Lecture 11 Reaction Types and Mechanisms for Inorganic Complexes

Transcription

1 2P32 Principles of Inorganic Chemistry Dr. M. Pilkington Lecture 11 Reaction Types and Mechanisms for Inorganic Complexes Variations in reactivity Reaction types substitution, dissociation, addition and redox reactions, reactions of coordinated ligands. Equilibrium constants for substitution reactions. Kinetic vs. thermodynamic stability of metal complexes Measuring rates of water exchange in aquometal ions. Another Consequence of Crystal field Stabilization Energy 1. Variations in Reactivity: Let's take as a single example the difference in ligand substitution rates for the complexes [Co(NH 3 ) 6 ] 3+ and [Ni(NH 3 ) 6 ] 2+. Each complex has a negative enthalpy for the following reaction, so both complexes are thermodynamically capable of undergoing ligand exchange with water. [M(NH 3 ) 6 ] n+ + 6 H > [M(H 2 ) 6 ] n+ + 6 NH 4 + However, the exchange reaction for the nickel complex is very fast, while the reaction for the cobalt complex takes days or weeks to go to completion. Why is this? 1

2 The short answer is found if we consider the reaction mechanism, i.e., how the reaction must take place. There are two limiting possibilities, both of them involving a 2-step reaction as shown below. Mechanism 1. ne of the NH 3 ligands dissociates from the complex to give a 5- coordinate intermediate. This intermediate then forms a complex with a water molecule. [M(NH 3 ) 6 ] n+ ---> [M(NH 3 ) 5 ] n+ + NH 3 [M(NH 3 ) 5 ] n+ + H 2 ---> [M(NH 3 ) 5 H 2 ] n+ Repetition of this process ultimately gives the hexaaquo ion. Mechanism 2. A water molecule coordinates to the complex to give a 7-coordinate intermediate. This intermediate then loses an ammonia molecule. [M(NH 3 ) 6 ] n+ + H 2 ---> [M(NH 3 ) 6 H 2 ] n+ [M(NH 3 3) H ] n+ ---> [M(NH 3 H ] n ) NH 3 The process repeats itself until the hexaaquo ion is formed. Each of these mechanisms involves a reaction to give a reactive intermediate with a different coordination number (5 or 7). When M = Co(III) (low-spin d 6 ), considerable crystal field stabilization is lost because the CFSE is so much greater for octahedral coordination than the CFSE than any possible geometry in 5- or 7-coordination. The loss in CFSE on going from 6-coordination to 5- or 7-coordination increases the activation energy for the reaction, thus slowing it. 2

3 For most other ions, there is less loss (or even a gain) in CFSE on going from 6- to 5- or 7-coordinate intermediates. Thus exchange reactions for these ions are not retarded by the loss in CFSE. 2. Reaction Types for Inorganic Complexes 1. Substitution Reactions - one (or more) ligands replace another ligand in the coordination sphere of a metal e.g. [Pt(NH ) 4 ] Cl [Pt(NH 3 ) 3 Cl] + NH 3 These are the most common reactions of coordination compounds. They involve substitution of one ligand in a coordination sphere for another. 2. Dissociation Reactions reactions which decrease in the coordination number of the metal. e.g. [Co(H 2 0) 6 ]Cl 2 6H 2 + CoCl 2 6-coordinate 2-coordinate 3

4 3. Addition Reactions increases the coordination number of the metal. e.g. Cu(acac) 2 + py Cu(acac) 2 py 4-coordinate 5-coordinate A square planar bis(acetylacetonato)copper(ii) molecule accepts a pyridine (py) ligand to form a square pyramidal product. Acac N Py 4. Redox Reactions oxidation-reduction or electron transfer reactions. [Ru(NH 3 ) 6 ] 3+ + [Cr(H 2 ) 6 ] 2+ [Ru(NH 3 ) 6 ] 2+ + [Cr(H 2 ) 6 ] 3+ The hexammineruthenium(iii) ion is reduced by the reaction with the Cr(II) ion, 5. Reactions of Coordinated Ligands [Cr(H 2 ) 6 ] 3+ + H - [Cr(H 2 ) 5 (H)] 2+ + H 2 Reactions of a ligand that take place without breaking the M-L bond. e.g. water ligand in the hexaaquochromium(iii) reacting with a hydroxide ion to produce the corresponding hydroxo complex. r replacement of the central hydrogen of acac with a Br atom. Cr + 3Br 2 Cr Br + 3HBr Acac 3 3 4

