Coordination Compounds - Class 12 Medical Chemistry - Extra Questions

On adding,  $$Pt^{4+}+4Cl^-\rightarrow PtCl_4$$

$$\displaystyle  { \left[ Co{ \left( en \right)  }_{ 3 } \right]  }^{ 3+ }$$  is more stable than $$\displaystyle { \left[ Co{ \left( N{ H }_{ 3 } \right)  }_{ 6 } \right]  }^{ 3+ } $$ as $$\displaystyle  { \left[ Co{ \left( en \right)  }_{ 3 } \right]  }^{ 3+ }$$  is a metal chelate due to presece of bidentate ligand ethylene diamine ligand. Metal chelates are more stable than complexes containing only monodentate lignads.

Two moles of $$AgCl$$ will be precipitated when an excess of $$\displaystyle AgNO_3$$ solution is added to one molar solution of $$\displaystyle [CrCl(H_2O)_5]Cl_2$$

$$\displaystyle [CrCl(H_2O)_5]Cl_2+2AgNO_3 \rightarrow [CrCl(H_2O)_5](NO_3)_2 +2AgCl\downarrow$$

Only chloride ions present in ionization sphere will be precipitated. The chlorine present in coordination sphere will not be precipitated.

 $$[Co(NH_{3})_{6}]Cl_3$$ shows solvate isomerism. There is exchange of ammonia molecules inside and outside co-ordination sphere.
$$ \displaystyle [Co(NH_{3})_{6}]Cl_3$$
$$ \displaystyle [Co(NH_{3})_{6}Cl]Cl_2  \cdot NH_{3}  $$
$$ \displaystyle [Co(NH_{3})_{6}Cl_2]Cl \cdot  2NH_{3}$$
$$ \displaystyle [Co(NH_{3})_{6}Cl_3]   \cdot 3NH_{3}$$

Mn(II) has an electronic configuration of $$3d^54s^0$$. In the case of hexaaquomanganese (II) ion the ligand attached to metal is water which is a weak field ligand and is not able to pair up the electrons of 3d subshell and the configuration is $$t_2g^3$$ and $$eg^2$$. whereas in the case of hexacyano ion the $$CN^-$$ is a strong field ligand and it causes pairing of electrons in the d orbitals and hence thee exist only one unpaired electron in $$t_2g^5$$$$eg^0$$.

a) [Co(NH3)6]3+[Co(NH3)4]3+ low spin d6d6

Therefore, first option is incorrectly matched.

The CFSE in the case of octahedral complexes are found using the formula -0.4$$\Delta_o$$$$\times$$number of electrons in $$t_2g$$ orbital + 0.6$$\Delta_o$$x number of electrons in eg orbitals. Since CFSE is given to be -0.8$$\Delta_o$$ this means the number of electrons in $$t_2g$$ orbital  is 2 and there is no electron in eg orbital.

The crystal field theory is highly useful and more significant as compared to the valence bond theory. Even after such useful properties, it has many limitations. The following points will clearly state the limitations of crystal field theory:

1. The assumption that the interaction between metal-ligand is purely electrostatic cannot be said to be very realistic.
2. This theory takes only d-orbitals of a central atom into account. The s and p orbits are not considered for the study.
3. The theory fails to explain the behavior of certain metals which cause large splitting while others show small splitting. For example, the theory has no explanation as to why H2O is a stronger ligand as compared to OH–.
4. The theory rules out the possibility of having p bonding. This is a serious drawback because p bonding is found in many complexes.
5. The theory gives no significance to the orbits of the ligands. Therefore it cannot explain any properties related to ligand orbitals and their interaction with metal orbitals.

