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 Metal aqua ions are formed in aqueous solution. [M(H2O)6 ] 2+, limited to M = Fe (green) and Cu (blue); [M(H2O)6 ] 3+, limited to M = Al (colourless), and Fe (violet) In solution and Fe(III) appears yellow/brown due to hydrolysis reactions. The violet colour is only really seen in solid hydrated salts that contain these complexes. Acidity or hydrolysis reactions [M(H2O)6 ] 3+ + H2O [M(H2O)5 (OH)]2+ + H3O+ The following equilibria happen in aqueous solutions of metal ions. The equilibria lead to generation of acidic solutions with M3+ ions, and very weakly acidic solutions with M2+ ions. The 3+ ions are noticeably more acidic. The acidity of [M(H2O)6 ] 3+ is greater than that of [M(H2O)6 ] 2+ 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-H bond so it breaks more easily.
3.2.6 Reactions of ions in aqueous solution (A-level only)
In aqueous solution, the following metal-aqua ions are formed: [M(H2O)6] 2+, limited to M = Fe and Cu [M(H2 O)6 ] 3+, limited to M = Al and Fe
The acidity of [M(H2O)6] 3+ is greater than that of [M(H2O)6] 2+
Students should be able to:
• explain, in terms of the charge/size ratio of the metal ion, why the acidity of [M(H2O)6 ] 3+ is greater than that of [M(H2O)6 ] 2+
Reaction with limited OH- and limited NH3 The bases OH- and ammonia when in limited amounts form the same hydroxide precipitates. They form in deprotonation acid base reactions M(OH)2 (H2O)4 (s) : Cu blue ppt, Fe (II) green ppt M(OH)3 (H2O)3 (s) : Fe(III) brown ppt, Al white ppt [Cu(H2O)6 ] 2+ (aq) + 2OH- (aq) Cu(H2O)4 (OH)2 (s) + 2H2O (l) [Al(H2O)6 ] 3+ (aq) + 3OH- (aq) Al(H2O)3 (OH)3 (s) + 3H2O (l) [Fe(H2O)6 ] 2+ (aq) + 2NH3 (aq) Fe(H2O)4 (OH)2 (s) + 2NH4 + (aq) [Fe(H2O)6 ] 3+ (aq) + 3NH3 (aq) Fe(H2O)3 (OH)3 (s) + 3NH4 + (aq) This process can happen step wise removing one proton at a time. Be able to write equations for this too. e.g. [Al(H2O)6 ] 3+ (aq) + OH- (aq) [Al(H2O)5 (OH)]2+ (aq) + H2O (l) Reaction with excess OHWith excess NaOH the Al hydroxide dissolves. Al becomes [Al(OH)4 ] – (aq) colourless solution. This hydroxides is classed as amphoteric because it reacts and dissolves in both acids and bases. Al(H2O)3 (OH)3 (s) + OH- (aq ) [Al(OH)4 ] – (aq) + 3H2O (l) Al(H2O)3 (OH)3 (s) + 3H+ (aq ) [Al(H2O)6 ] 3+ (aq) Reaction with excess NH3 With excess NH3 a ligand substitution reaction occurs with Cu and its precipitate dissolve This substitution is incomplete with Cu. Cu becomes [Cu(NH3 ) 4 (H2O)2 ] 2+ deep blue solution Cu(OH)2 (H2O)4(s) + 4NH3 (aq) [Cu(NH3 ) 4 (H2O)2 ] 2+ (aq) + 2H2O (l) + 2OH- (aq) [Al(H2O)5 (OH)]2+ (aq) + OH- (aq) [Al(H2O)4 (OH)2 ] + (aq) + H2O (l) Here the NH3 and OHions are acting as Bronsted-Lowry bases accepting a proton In this reactions NH3 is acting as a Lewis base donating an electron pair.
