Download Lecture 11 – Reaction Types and Mechanisms for Inorganic

2P32 – Principles of Inorganic Chemistry
Dr. M. Pilkington
Lecture 11 – Reaction Types and Mechanisms for Inorganic
Complexes


Variations in reactivity



Equilibrium constants for substitution reactions.
Reaction types – substitution, dissociation, addition and redox
reactions, reactions of coordinated ligands.
Kinetic vs. thermodynamic stability of metal complexes
M s in rates
Measuring
t s of
f water
t exchange
x h n in aquometal
m t l iions.
ns
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(NH3)6]3+ and [Ni(NH3)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(NH3)6]n+ + 6 H3O+ ---> [M(H2O)6]n+ + 6 NH4+

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?
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
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. One of the NH3 ligands dissociates from the complex to give a 5coordinate
coord
nate intermediate.
ntermed ate. This
h s intermediate
ntermed ate then forms a complex w
with
th a water
molecule.
[M(NH3)6]n+ ---> [M(NH3)5]n+ + NH3
[M(NH3)5]n+ + H2O ---> [M(NH3)5H2O]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(NH3)6]n+ + H2O ---> [M(NH3)6H2O]n+
[[M(NH
(
]n+ ---> [[M(NH
(
]n+ + NH3
3)6H2O]
3)5H2O]



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 d6), 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.
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
For most other ions, there is less loss (or even a gain) in CFSE on going from 6to 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(NH3)4]2+ + Cl-
[Pt(NH3)3Cl]+ + NH3


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(H20)6]Cl2
6-coordinate
6H2O + CoCl2
2-coordinate
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3.
Addition Reactions –increases the coordination number of the
metal. e.g.
Cu(acac)2 + py
Cu(acac)2py
4-coordinate
5-coordinate
A square planar bis(acetylacetonato)copper(II) molecule accepts
a pyridine (py) ligand to form a square pyramidal product.
O
O
N
Acac
4.
Py
Redox Reactions – oxidation-reduction or electron transfer
reactions.
[Ru(NH3)6]3+ + [Cr(H2O)6]2+
[Ru(NH3)6]2+ + [Cr(H2O)6]3+
The hexammineruthenium(III) ion is reduced by the reaction with
the Cr(II) ion,
5.
Reactions of Coordinated Ligands
[Cr(H2O)5(OH)]2+ + H2O
[Cr(H2O)6]3+ + OHReactions 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.
Or replacement of the central hydrogen of acac with a Br atom.
O
O
+
Cr
O
3
3Br2
Cr
+ 3HBr
Br
O
3
Acac
4
3.

Equilibrium Constants for Metal Complex Formation
Stepwise constants – one ligand at a time is replaced:
For Example for the formation of : [Ni(NH3)6]2+
[Ni(H2O)6]2+
[Ni(H2O)5NH3]2+
K1
NH3
[Ni(H2O)5NH3]2+ + H2O
K2
NH3
[Ni(H2O)4(NH3)2]2+ + H2O
and so on for K3, k4, k5
[Ni(H2O)(NH3)5]2+ + NH3
K6
[Ni(NH3)6]2+ + H2O
K1 = 1st stepwise formation constant (equilibrium constant for
the formation of the complex).
K1 =
[Ni(H2O)5NH3]22+
[Ni(H2O)6]2+[NH3]
K2 =
[Ni(H2O)4(NH3)2]2+
[Ni(H2O)5(NH3)]2+[NH3]
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
If we only want the reaction with addition of 2NH3
[Ni(H2O)6]2+ + 2NH3
[Ni(H2O)4(NH3)2]2+ + 2H2O
Overall Reaction
Overall Constant
2+
= [Ni(H2O)4(NH3)2]
[Ni(H2O)6]2+[NH3]2
the formation
constant = beta

We can also arrive at the previous result by multiplying the first and
second equilibrium (formation) constants:
K1 x K2 =
[Ni(H2O)5NH3]2+
[Ni(H2O)4(NH3)2]2+
[Ni(H2O)6]2+[NH3]
[[Ni(H
( 2O))5((NH3)]2+[[NH3]
[Ni(H2O)4(NH3)2]2+
[Ni(H2O)6]2+[NH3]2
=

6

So can we make this compound ….?
[Ni(NH3)6]2+ - to decide we have to look at the formation constant.
6 =
[Ni(NH3)6]2+
[Ni(H2O)6]2+[NH3]6
= K1 x K2 x K3 x K4 x K5 x K6
For [Ni(NH3)6]2+
K1 = 6.3 x 102
K2 = 1.6 x 102
K3 = 5 x 101
K4 = 16
K5 = 6.3
K6 = 1.1
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6 = 5.4 x 108
i.e the formation constant is very large which means the complex
will be formed




