case with Cu. Cu( (s) + 4N 3 (aq) [Cu(N 3 (aq) + 2 2 + 2O - (aq) In these reactions N 3 is acting as a Lewis base donating an electron pair. Reactions with Chloride Ions Addition of a high concentration of chloride ions (from conc Cl or saturated NaCl) to an aqueous ion leads to a ligand substitution reaction. The Cl - ligand is larger than the uncharged 2 O and N 3 ligands so therefore ligand exchange can involve a change of co-ordination number. Addition of conc Cl to aqueous ions of Cu, Co and Ni lead to a change in coordination number from 6 to 4. [CuCl 4 ] yellow/ution [CoCl 4 ] blue solution These are tetrahedral in shape [NiCl 4 ] [Cu + 4Cl - [CuCl 4 ] + 6 2 O [Co + 4Cl - [CoCl 4 ] + 6 2 O N Goalby chemrevise.org 4
Reactions with Carbonate solution The 2+ ions react differently to the 3+ ions with carbonate solutions. If + ions are added in sufficient concentrations to carbonate ions the following equilibria are pushed to towards products and CO 2 is produced + + CO 3 CO 3 - + + CO 3-2 CO 2 The metal aqua 2+ ions are acidic but not sufficiently acidic to bring about the changes mentioned above as the following equilibrium does not produce enough + ions [M + 2 O [M O) 5 (O)] + + 3 So when 2+ metal aqua ions are added to carbonate ions they form coloured carbonate precipitates The 2+ ions with carbonate solution results in MCO 3 ppt being formed Cu = blue/, Co = pink ppt, Fe(II)=, Zn = white ppt, Mn= pink ppt, Ni= Mn 2+ (aq) + CO 3 (aq) MnCO 3 (s) Cu 2+ (aq) + CO 3 (aq) CuCO 3 (s) Fe 2+ (aq) + CO 3 (aq) FeCO 3 (s) Co 2+ (aq) + CO 3 (aq) CoCO 3 (s) [Cu + CO 3 CuCO 3 + 6 2 O [Fe + CO 3 FeCO 3 + 6 2 O [Co + CO 3 CoCO 3 + 6 2 O These are precipitation reactions CoCO 3 CuCO 3 The 3+ ions with carbonate solution form a M(O) 3 ppt and CO 2 gas is evolved. MCO 3 is formed with 2+ ions but M 2 (CO 3 ) 3 is not formed with 3+ ions. The difference is explained by the greater polarising power of the 3+ ion due to its higher charge density. The greater the polarising power, the more strongly it attracts the water molecule. This weakens the O- bond so it breaks more easily, producing a high concentration of + ions. The metal aqua ion reacts releasing + ions and forms the hydroxide precipitate M(O) 3 This acidity decomposes the CO 3 ions CO 3 + 2 + 2 CO 2 Al forms white ppt of Al(O) 3 + CO 2 Cr (III) forms of Cr(O) 3 + CO 2 Fe(III) forms of Fe(O) 3 + CO 2 2[Cr (aq) + 3CO 3 (aq) 2Cr(O) 3 (s) +3CO 2 + 3 2 2[Fe (aq) + 3CO 3 (aq) 2Fe(O) 3 (s) +3CO 2 + 3 2 2[Al (aq) + 3CO 3 (aq) 2Al(O) 3 (s) +3CO 2 + 3 2 These are classed as acidity reactions. N Goalby chemrevise.org 5
2+ Ion Summary [MCl 4 ] [CuCl 4 ] [CoCl 4 ] blue sol [Zn(O ] colourless sol water Conc Cl - excess O - [M [Cu [Co [Fe [Ni [Mn [Zn Few drops O - or N 3 M( excess N 3 [Cu(N 3 + blue sol pink sol very pale pink sol colourless sol Cu( blue ppt Co( blue ppt Fe( Ni( Mn( light Zn( white ppt deep blue solution [Co(N 3 ) 6 yellow/light brown solution [Ni(N 3 ) 6 blue sol [Zn(N 3 colourless sol Air or 2 O 2 CO 3 Air or 2 O 2 [Co(N 3 ) 6 brown sol MCO 3 CuCO 3 blue/ FeCO 3 CoCO 3 pink ppt ZnCO 3 = white ppt MnCO 3 = pink ppt NiCO 3 = M(O) 3 Fe(O) 3 Co(O) 3 Mn(O) 3 dark Cr 2 O 7 orange sol 3+ Ion Summary + O - [Cr(O) 6 ] 3- [Al(O ] - colourless sol 2 O 2 CrO 4 yellow sol excess O - Few drops [M O - or N 3 excess N M(O) 3 3 [Cr(N3 ) 6 violet sol [Cr [Fe [Al yellow sol colourless sol + Cr(O) 3 Fe(O) 3 Al(O) 3 white ppt CO 3 (+ CO 2 ) N Goalby chemrevise.org 6