How Cu reacts with NH3

An equilibrium involving copper(II) ions Students add ammonia to a solution of copper(II) sulfate, observe the colour changes taking place, and then reverse the reaction by the addition of sulfuric acid. Read our standard health & safety guidance Lesson organisation This experiment is best carried out by students working individually. It takes about 15 - 20 minutes. Apparatus and chemicals Each student or working group will require: Eye protection Test-tubes, 3 Test-tube rack Test-tube holder Dropping pipettes, 2 Copper(II) sulfate solution, 1.0 mol dm-3 (Harmful at this concentration), about 3 cm3 (see note 1) Ammonia solution, 1.0 mol dm-3 (Low hazard at this concentration), about 10 cm3 Dilute sulfuric acid, 1.0 mol dm-3 (Irritant at this concentration), about 10 cm3 Technical notes Copper(II) sulfate solution (Harmful at concentration used) Refer to CLEAPSS Hazcard 27C Ammonia solution (Low hazard at concentration used) Refer to CLEAPSS Hazcard 6 and Recipe Card 4 Dilute sulfuric acid (Irritant at concentration used) Refer to CLEAPSS Hazcard 98A and Recipe Card 69 1 The copper(II) sulfate solution is most conveniently supplied in a bottle fitted with teat pipette. Procedure a Put 10 drops of copper(II) sulfate solution into each of two test-tubes. b Add ammonia solution drop-by-drop to the first test-tube. Shake the tube gently from side to side after adding each drop. What happens as you add a few drops of the solution? c Add more drops of ammonia solution. What happens? Continue until you have a clear blue solution. d Divide the solution from c into two test-tubes. Add dilute sulfuric acid drop-by-drop to one of the solutions from c. Shake the tube gently from side to side after adding each drop. Do you get back to where you started – compare the three test-tubes? e Can you repeat the whole process by adding ammonia again to the acidified solution? Teaching notes If this experiment is being carried out with pre-A-level students, the reactions occurring can simply be explained by reference to the addition of an alkali (containing hydroxide ions) being added to a solution of a copper compound, producing copper(II) hydroxide initially and later a complex compound of ammonia. The reversal of the process is easy to explain since sulfuric acid is capable of neutralising the alkaline ammonia and causing the reaction to reverse back to the start: CuSO4(aq)(pale blue solution) + 2NH3(aq) + 2H2O(l) → Cu(OH)2(s) + (NH4)2SO4(aq)(pale blue precipitate) Cu(OH)2(s)(pale blue precipitate) + ammonia → complex copper compound (dark blue solution) A rather more advanced treatment in terms of complexes and ligand exchange would involve the following explanation: 1 Ammonia is a weak base and forms a few ammonium and hydroxide ions in solution NH3(g) + H2O(l) ⇌ NH4+(aq) + OH-(aq) 2 The hexaaquacopper(II) ions react with hydroxide ions to form a precipitate. This involves deprotonation of two of the water ligand molecules: [Cu(H2O)6]2+(aq)(pale blue) + 2OH-(aq) → [Cu(H2O)4(OH)2](s)(pale blue precipitate) + 2H2O(l) 3 The copper(II) hydroxide precipitate reacts with ammonia molecules to form tetraamminediaquacopper(II) ions This involves ligand exchange: [Cu(H2O)4(OH)2](s)(pale blue precipitate) + 4NH3(aq) ⇌ [Cu(NH3)4(H2O)2)]2+(aq)(dark blue solution) + 2OH-(aq) + 2H2O(l) 4 Thus the overall reaction, combining 2 with 3, gives: [Cu(H2O)6]2+(aq) + 4NH3(aq) ⇌ [Cu(NH3)4(H2O)2)]2+(aq) + 4H2O(l) 5 Addition of dilute sulfuric acid introduces H+ ions, which react with NH3 molecules to form NH4+ ions, and this draws the equilibrium in 4 back to the left-hand side, regenerating the hexaaquacopper(II) ions in the process. Health & Safety checked, August 2008 Web Links This link describes how to obtain a solid sample of a salt containing the dark blue tetraamminediaquacopper(II) ions: http://firstyear.chem.usyd.edu.au/LabManual/E16.pdf (Websites accessed August 2008)