Class 12 Medical Chemistry - Coordination Compounds

On the basis of crystal field theory write the electronic configuration for ${d}^{4}$ if ${\Delta}_{0}< P$ .

1305948_1372300_ans_444d38c051fe431ca0b3e25117bb53ae.PNG

What is $pK$ value of the complex $[PtCl_4]$ if__________.
$Pt^{4+}+Cl^-\rightarrow [PtCl]^{3+}$ $K_1=10^{-2}$
$[PtCl]^{3+}+Cl^-\rightarrow [PtCl_2]^{2+}$ $K_2=10^{-3}$
$[PtCl_2]^{2+}+Cl^-\rightarrow [PtCl_3]^+$ $K_3=10^{-4}$
$[PtCl_3]^++Cl^-\rightarrow [PtCl_4]$ $K_4=10^{-4}$

On adding,

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

$\therefore$ Overall stability constant

$K=K_1K_2K_3K_4=10^{-13}$

$log K=log K_1+log K_2 + log K_3+ log K_4=-13$

$-log K=13$

$pK=13$

Which of the following is more stable complex and why? ${ \left[ Co{ \left( N{ H }_{ 3 } \right) }_{ 6 } \right] }^{ 3+ }$ and ${ \left[ Co{ \left( en \right) }_{ 3 } \right] }^{ 3+ }$

$\displaystyle { \left[ Co{ \left( en \right) }_{ 3 } \right] }^{ 3+ }$ 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.

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

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.

Which isomerism is shown by $[Co(NH_{3})_{6}]Cl_3$ . Write name?

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

$\displaystyle [Co(NH_{3})_{6}]Cl_3$

$\displaystyle [Co(NH_{3})_{6}Cl]Cl_2 \cdot NH_{3}$

$\displaystyle [Co(NH_{3})_{6}Cl_2]Cl \cdot 2NH_{3}$

$\displaystyle [Co(NH_{3})_{6}Cl_3] \cdot 3NH_{3}$

Which one is more stable in the following pairs of complexes? Give reasons.
i) $[Co(NO_{2})_{6}]^{4-}$ and $[Co(NO_{2})_{6}]^{3-}$
ii) $K_{4}[Fe(CN)_{6}]$ and $K_{3}[Fe(CN)_{6}]$ .
(iii) $[CO(H_{2}O)_{6}]^{2+}$ and $[CO(NH_{3})_{6}]^{2+}$

i) Valence shell electronic configuration of

$Co$ is

$4s^23d^7$

Oxidation state of

$Co$ in

$[Co(NO_{2})_{6}]^{4-}$ is

$x+6\times(-1)=-4$ and

$x=+2$ i.e.

$3d^7$

and in

$[Co(NO_{2})_{6}]^{3-}$ is:

$x+6\times(-1)=-3$ and

$x=+3$ i.e.

$3d^6$

Since

$NO_2^-$ is a strong field ligand, therefore it will completely fill the

$t_{ 2g }$ orbital with

$6$ electrons and then fill

$e_g$ orbital.

Therefore

$[Co(NO_{2})_{6}]^{3-}$ will be more stable than

$[Co(NO_{2})_{6}]^{4-}$ complex as all electrons are in low energy orbital.

ii) Valence shell electronic configuration of

$Fe$ is

$4s^23d^6$

Oxidation state of

$Fe$ in

$K_{4}[Fe(CN)_{6}]$ is

$x+6\times(-1)+4=0$ and

$x=+2$ i.e.

$3d^6$

and in

$K_{3}[Fe(CN)_{6}]$ . is:

$x+6\times(-1)+3=0$ and

$x=+3$ i.e.

$3d^5$

Since

$CN^-$ is a strong field ligand, therefore it will completely fill the

$t_{ 2g }$ orbital with

$6$ electrons and then fill

$e_g$ orbital.

