1. Unless otherwise stated, all images in this file have been reproduced from: Blackman, Bottle, Schmid, Mocerino and Wille, Chemistry, 2007 (John Wiley) ISBN: 9 78047081 0866

2. Slide 2/16 e CHEM1002 [Part 2] A/Prof Adam Bridgeman (Series 1) Dr Feike Dijkstra (Series 2) Weeks 8 – 13 Office Hours: Monday 2-3, Friday 1-2 Room: 543a e-mail: adam.bridgeman@sydney.edu.au e-mail: feike.dijkstra@sydney.edu.au

3. Slide 3/16 e Complexes II For octahedral complexes with formulae [MX 2 Y 4 ], cis and trans geometrical isomers are possible For square planar complexes with formulae [MX 2 Y 2 ], cis and trans geometrical isomers are possible For octahedral complexes with bidentate ligands, optical isomerism is also possible Metal complex formation can greatly increase solubility Summary of Last Lecture

4. Slide 4/16 e Lecture 14 Transition Metals Electron Configuration Oxidation States Colours Magnetism Blackman Chapter 13, Sections 13.4 and 13.7 Lecture 15 Metals in Biological Processes Essential Elements Toxic Elements Medicinal Uses Blackman Chapter 13 Complexes III

5. Slide 5/16 e Transition (or d-block) Metals

6. Subshells (from CHEM1001 lecture 14) Each shell is divided into subshells called s, p, d, f…. There is one extra subshell for each new shell –First shell: 1s –Second shell: 2s and 2p –Third shell: 3s, 3p and 3d –Fourth shell: 4s, 4p, 4d and 4f 1 st shell1s1s 2 nd shell2s2s2p2p 3 rd shell3s3s3p3p3d3d 4 th shell4s4s4p4p4d4d4f4f 5 th shell5s5s5p5p5d5d5f5fetc 6

7. Filling Subshells (from CHEM1001 lecture 14) 2 electrons can fit into a s subshell 6 electrons can fit into a p subshell 10 electrons can fit into a d subshell 14 electrons can fit into a f subshell 1s1s 2s2s2p2p 3s3s3p3p3d3d 4s4s4p4p4d4d4f4f 5s5s5p5p5d5d5f5fetc energy increases 7

8. Slide 8/16 e Electronic Configurations of Atoms Each sub shell is made from “orbitals” which can each accommodate up to 2 electrons: 2 electrons can fit into a s subshell in one orbital 6 electrons can fit into a p subshell in three orbitals 10 electrons can fit into a d subshell in five orbitals s 3d3d p

9. Slide 9/16 e Electronic Configurations of Cations Main group metals like Na, K, Mg and Ca lose all of their valence electrons  Na: [Ne] 3s 1 Na + : [Ne] 3s 0 or just [Ne]  Mg: [Ne] 3s 2 Mg 2+ : [Ne] 3s 0 or just [Ne]  K: [Ar] 4s 1 K + : [Ar] 4s 0 or just [Ar]  Ca: [Ar] 4s 2 Ca 2+ : [Ar] 4s 0 or just [Ar] Transition metal cations may have valence electrons left over and can form more than one oxidation number

10. Slide 10/16 e Electronic Configurations of Cations Left over valence electrons fill 3d only Group number gives number of valence electrons Cation has (group number – oxidation number) electrons  Mn 7+ : group 7 and oxidation number 7 so has: (7 – 7) = 0 electrons: [Ar](3d) 0  Mn 2+ : group 7 and oxidation number 2 so has (7 – 2) = 5 electrons: [Ar](3d) 5  Ni 2+ : group 10 and oxidation number 2 so has (10 – 2) = 8 electrons: [Ar](3d) 8  Ni 3+ : group 10 and oxidation number 3 so has (10 – 3) = 7 electrons: [Ar](3d) 7

11. Slide 11/16 e Electronic Configurations of Cations To minimize repulsion between electrons, they occupy d- orbitals singly with until they have to pair up: 3d3d  Mn 2+ : [Ar](3d) 5 3d3d  Ni 2+ : [Ar](3d) 8 If the metal cation has unpaired electrons, the complex will be attracted to a magnet: paramagnetic

12. Slide 12/16 e Aqueous Oxoanions of Transition Metals One of the most characteristic chemical properties of these elements is the occurrence of multiple oxidation numbers, often associated with different colours. Ion Ox. No.Colour VO 3 - +5yellow VO 2+ +4green V 3+ +3blue V 2+ +2violet

13. Slide 13/16 e Colourful Complexes Aqueous solutions of the Co(III) complexes (from left to right): [Co(NH 3 ) 5 OH 2 ] 3+, [Co(NH 3 ) 6 ] 3+, trans-[Co(en) 2 Cl 2 ] +, [Co(en) 2 O 2 CO] + and [Co(NH 3 ) 5 Cl] 2+. All contain Co(III): colour influenced by the ligand

14. Slide 14/16 i Absorbed and Observed Colours Unless the d-orbitals are empty, half full or full, electrons can be excited from one d-orbital to another: absorption of light which we see as colour

15. Slide 15/16 e Summary: Complexes II Learning Outcomes - you should now be able to: Complete the worksheet Work out the electron configurations of atoms and cations Work out the number of unpaired electrons Answer review problems 13.59-13.62 in Blackman Next lecture: The Biological Periodic Table

16. Slide 16/16 x Practice Examples 1.How many d-electrons and how many unpaired electrons are there in the following complexes? A.K 2 [NiCl 4 ] B.[Co(en) 3 ]Cl 3 C.[CrCl 2 (OH 2 ) 4 ] + D.K 2 [Zn(OH) 4 ] E.[PtCl 2 (NH 3 ) 2 ] 2. Consider the compound with formula [CoCl 2 (NH 3 ) 4 ]Br  2H 2 O (i) Write the formula of the complex ion. (ii) Write the symbols of the ligand donor atoms. (iii) What is the d electron configuration of the metal ion in this complex?