5 3. Equilibrium Constants for Metal Complex Formation Stepwise constants one ligand at a time is replaced: For Example for the formation of : [Ni(NH 3 ) 6 ] 2+ [Ni(H 2 ) 6 ] 2+ K 1 NH 3 [Ni(H 2 ) 5 NH 3 ] 2+ + H 2 [Ni(H 2 ) 5 NH 3 ] 2+ K 2 NH 3 [Ni(H 2 ) 4 (NH 3 ) 2 ] 2+ + H 2 and so on for K 3, k 4, k 5 K 6 [Ni(H 2 )(NH 3 ) 5 ] 2+ + NH 3 [Ni(NH 3 ) 6 ] 2+ + H 2 K 1 = 1 st stepwise formation constant (equilibrium constant for the formation of the complex). K 2+ 1 = [Ni(H 2 ) 5 NH 3 ] [Ni(H 2 ) 6 ] 2+ [NH 3 ] K 2 = [Ni(H 2 ) 4 (NH 3 ) 2 ] 2+ [Ni(H 2 ) 5 (NH 3 )] 2+ [NH 3 ] 5

6 If we only want the reaction with addition of 2NH 3 [Ni(H 2 ) 6 ] NH 3 [Ni(H 2 ) 4 (NH 3 ) 2 ] H 2 verall Reaction verall Constant = [Ni(H 2 ) 4 (NH 3 ) 2 ]2+ the formation constant = beta [Ni(H 2 ) 6 ] 2+ [NH 3 ] 2 We can also arrive at the previous result by multiplying the first and second equilibrium (formation) constants: K 1 x K 2 = [Ni(H 2 ) 5 NH 3 ] 2+ [Ni(H 2 ) 4 (NH 3 ) 2 ] 2+ [Ni(H 2) 6] 2+ [NH 3 ] [Ni(H 2 ) 5 (NH 3 )] 2+ [NH 3 ] [Ni(H 2 ) 4 (NH 3 ) 2 ] 2+ [Ni(H 2 ) 6 ] 2+ [NH 3 ] 2 = 6

7 So can we make this compound.? [Ni(NH 3 ) 6 ] 2+ - to decide we have to look at the formation constant. 6 = [Ni(NH 3 ) 6 ]2+ [Ni(H 2 ) 6 ] 2+ [NH 3 ] 6 = K 1 x K 2 x K 3 x K 4 x K 5 x K 6 For [Ni(NH 3 ) 6 ] 2+ K 1 = 6.3 x 10 2 K 2 = 1.6 x 10 2 K 3 = 5 x 10 1 K 4 = 16 K 5 = 6.3 K 6 = = 5.4 x 10 8 i.e the formation constant is very large which means the complex will be formed It is generally true that: K 1 >K 2 >K 3 >K 4 >K 5 >K 6 This is partly statistical the fewer H 2 s to replace, the less likely the reaction will occur. For [Ni(H 2 ) 6 ] H 2 s and anyone of them can be replaced. For [Ni(H 2 )(NH 3 ) 5 ] 2+ -only one H 2 can be replaced. 4. Kinetics versus Thermodynamics Stabilty (Stable or Unstable) 1. Thermodynamic stability: G = H T S where, H = enthalpy or heat change; T is the absolute temperature, and S is the entropy change/disorder. G = - RTlnK where K is the formation constant For thermodynamic stability the requirements are K > 1, G < this means that the products are more stable than the reactants. 7

8 2. Kinetic Stability refers to the rate of reaction. Metal complexes that undergo substitution reactions very slowly are said to be INERT. Metal complexes that undergo substitution reactions very quickly are said to be LABILE. Examples: 1. [Co(NH 3 ) 6 ] H 3 + [Co(H 2 ) 6 ] NH 4 + K- the equilibrium (formation) constant for the reaction is extremely large (10 30 ) so we would say that the cation is unstable toward reaction with acid, but it takes weeks or months to see any evidence of reaction. Accordingly [Co(NH 3 ) 6 ] 3+ must be classified as being Unstable thermodynamically but Inert kinetically. 2. [Co(NH 3 ) 6 ] H 3 + [Co(H 2 ) 6 ] NH 4 + The equilibrium constant K for the reaction is very large and the above reaction is instantaneous. Hence [Co(NH 3 )] 2+ is kinetically Labile and Thermodynamically Unstable. 3. [Ni(CN) 4 ] 2- is exceptionally stable (thermodynamically). The equilibrium constant for its formation is in the vicinity of Ni CN [Ni(CN) 4 ] 2- K = At the same time the complex anion is labile, that is the cyanide ligands in the coordination sphere exchange rapidly with those found free in aqueous solution. [Ni(CN) 4 ] CN - [Ni(CN) 3 ( 14 CN)] 2- +CN - [Ni(CN) 4 ] 2- is Stable but Labile. 8