It is generally true that:
K1>K2>K3>K4>K5>K6
This is partly statistical – the fewer H2O’s to replace, the less likely the
reaction will occur.
For [Ni(H2O)6]2+ - 6H2O’s and anyone of them can be replaced.
For [Ni(H2O)(NH3)5]2+ - only one H2O 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 < O this means that
the products are more stable than the reactants.
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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(NH3)6]3+ + 6H3O+
[Co(H2O)6]3+ + 6NH4+
K- the equilibrium (formation) constant for the reaction is extremely
large (1030) so we would say that the cation is unstable toward
reaction
ti with
ith acid,
id b
butt it ttakes
k weeks
k or months
th tto see any evidence
id
of reaction.
Accordingly [Co(NH3)6]3+ must be classified as being Unstable
thermodynamically but Inert kinetically.
2. [Co(NH3)6]2+ + 6H3O+
[Co(H2O)6]2+ + 6NH4+
The equilibrium constant K for the reaction is very large and the above
reaction is instantaneous.
Hence [Co(NH3)]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 
Ni2+ + 4CN
[Ni(CN)4]2-
K = 1030
At the same time the complex anion is labile, that is the cyanide ligands
in the coordination sphere
p
exchange
g rapidly
p y with those found
f
free
f
in
aqueous solution.
[Ni(CN)4]2- +
14CN-
[Ni(CN)3(14CN)]2- +CN-
[Ni(CN)4]2- is Stable but Labile.
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



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
Cr3+ and Co3+ are generally labile, whereas most second and third row
transition metal ions are inert.
In the case of H2O exchange you can experimentally label with 17O or
18O 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. CaCl2
Ca2+
Ca
H
O
2Cl-
O
ion-dipole interactions.
Lewis acid-base if Ca2+ is assumed to
coordinate to the H2O
H
H
H
Ca2+
O
H
H
We can measure the rates of exchange of water
water:
[M(H2O)n]x+ + 18OH2
[M(H2O)n-118(OH2)]x+ + H2O
Water exchange rates (in text) page 103, Table 5.5
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Rate constants for water exchange for various ions
[M(H2O)n]x+ +
Very slow
K > 10-3 to 10-6 sec-1
18OH
2
[M(H2O)n-118(OH2)]x+ + H2O
Very Fast
K > 108 sec-1
Group 1A – as we go down the group the cations are getting larger and the charge
density decreases so the Mn+-OH2 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 Cu2+ and Cr2+; k >
108. Diffusion controlled reactions
Ionic charge and size are important.
g
Rates of Exchange:
Li+< Na+< K+< Cs+
Group 1A
Be2+< Mg2+< Ca 2+< Sr2+< Ba2+
Group 2A
In these two series, the smallest ions exchange the slowest:
Within a group e.g. Group 1A, size is important:
Cs+•••OH2
Li+•••OH2
+
Stonger
g because the Li
Less strong
g attracting
g
Ion is smaller
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Cs+
O2-
Li+
d
O2-
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+-OH2.
This means the Li-OH2 bond is stronger than the Cs-OH2 since the
Lithium ion holds onto the electrons of the O2- and does not want to
break to release H2O as easily as Cs-OH2
Moving from Group 1A to Group 2A – charge is important
Increasing charge on the ion – H2O 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.
Cu2+ and Cr3+
Cu2+(d9) and Cr2+(d4) 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 eg 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
the electron pair on the ligand and any electrons with
orbitals in the z-component, thus lowering the energy
of the complex.
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Classes 2 and 3 - Includes most of the first row TM ions and the
lanthanides plus Be2+, Al3+, V2+. Rate constants 1 to 108 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. Cr2+, Ni2+,
and Cu2+ have identical radii.

Mn2+ > Fe2+ > Co2+ substitution rates decrease across the series. This is
due to the increase in Zeff
ff across the series and increase in E(M--L).
 d- electron configurations are important because the CFSE will affect
the rates of exchange here.
Class 4 - Rate constants are in the range 10-3-10-6 sec-1. This includes
Cr3+, Co3+, Rh3+, Ir3+, Pt2+.

For these metal ions the rate of exchange is partially related to the size of the
cations and p
partly
y 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.

d3 and d6 ions are predicted to be inert e.g. Rh3+, Cr3+ and Co2+
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