1. Chemistry of Coordination Compounds Brown, LeMay Ch 24 AP Chemistry Monta Vista High School To properly view this presentation on the web, use the navigation arrows below and left-click on each page to view any animations.

2. 2 24.1: Structure of Complexes Complex: species in which a central metal ion (usually a transition metal) is bonded to a group of surrounding molecules or ions Coordination compound: compound that contains a complex ion or ions.

3. 3 A coordination compound, or complex, consists of: Metal ion Acts as a Lewis acid (e - pair acceptor) Electrophile: species that is “e - poor” and seeks e - (gets attacked by nucleophile) Ligand or complexing agent: molecule or ion with a lone pair of e - that bonds to a metal ion Acts as a Lewis base (e- pair donor) Coordinate covalent bond: metal-ligand bond Nucleophile: species that is “e - rich” and seeks an e - poor area of a molecule (seeks an electrophile)

4. 4 Lewis Structures of common ligands NH 3 CN - S 2 O 3 2- SCN - H 2 O (not always included in formula, however)

5. 5 Complexation reactions Ligand usually added “in excess” on AP Usually result in color changes (colors generally originate from e - transitions in a partially filled d shell) Change properties of metal ion Thermodynamic (  H,  S,  G) Electrochemical (Eº)

6. 6 The golden-orange compound is CoCl3*6NH3 while the purple compound only has 5 ammonia molecules in the coordinated compound. As shown in the ball-and-stick model, the chlorides serve as counter ions to the cobalt/ammonia coordiation complex in the orange compound, while one of the ammonia molecules is replaced by Cl in the purple compound. In both cases, the coordination geometry is octahedral around Co.

7. 7 Notation Write complexes in square brackets, with charge on outside Ex: Cu 2+ (aq) + 4 NH 3 (aq) → [Cu(NH 3 ) 4 ] 2+ (aq) H | Cu 2+ (aq) + 4 :N ─ H (aq) → Cu | H :NH 3 H 3 N: :NH 3 2+

8. 8 Coordination number Number of positions where a ligand can bond. Similar to oxidation state Each metal ion has a characteristic (i.e., typical) coordination number, which can be predicted according to crystal field theory. Ag + : coordination number = 2 (2 ligand bonding positions); results in a linear complex [Ag(NH 3 ) 2 ] + (aq) H H H | | | Ag+ (aq) + 2 :N ─ H (aq) → H─ N:Ag:N─H(aq) | | | H H H +

9. 9 Zn 2+ & Cu 2+ : coordination number = 4; tetrahedral complex Ex: [Zn(H 2 O) 4 ] 2+ (aq) Pt 2+ : coordination number = 4; square planar complex (d 8 e- structure) Ex: [Pt(CN) 4 ] 2- (aq)

10. 10 Al 3+, Cr 3+, and Fe 3+ : coordination number = 6; octahedral complex Ex: [Cr(NH 3 ) 5 Cl] 2+ (aq)

11. 11 Is dependent on: Charge of ligand: Ni 2+ : 6 NH 3 or 4 CN - (since CN - transfers more negative charge) Size of ligand: Fe 3+ : 6 F - or 4 Cl - (larger ions take up more space)

12. 12 24.2: Chelates & Polydentate ligands http://chemlabs.uoregon.edu/GeneralResources/models/bidentate.html Ligands with more than one bonding position Ethylenediamine (“en”, C 2 H 4 N 2 ), or oxalate, C 2 O 4 2- Ex: Cr 3+ (aq) + 3 C 2 O 4 2- (aq) → [Cr(C 2 O 4 ) 3 ] 3-

13. 13 24.3: Nomenclature 1.Name cation before anion; one or both may be a complex. (Follow standard nomenclature for non- complexes.) 2.Within each complex (neutral or ion), name all ligands before the metal. Name ligands in alphabetical order If more than one of the same ligand is present, use a numerical prefix: di, tri, tetra, penta, hexa, … Ignore numerical prefixes when alphabetizing.

14. 14 Neutral ligands: use the name of the molecule (with some exceptions) NH 3 ammine-H 2 Oaqua- Anionic ligand: use suffix –o Br - bromo-CN - cyano- Cl - chloro- OH - hydroxo- 3.If the complex is an anion, use –ate suffix Record the oxidation number of the metal in parentheses (if appropriate). Ex: [Co(NH 3 ) 5 Cl]Cl 2 pentamminechlorocobalt (III) chloride

15. 15 Nomenclature practice 1. K 4 [Fe(CN) 6 ] 2. [Cr(NH 3 ) 4 (H 2 O)CN]Cl 2 3. Na[Al(OH) 4 ] potassium hexacyanoferrate (II) tetrammineaquacyanochromium (III) chloride sodium tetrahydroxoaluminate

16. 16 * 24.5: Color & Magnetism Atoms or ions with a partially filled d-shell usually exhibit color because the e - transitions fall within the visible part of the EM spectrum. Ex: transition metals such as Cu 2+ (blue) and Fe 3+ (orange) Therefore, those with empty or filled d-shells are usually colorless. Ex: alkali & alkaline earth halides, Al 3+

17. 17 * 24.6: Crystal Field Theory Created to explain why transition metal ions in complexes (having unfilled d-shells) are not necessarily paramagnetic. With coordination bonding, valence d-orbitals are not truly degenerate. Instead, they “split”. Some are lower in energy (more stable) and some higher. http://scienceworld.wolfram.com/chemistry/CrystalFieldTheory.html

18. 18 The gap between the higher and lower energy levels is called the crystal-field splitting energy, which varies with each ligand, yielding different E, (different, different colors). e- in an “unfilled” d-shell can actually be all paired (i.e., diamagnetic). Ex: Co 3+ (has 6 d e-)