Therefore

$K_{4}[Fe(CN)_{6}]$ will be more stable than

$K_{3}[Fe(CN)_{6}]$ complex as it has

$6$ paired electrons and each stabilizes the complex by

$0.4\Delta _{ oh }$ in low energy orbital.

The hexaaquomanganese (II) ion contains five unpaired electrons while the hexacyano ion contains only one unpaired electron. Explain using Crystal Field Theory.

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

$t_2g^5$

$eg^0$ .

Which of the following is not correctly matched?

Complex ion CFSE
a $[Co(NH_3)_6]^{3+}$ $24 \ Dq$
b $[Cr(NH_3)_6]^{3+}$ $12 \ Dq$
c $[FeF_6]^{3-}$ $4 \ Dq$
d $[Fe(CN)_6]^{3-}$ $20 \ Dq$

a) [Co(NH3)6]3+ low spin d6

Therefore, first option is incorrectly matched.

If crystal field stabilization energy of $[ML_6]^{+n}$ is $-0.8 \triangle_0$ . Find minimum number of electrons in $t_{2g}$ orbitals of metal ion?

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.

On the basis of crystal field theory write the electronic configuration for $d^{4}$ ion, if $\Delta_{0} < P$ .

1194499_1318532_ans_a4e1fe1e6887446abacfe5135af7d4cc.jpg

What is crystal field stabilization energy?
On the basic of crystal field theory, write the electronic configuration for $d^{4}$ ion if $\Delta_{o}<P$

1375394_1317751_ans_b2d4ed98f7cb4de7ad96f520b08cf3a7.jpg

Write what are the limitation of crystal field theory.

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:

The assumption that the interaction between metal-ligand is purely electrostatic cannot be said to be very realistic.
This theory takes only d-orbitals of a central atom into account. The s and p orbits are not considered for the study.
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–.
The theory rules out the possibility of having p bonding. This is a serious drawback because p bonding is found in many complexes.
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.

How would you account for the following:
The ${d}^{1}$ configuration is very unstable in ions.

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

List and explain the factors which affect the stability of coordination complexes.

Draw an energy level diagram to show the lifting of the degeneracy of the 3d orbitals in a tetrahedral ligand field.

1555440_1713336_ans_9e6479d786fb45889f3835c9ba6f6f1a.jpg

Explain, why $Mn^{3+}$ is less stable than $Mn^{2+}$ and $Mn^{4+}$ ions?
Explain, why $Mn^{3+}$ disproportionate into $Mn^{2+}$ and $Mn^{4+}$ ions?

1731747_1732016_ans_339de18e79c0411e8fc00c0a5168e546.jpg

Explain why $[Cu(en)_2]^{2+}$ is less stable than $[Fe(EDTA)]^-$ ?

$[Cu(en)_{2}]^{2+}$ 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.

Why is the complex $[Co(en)_{3}]^{3+}$ more stable than the complex $[CoF_{6}]^{3-}$ ?

$[Co(en)_{3}]^{3+}$ is more stable than the complex

$[CoF_{6}]^{3-}$ due to chelate effect as the complex

$[Co(en)_{3}]^{3+}$ contains

the following chelating ligand.

$\ddot { N } H_2-CH_2-CH_2-\ddot { N } H_2$

What is spectrochemical series?

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

The order is:

$I^{-} < Br^{-} < SCN^{-} < Cl^{-} < S^{2-} < F^{-} < OH^{-} < C_{2}O_{4}^{2-} < H_{2}O < NCS < EDTA^{4-}$

$< NH_{3} < en < CN^{-} < CO$

Which of the following complex is more stable and why ?
$[Co(NH_{3})_{6}]^{3+}$ and $[Co(en)_{3}]^{3+}$

$[Co(en)_{3}]^{3+}$ is more stable than

$[Co(NH_{3})_{6}]^{3+}$ because of chelation effect.

Why do compounds having similar geometry have different magnetic moment?

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,

e.g.,

$[CoF_6]^{3-}$ and

$[Co(NH_3)_6]^{3+}$ , the former is paramagnetic and the latter is diamagnetic.