Reactions with Carbonate solution The 2+ ions react differently to the 3+ ions with carbonate solutions. The 2+ ions with carbonate solution results in MCO3 ppt being formed (Cu blue/green, Fe(II) green ) Cu2+ (aq) + CO3 2- (aq) CuCO3 (s) These are precipitation reactions The 3+ ions with carbonate solution form a M(OH)3 ppt and CO2 gas is evolved. MCO3 is formed with 2+ ions but M2 (CO3 )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. Al forms white ppt of Al(OH)3 (H2O)3 + CO2 Fe(III) forms brown ppt of Fe(OH)3 (H2O)3 + CO2 These are classed as acidity reactions. Fe2+ (aq) + CO3 2- (aq) FeCO3 (s) 2[Fe(H2O)6 ] 3+ (aq) + 3CO3 2- (aq) 2Fe(OH)3 (H2O)3(s) +3CO2 + 3H2O(l) 2[Al (H2O)6 ] 3+ (aq) + 3CO3 2- (aq) 2Al(OH)3 (H2O)3(s) +3CO2 + 3H2O(l) [Cu(H2O)6 ] 2+ + CO3 2- CuCO3 + 6H2O [Fe(H2O)6 ] 2+ + CO3 2- FeCO3 + 6H2O
3.2.6 Reactions of ions in aqueous solution (A-level only)
Some metal hydroxides show amphoteric character by dissolving in both acids and bases (eg hydroxides of Al3+).
Students should be able to:
• describe and explain the simple test-tube reactions of: M2+(aq) ions, limited to M = Fe and Cu, and of M3+(aq) ions, limited to M = Al and Fe, with the bases OH– , NH3 and CO3 2–
A-level Chemistry exemplar for required practical No. 11 Carry out simple test-tube reactions to identify transition metal ions in aqueous solution: An investigation of some transition metal compounds Most transition metal compounds are coloured. Some of them are used as dyes and pigments. A dye is a soluble coloured compound. A pigment is an insoluble coloured compound. Both dyes and pigments have to be resistant to chemical change. Four solutions, labelled P, Q, R and S, have been provided by a supplier as possible dyes. You will carry out tests on these solutions. You will record what you observe for each test. You should ensure that you record observations on dropwise addition, on addition to excess and on standing. Where no visible change is observed, write ‘no visible change’. In this task, you are not required to identify any of the solutions or any of the reaction products. Requirements You are provided with the following: four solutions – labelled ‘Solution P’, ‘Solution Q’, ‘Solution R’ and ‘Solution S’ sodium hydroxide solution sodium carbonate solution silver nitrate solution 12 test tubes 7 dropping pipettes test-tube rack 250 cm3 beaker access to hot water plentiful supply of distilled or deionised water. Suggested method Test 1(a) a) Place about 10 drops of solution P in a test tube. b) Add sodium hydroxide solution, dropwise with gentle shaking, until in excess. c) Do not discard this mixture. d) Repeat this test with solution Q and then solution R and then solution S.Test 1(b) a) Half fill a 250 cm3 beaker with the freshly boiled water provided. b) Allow the four test tubes containing the mixtures from Test 1(a) to stand in the beaker of hot water for about 10 minutes. c) While you are waiting, begin Test 2. Test 2 a) Place about 10 drops of sodium carbonate solution in a test tube. b) Add about 10 drops of solution P and shake the mixture gently. c) Repeat this procedure with solution Q and then with solution R and then solution S. Test 3 a) Place about 10 drops of solution P in a test tube. b) Add about 10 drops of silver nitrate solution and shake the mixture gently. c) Repeat this procedure with solution Q and then with solution R and then solution S. d) Allow the four test tubes to stand for about 10 minutes. Sample results Test 1(a) and (b) P Q R S Initial blue solution yellow solution light blue solution pale green solution Add NaOH white precipitate orange/brown precipitate deep blue precipitate grey/green precipitate On standing in hot water green solution no visible change no visible change no visible change Test 2 P Q R S Addition of sodium carbonate grey precipitate orange/brown precipitate and effervescence blue/green precipitate grey/green precipitate Test 3 P Q R S Addition of silver nitrate no visible change white precipitate white precipitate light brown precipitate
3.2.6 Reactions of ions in aqueous solution (A-level only)
Required practical 11
Carry out simple test-tube reactions to identify transition metal ions in aqueous solution.