The ions in $$d^1 $$ configuration tend to lose one more electrons to get into stable $${d}^{0}$$ configuration. Also, the hydration or lattice energy is more than sufficient to remove the only electron present in the d-orbital of these ions. Therefore, they act as reducing agents

a

$$[Cu(en)_{2}]^{2+}$$ is less stable than $$[Fe(EDTA)]^{-}$$ due to chelating effect of EDTA. EDTA is a versatile chelating agent. It can form four or six bonds with a metal ion and it forms chelates with both transition-metal ions and main group ions. The chelate effect is in complexes resulting from coordination with the chelating ligand and is much more stable than complexes with non-chelating ligands.

$$[Co(en)_{3}]^{3+}$$ is more stable than the complex $$[CoF_{6}]^{3-}$$ due to chelate effect as the complex $$[Co(en)_{3}]^{3+}$$ contains 

The series in which ligands are arranged in the order of increasing field strength is called spectrochemical series. 

$$[Co(en)_{3}]^{3+}$$ is more stable than $$[Co(NH_{3})_{6}]^{3+}$$ because of chelation effect.

It is due to the presence of weak and strong ligands in complexes. If CFSE is high, the complex will show low value of magnetic moment and vice versa, 

Crystal filed splitting energy increases in the order

First obtain the stepwise formation constants from logK values.
$$logK_1=4.34$$. Hence, $$K_1=10^{4.34}$$
$$logK_2=3.31$$. Hence, $$K_2=10^{3.31}$$
$$logK_3=2.05$$. Hence, $$K_3=10^{2.05}$$

First, obtain the stepwise formation constants from log K values.

The crystal field stabilisation energy (CFSE)  is the gain in the energy achieved by preferential filling up of orbitals by electrons. It is usually less than or equal to 0. When it is equal to 0, the complex is unstable. The magnitude of CFSE depends on the number and nature of ligands and the geometry of the complex.

Consider octahedral $$\displaystyle d^4$$ system. Three electrons are in lower $$\displaystyle t_{2g}$$ level. Fourth electron will enter into higher $$\displaystyle e_g$$ level if $$\displaystyle \Delta _o < P$$. Fourth electron will enter into lower $$\displaystyle t_{2g}$$ level if $$\displaystyle \Delta _o > P$$. Here, P is the pairing energy. It is the energy required to pair two electrons against electron electron repulsion in the same orbital.

The stability of coordination compounds in solution means the degree of association between the metal ion and the ligands involved in the state of equilibrium. Quantitatively, the stability is expressed by the equilibrium constant for the association.
$$\displaystyle Cu^{2+} (aq) + 4NH_3(aq) \rightleftharpoons [Cu(NH_3)_4^{2+}] (aq) \\ K = \dfrac {[Cu(NH_3)_4]^{2+}}{[Cu^{2+}]  [NH_3]^4}$$
Factors affecting the stability of a complex ion:
(a) The nature of the central metal atom
(1) Charge on the central metal ion
(2) Size of metal ion
(3) Electronegativity or charge distribution of metal ion
(4) Chelate effect
(5) Macrocylic effect
(b) The nature of the ligand
(1) Basic strength
(2) Size and charge of ligands

$$\begin{array}{l}\Delta E = hc\overline v  - 34\\ = 6.63 \times 10Js \times 3.00 \times {10^8}m/s \times 20300c{m^{ - 1}} \times \frac{{100cm}}{{1m}}\\ = 4.037 \times {10^{ - 19}}J\end{array}$$
This is energy change for one ion 
For one mole
$$\begin{array}{l}\Delta E = 4.037 \times {10^{ - 19}}J/ion \times 6.02 \times {10^{23}}ions/mol\\ = 243067J/mol\end{array}$$
But, $$1kJ = 1000J$$
$$\therefore \Delta E = 243kJ/mol$$
But, $$T_1$$ ion has d' electron configuration
$$CFSE = 0.4\Delta E$$

Solvate isomerism is known as 'hydrate isomerism' in case where water is involved as a solvent. Solvate isomers differ by whether or not a solvent molecule is directly bonded to the metal ion or merely present as free solvent molecules in the crystal lattice.

$$Cu^{2+} + NH_{3} \overset{K_{1}} \rightleftharpoons [Cu(NH_{3})]^{2+} $$