2P32 – Principles of Inorganic Chemistry
Dr. M. Pilkington
Lecture 13 – Redox Reactions Between Metal Ions,the Chelate Effect, Hard and Soft Acids and Bases
1.
1.
Outer-sphere reactions “through space”.
2.
Inner-sphere reactions “through bonds”.
3.
The chelate effect
4.
Hard and soft acids and bases
Oxidation-Reduction Reactions of Transition Metal Complexes There are two main mechanisms: i) Outer Sphere: electron transfer is “through space” – the coordination sphere of the metal stays intact. ii) Inner Sphere: electron transfer takes place through a bridging ligand (through bonds) – the coordination sphere of the metals are altered in some way.
i) Outer Sphere Reactions Example: [Ru(H2O)6]2+ + [Ru(H2O18)6]3+ d6 low spin
[Ru(H2O)6]3+ + [Ru(H2O18)6]2+
d5 low spin O18 labeled H2O
The two coordinated ions are the same in everyway accept their oxidation state “self – exchange” – the metal is the same, just different oxidation state
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Outer Sphere Reactions [Ru(H2O)6]2+ + [Ru(H2O18)6]3+
[Ru(H2O)6]3+ + [Ru(H2O18)6]2+
Characteristics Measured: 1.
1st- order reaction with respect to both reactants
2.
No scrambling of H2O18 with H2O (i.e. no complexes with [Ru(H2O)5(H2O18)]
3.
Rate of H2O exchange with solvent water is very slow compared to the rate of electron transfer.
Mechanism of the Outer-Sphere Reaction – 3 steps 1.
[Ru(H2O18)6]3+ + [Ru(H2O)6]2+
{ [Ru(H2O18)6]3+ + Ru(H2O)6]2+} ion pair
Collision of the two reactant coordinated ions to form an outersphere complex (transient ion pair). electron transfer
2.
[Ru(H2O18)6]3+
+ [Ru(H2O)6]2+
{ [Ru(H2O18)6]2+ + Ru(H2O)6]3+}
Electron transfer takes place instantaneously within this ion pair 3.
The ion pair dissociates to give the product.
N.B. ONLY the electrons are transferred, there is no exchange of H2O ligands.
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How fast are these reactions? – The reactions can be very fast if the reactants have similar bond distances. For example:
1.
[Co(NH3)6]3+ / low spin d6 Co-N = 1.94A (194pm)
2.
[Ru(H2O)6]3+ / [Ru(H2O)6]2+ low spin d6 high spin d7 Ru-O = 2.03A Co-N = 2.12A
44 M-1s-1 (fast)
3.
[FeIII(CN)6]3-/ [FeII(CN)6]4low spin d5 low spin d6 Fe-CN = 1.92A Fe-CN = 1.95A
700 M-1s-1 (very fast)
[Co(NH3)6]2+ high spin d7 Co-N = 2.11A (211pm)
Rate Constant 1x10-6 M-1s-1 (very slow)
If the M-L distances are very different in the two complexes, there will be a large energy barrier, that is the energy of activation involved in bringing about the electron transfer.
Ea is a measure of how much stretching and compressing of bonds must take place during the reaction.
Frank-Condon principle: electron exchange is most rapid if no movement of atoms in reactants and products is necessary.
So the more stretching and compressing that must be done the slower the reaction.
So we can conclude then that when geometry and bond distances in reactants and products are similar, electron exchange is fast.
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ii) Inner Sphere Reactions – electron transfer takes place through a bridging ligand.
For example: H+ [CoIII(NH3)5Cl]2+ + [CrII(H2O)6]2+ inert: low spin d6 labile: high spin d4 does not exchange gives H2O readily water quickly
Reaction Mechanism 1.
[CrII(H2O)6]2+
fast
[CoII(H2O)6]2+ + [CrIII(H2O)5Cl]2+ + 5NH4+ labile: d7 high spin inert: d3 exchange H2O at fast rate
[CrII(H2O)5]2+ + H2O
labile d4 high spin dissociation of a water from the labile complex. 2.
[CoIII(NH3)5 Cl ]2+ + [CrII(H2O)5]2+ the lone pairs can
[(NH3)5CoIII- Cl-CrII(H2O)5]4+ bridged molecule
bond to the Cr(II) ion when H2O is removed Inner sphere electron transfer reactions involve the formation of a bridged complex where the two metals are connected by a bridging ligand that helps promote electron transfer.
e- move 3.
[(NH3)5CoIII- Cl-CrII(H2O)5]4+
[(NH3)5CoII- Cl-CrIII(H2O)5]4+ since the bond is labile then this bond breaks between the Co-Cl H+, H2O
4.
[CoII(H2O)6]2+ + 5NH4+ + [CrIII(H2O)5Cl]2+ If we carry out the reaction in a solution containing 36Cl-, then no 36Cl- appears in the Cr product. This proves that the Cl is coming from the Co complex.
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Bridging ligands – must have at least two lone pairs of electrons (on the same atom or different atoms).
For example
Cl- OH- C N
N
CS
can bridge through N and S M-N
C-S-M
or can bridge through S only M-S-C
N
M'
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3. The Chelate Effect Can be defined as the unusual stability of a coordination compound involving a chelating multidentate ligand, as compared with equivalent compounds involving monodentate ligands.