9 Some coordination compounds are kinetically inert, whereas others turn out to be labile. This lability seems to be unrelated to the thermodynamic stability of the compound. Complexes of the first row transition metal ions with the exception of Cr 3+ and Co 3+ are generally labile, whereas most second and third row transition metal ions are inert. In the case of H 2 exchange you can experimentally label with 17 or 18 and you can then see which ions are kinetically labile Labile half life < 1 min Inert half life > 1 min 5. Rates of water exchange in aquometal ions. e.g. CaCl 2 Ca 2+ 2Cl - H Ca H H ion-dipole interactions. Lewis acid-base if Ca 2+ is assumed to coordinate to the H 2 H Ca 2+ H H We can measure the rates of exchange of water: [M(H 2 ) n ] x H 2 [M(H 2 ) n-1 18 (H 2 )] x+ +H 2 Water exchange rates (in text) page 103, Table 5.5 9

10 Rate constants for water exchange for various ions [M(H Very slow 2 ) n ] x H 2 K > 10-3 to 10-6 sec -1 [M(H 2 ) n-1 18 (H 2 )] x+ + H 2 Very Fast K > 10 8 sec -1 Group 1A as we go down the group the cations are getting larger and the charge density decreases so the M n+ -H 2 bond is getting weaker and more easily broken Group 2A the charge density is larger (doubly charged) so the strength of the bond is greater so the rate of exchange is slower Class 1 Groups IA, IIA and IIB as well as Cu 2+ and Cr 2+ ; k > Diffusion controlled reactions Ionic charge and size are important. Rates of Exchange: Li + < Na + < K + < Cs + Group 1A Be 2+ < Mg 2+ < Ca 2+ < Sr 2+ < Ba 2+ Group 2A In these two series, the smallest ions exchange the slowest: Within a group e.g. Group 1A, size is important: Li + H 2 Cs + H 2 Stonger because the Li + Less strong attracting Ion is smaller 10

11 Li+ Cs d d Coloumb s Law E = Q+,Q-/r if we measure the distance between the centres of the atoms as used in Coloumb's Law then we have much greater distance for Cs + -H 2. This means the Li-H 2 bond is stronger than the Cs-H 2 since the Lithium ion holds onto the electrons of the 2- and does not want to break to release H 2 as easily as Cs-H 2 Moving from Group 1A to Group 2A charge is important Increasing charge on the ion H 2 exchange rates become slower. This is because as the charge goes up the bond strength increases so we have a stronger bond that is more reluctant to break to release the water. Hence the higher +2 charged ions of Group 2A have slower rates than the +1 charged Group 1A ions. Cu 2+ and Cr 3+ Cu 2+( d 9 ) and Cr 2+ (d 4 ) are structurally distorted by the Jahn-Teller effect, with bond to the axial ligands longer and weaker than bonds to the equatorial groups. Therefore the ground state structures are not far removed from the transition state structures. Jaan-Teller effect is most often encountered in octahedral complexes most commonly 6-coordinate Cu(II). The electronic configuration of the ion gives 3 electrons in the two degenerate e g orbitals, The complex distorts along the z-axis which lowers the overall energy. This distortion normally takes the form of elongating the bonds to the ligands lying along the z-axis. This lowers the replusion between theelectronpairontheligandandanyelectronswith orbitals in the z-component, thus lowering the energy of the complex. 11

12 Classes 2 and 3 - Includes most of the first row TM ions and the lanthanides plus Be 2+, Al 3+, V 2+. Rate constants 1 to 10 8 sec -1. These elements who tend to undergo water exchange and ligand exchange reactions. For T.M. metal ions the correlation of rate with size is not obeyed, e.g. Cr 2+, Ni 2+, and Cu 2+ have identical radii. Mn 2+ > Fe 2+ > Co 2+ substitution rates decrease across the series. This is due to the increase in Z eff across the series and increase in E(M--L). ff d- electron configurations are important because the CFSE will affect the rates of exchange here. Class 4 - Rate constants are in the range sec -1. This includes Cr 3+, Co 3+, Rh 3+, Ir 3+, Pt 2+. For these metal ions the rate of exchange is partially related to the size of the cations and partly to the CFSE. The basic assumption is that the significant contribution to the activation energy in a substitution reaction is the change in d-orbital energy on going from the ground state of the complex to the transition state. d 3 and d 6 ions are predicted to be inert e.g. Rh 3+, Cr 3+ and Co 2+ 12