Arrange following complex ions in increasing order of crystal field splitting energy $(\Delta_0)$ :
$[Cr(Cl)_6]^{3-}, [Cr(CN)_6]^{3-}, [Cr(NH_3)_6]^{3+}$

Crystal filed splitting energy increases in the order

$[Cr(Cl)_6]^{3-} < [Cr(CN)_6]^{3+} < [Cr(NH_3)_6]^{3-}$ .

Ethylenediamine displaces $H_2O$ in the complex $[Fe(H_2O)_6]^{3+}$ in three steps.
$[Fe(H_2O)_6]^{3+}+3en\rightarrow [Fe(en)_3]^{3+}+6H_2O$
Stepwise formation constants are
          $log K_1=4.34$
          $log K_2=3.31$
          $log K_3=2.05$
Overall formation constant is $A\times 10^B$ .
$A+B$ is ............

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,

$K_3=10^{2.05}$

Now calculate the overall formation constant.

$K=K_1 \times K_2 \times K_3 = 10^{4.34} \times 10^{3.31} \times10^{2.05} ={10}^{9.7}= 5 \times 10^9$ , this is equal to

$A \times 10^B$ .

Hence,

$A=5$ and

$B=9$

$A+B = 9+5=14$

Ethylenediamine displaces $H_2O$ in the complex $[Fe(H_2O)_6]^{3+}$ in three steps.
$[Fe(H_2O)_6]^{3+}+3en\rightarrow [Fe(en)_3]^{3+}+6H_2O$
Stepwise formation constants are:
$\log K_1=4.34$
$\log K_2=3.31$
$\log K_3=2.05$
Overall formation constant is $A\times 10^{2B}$ .
$\text{B}$ in question is:

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

$\log K_1=4.34$ . Hence,

$K_1=10^{4.34}$

$\log K_2=3.31$ . Hence,

$K_2=10^{3.31}$

$\log K_3=2.05$ . Hence,

$K_3=10^{2.05}$

Now calculate the overall formation constant.

$K=K_1 \times K_2 \times K_3$

$= 10^{4.34} \times 10^{3.31} 10^{2.05}$

$= 5 \times 10^{10}$ .

But this is equal to

$A \times 10^{2B}$ .

Hence,

$B=5$

Which complex is more stable in each pair?

(a) $[Ni(CN)_4]^{2-}$ , $[Pt(CN)_4]^{2-}$ .

$[Pt(CN)_4]^{2-}$ has

$Pt$ in

$+2$ oxidation state. It forms a square planar structure with

$dsp^2$ hybridisation.

$CN^-$ bring a strong field ligand causes pairing of unpaired electrons. Hence,

$[Ni(CN)4]^{-2}$ is more stable than

$Pt(CN)_4]^{2-}$ .

1028773_166188_ans_76f690d87f484614b4cd918b57f2ba26.PNG

What is crystal field splitting energy? How does the magnitude of $\Delta_{0}$ decide the actual configuration of d orbitals in a coordination entity?

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.

What is meant by stability of a coordination compound in solution? State the factors which govern stability of complexes.

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

Find the crystal field stabilization energy (CFSE) (in kJ/mol) for complex, ${\left[ {Ti{{\left( {{H_2}O} \right)}_6}} \right]^{3 + }}$ . According to CFT, the first absorption maximum is obtained at $20,3000\, cm^{-1}$ for the transition

$\begin{array}{l}\Delta E = hc\overline v - 34\\ = 6.63 \times 10Js \times 3.00 \times {10^8}m/s \times 20300c{m^{ - 1}} \times \frac{{100cm}}{{1m}}\\ = 4.037 \times {10^{ - 19}}J\end{array}$
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

$CFSE = 0.4\Delta E$

Explain hydrate (solvate) isomerism and linkage isomerism with suitable examples.

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.