The added stability of a complex that results from the formation of chelate rings. For Example Ni2+ + NH3
[Ni(NH3)]2+
[Ni(NH3)]2+ + NH3
[Ni(NH3)2]2+
Overall Reaction [Ni(NH3)2] 2+ Ni2+ + 2NH3 β2 = 1.0 x 105
N.B it is not possible to measure these steps seperately Ni2+ + N
N
2+
Ni
(en)
N
two nitrogens attached to the metal
N
K (formation constant) = 3 x 107
5-membered chelate ring
If we now compare this to: Ni2+ + 2NH3
[Ni(NH3)2]2+ β2 = 1.0 x 105
Ni(en) is 300 x more stable than [Ni(NH3)2]2+ due to the formation of a 5-membered chelate ring.
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Why does a chelate ring result in greater stability? Two Explanations: 1. Effective Concentration Ni
2+
Ni
N
N N
N
The second Nitrogen is held close to the metal by the ethyl (CH2CH2) group, allowing the second bond to form easily. whereas for the monodentate case: Ni-NH3
The second N must keep moving until it comes within close proximity of a metal to bond, i.e it is less likely to form the second bond with the Ni2+ atom.
+ NH3
5-membered ring are generally more stable than 6-membered rings. 4 membered rings are much rarer as they suffer from bond strain.
The order of stability is as follows: L
L
<<
M L
Too much strain within the bonds to make 4membered
L
>
M L
>>
M
7-membered ring
L
Most stable since the bond strain is less and the ligand allows for an increase in bonding ability
Increase in bond strain and in this case the second ligand is a little further away so it is not as likely to bond than the 5-membered ring case
When rings become much larger than 6 the enhancement of the local concentration is diminished. Ligands are positioned further away from the metal in the larger rings.
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2.
Entropy Considerations
a)
[Ni(H2O)6] 2+ + 2NH3
[Ni(NH3)2(H2O)4]2+ + 2H2O
no change in the number of particles i.e. ΔS = O b)
[Ni(H2O)6] 2+ +
N
[
N
N Ni
(H2O)4]2+ + 2H2O
N
two particles produce three particles ΔSo = +ve which increases disorder so the forward reaction is favored. N.B when ΔSo = +ve then ΔGo becomes more negative.
Example 2: Consider the following reactions: [Ni(H2O)6]2+ + 6NH3(aq) 1 6 [Ni(H2O)6]2+ + 3en(aq) 1 3
[Ni(NH3)6]2+ + 6H2O 1 6 [Ni(en)3]2+ + 6H2O 1 6
β = 4.0 x 108 β = 2.0 x 1018
The larger the number of particles distributed randomly, the higher the entropy of the reaction.
ΔGo = ΔHo –TΔSo
A large increase in ΔSo results in a more negative ΔGo.
Since ΔGo = -RTlnK then a more negative ΔGo leads to a more positive equilibrium (formation) constant.
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4. “Hard and Soft” Acids and Basis
Metal ion – lewis acid (electron pair acceptor).
Ligand – lewis base (electron pair donor).
M Lewis Acid
:L Lewis Base
Long known that some metal ions form their most stable complexes with F- and Odonors or N- donors rather than Cl-, S or P donors. For Example: Mg2+ and Al 3+ have a high affinity for F- and O- donors.
Other metal ions form their most stable complexes with heavier halides (Cl-, Br-, I-), S-donors or P-donors. For Example: Hg2+, Au3+ form stable complexes such as AuBr4, Hg(S-cys)2 The Mercaptans (mercury capturers) HgS
Hg2+ + S2K = 10-54 (would take 1000 L of H2O to dissociate 1 atom)
very stable HgS does not dissociate readily, very stable.
Hard acids and Bases are not easily polarizable
Hard Acids: low electronegativty and/or highly positively charged small ions. i.e. H+, alkali and alkaline earth metals transition metals with higher charges (Fe3+, Cr3+ , Cr6+ etc…)
Hard Bases: high electronegativity, small or negative charge on donor atom (ligand), small donors. i,e. F-, O2-, OH-, ROH-, NH3, RNH2, PO43- (the charge is spread over all 4 oxygen atoms that is why it is a hard base).
A large component of ionic bonding in complexes between hard acids and bases, i,e they form very stable complexes.
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Soft acids and Bases are easily polarizable Soft Acids: high electronegativity (for metal), large size and/or low positive
charge. e.g.
Hg 2+ Cd2+ Cu2+ Au2+ Pt2+
e-negativity plus
M0
Pt4+
2.00 1.69 1.90 2.54 2.28 2.28
(metal ions in zero oxidation states)
i.e. Ni(CO)4 (Ni0) Fe 0(CO)4 Soft Bases: low electronegativity (for a non metal), large size, p-
bond/delocalized bonding. i,e. H-, CN-, CO, Ph3P, SCN-, I-, RSSoft to soft interactions involve significant covalent character in the bonding.
“Borderline” characteristics between hard and soft
Acids:
Fe2+
Cu2+ there are others
Fe3+ (hard)
Cu+ (soft)
Bases: N2, NO2, NCS-
Pyridine is softer than NH3 because of the π-bonding H N N
softer
H
N
soft
Ambidentate Bases – Can have both hard and soft donor atoms: e.g. NO2-, NCS-
Soft donor O
S O
O
Hard donor
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You should review this for your midterm/final exam.
HSAB Theory - Summary: Advantages
HSAB theory: idea that soft acids preferentially bind with soft bases and hard acids with hard bases.
Rule used for rationalizing and predicting the relative stabilities of transition metal complexes and other compounds.
Handy way to organize ideas “useful rule of thumb”
Disadvantages
Some acids and bases do not follow the rules e.g. [Ag(NH3)2]+ (soft acid/hard base).
Other factors which contribute to the strength of the bonds between donor and acceptor are
1.
Size of cation and donor atom.
2.
Their charges and electronegativities and orbital overlap between them.
3.
Competing substitution reactions in aqueous solution e.g. ligands displace water.
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