Example :- The complex

$[Cr(H_2O)_6]Cl_3$ (Violet) and its solvate isomer

$[Cr(H_2O)_5Cl] Cl_2H_2O$ (grey- green)

Linkage isomerism arises in a coordination compound containing ambidentate ligands. A simple example is provided by the complexes containing the thiocyanate ligand, NCS which may bind through the nitrogen to give M-NCS or through sulphur to given M-SCN.

The step wise formation of $[Cu(NH_{3})_{4}]^{2+}$ is given below

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

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

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

$[Cu(NH_{3})_{3}]^{2+} + NH_{3} \overset{K_{4}} \rightleftharpoons [Cu(NH_{3})_{4}]^{2+}$

The value of stability constants $K_{1} , K_{2} , K_{3}$ and $K_{4}$ are $10^{4} , 1.58 \times 10^{3} , 5 \times 10^{2}$ and $10^{2}$ respectively. The overall equilibrium constants for dissociation of $[Cu(NH_{3})_{4}]^{2+}$ is $x \times 10^{-12}$ . The value of $x$ is _____________.
(Rounded off to the nearest integer)

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

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

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

$[Cu(NH_{3})_{3}]^{2+} + NH_{3} \overset{K_{4}} \rightleftharpoons [Cu(NH_{3})_{4}]^{2+}$

$K = K_{1} \times K_{2} \times K_{3} \times K_{4}$

$= 10^{4} \times 1.58 \times 10^{3} \times 5 \times 10^{2} \times 10^{2}$

$K = 7.9 \times 10^{11}$

Where

$K \rightarrow$ Equilibrium constant for formation of

$[Cu(NH_{3})_{4}]^{2+}$ . So equilibrium constant

$(K')$ for dissociation of

$[Cu(NH_{3})_{4}]^{2+}$ is

$\dfrac{1}{K}$

$K' = \dfrac{1}{K}$

$K' = \dfrac{1}{7.9 \times 10^{11}} =1.26\times 10^{-12}$

So, the value of

$x = 1.26$

OMR Ans =

$1$ (After rounded off to the nearest integer)

Coordination Compounds - Class 12 Medical Chemistry - Extra Questions

On the basis of crystal field theory write the electronic configuration for ${d}^{4}$ if ${\Delta}_{0}< P$ .

What is $pK$ value of the complex $[PtCl_4]$ if__________.
$Pt^{4+}+Cl^-\rightarrow [PtCl]^{3+}$ $K_1=10^{-2}$
$[PtCl]^{3+}+Cl^-\rightarrow [PtCl_2]^{2+}$ $K_2=10^{-3}$
$[PtCl_2]^{2+}+Cl^-\rightarrow [PtCl_3]^+$ $K_3=10^{-4}$
$[PtCl_3]^++Cl^-\rightarrow [PtCl_4]$ $K_4=10^{-4}$

Which of the following is more stable complex and why? ${ \left[ Co{ \left( N{ H }_{ 3 } \right) }_{ 6 } \right] }^{ 3+ }$ and ${ \left[ Co{ \left( en \right) }_{ 3 } \right] }^{ 3+ }$

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

Which isomerism is shown by $[Co(NH_{3})_{6}]Cl_3$ . Write name?

Which one is more stable in the following pairs of complexes? Give reasons.
i) $[Co(NO_{2})_{6}]^{4-}$ and $[Co(NO_{2})_{6}]^{3-}$
ii) $K_{4}[Fe(CN)_{6}]$ and $K_{3}[Fe(CN)_{6}]$ .
(iii) $[CO(H_{2}O)_{6}]^{2+}$ and $[CO(NH_{3})_{6}]^{2+}$

The hexaaquomanganese (II) ion contains five unpaired electrons while the hexacyano ion contains only one unpaired electron. Explain using Crystal Field Theory.

Which of the following is not correctly matched?

Complex ion CFSE
a $[Co(NH_3)_6]^{3+}$ $24 \ Dq$
b $[Cr(NH_3)_6]^{3+}$ $12 \ Dq$
c $[FeF_6]^{3-}$ $4 \ Dq$
d $[Fe(CN)_6]^{3-}$ $20 \ Dq$

If crystal field stabilization energy of $[ML_6]^{+n}$ is $-0.8 \triangle_0$ . Find minimum number of electrons in $t_{2g}$ orbitals of metal ion?

On the basis of crystal field theory write the electronic configuration for $d^{4}$ ion, if $\Delta_{0} < P$ .

What is crystal field stabilization energy?
On the basic of crystal field theory, write the electronic configuration for $d^{4}$ ion if $\Delta_{o}<P$

Write what are the limitation of crystal field theory.

How would you account for the following:
The ${d}^{1}$ configuration is very unstable in ions.

List and explain the factors which affect the stability of coordination complexes.

Draw an energy level diagram to show the lifting of the degeneracy of the 3d orbitals in a tetrahedral ligand field.

Explain, why $Mn^{3+}$ is less stable than $Mn^{2+}$ and $Mn^{4+}$ ions?
Explain, why $Mn^{3+}$ disproportionate into $Mn^{2+}$ and $Mn^{4+}$ ions?

Explain why $[Cu(en)_2]^{2+}$ is less stable than $[Fe(EDTA)]^-$ ?

Why is the complex $[Co(en)_{3}]^{3+}$ more stable than the complex $[CoF_{6}]^{3-}$ ?

What is spectrochemical series?

Which of the following complex is more stable and why ?
$[Co(NH_{3})_{6}]^{3+}$ and $[Co(en)_{3}]^{3+}$

Why do compounds having similar geometry have different magnetic moment?

Arrange following complex ions in increasing order of crystal field splitting energy $(\Delta_0)$ :
$[Cr(Cl)_6]^{3-}, [Cr(CN)_6]^{3-}, [Cr(NH_3)_6]^{3+}$

Ethylenediamine displaces $H_2O$ in the complex $[Fe(H_2O)_6]^{3+}$ in three steps.
$[Fe(H_2O)_6]^{3+}+3en\rightarrow [Fe(en)_3]^{3+}+6H_2O$
Stepwise formation constants are:
$\log K_1=4.34$
$\log K_2=3.31$
$\log K_3=2.05$
Overall formation constant is $A\times 10^{2B}$ .
$\text{B}$ in question is:

Which complex is more stable in each pair?

(a) $[Ni(CN)_4]^{2-}$ , $[Pt(CN)_4]^{2-}$ .

What is crystal field splitting energy? How does the magnitude of $\Delta_{0}$ decide the actual configuration of d orbitals in a coordination entity?

What is meant by stability of a coordination compound in solution? State the factors which govern stability of complexes.

Find the crystal field stabilization energy (CFSE) (in kJ/mol) for complex, ${\left[ {Ti{{\left( {{H_2}O} \right)}_6}} \right]^{3 + }}$ . According to CFT, the first absorption maximum is obtained at $20,3000\, cm^{-1}$ for the transition

Explain hydrate (solvate) isomerism and linkage isomerism with suitable examples.

Class 12 Medical Chemistry Extra Questions

	Complex ion	CFSE
a	$[Co(NH_3)_6]^{3+}$	$24 \ Dq$
b	$[Cr(NH_3)_6]^{3+}$	$12 \ Dq$
c	$[FeF_6]^{3-}$	$4 \ Dq$
d	$[Fe(CN)_6]^{3-}$	$20 \ Dq$

Coordination Compounds - Class 12 Medical Chemistry - Extra Questions

On the basis of crystal field theory write the electronic configuration for d4{d}^{4} if Δ0<P{\Delta}_{0}< P.

Which of the following is more stable complex and why? [Co(NH3)6]3+{ \left[ Co{ \left( N{ H }_{ 3 } \right) }_{ 6 } \right] }^{ 3+ } and [Co(en)3]3+{ \left[ Co{ \left( en \right) }_{ 3 } \right] }^{ 3+ }

How many moles of AgClAgCl will be precipitated when an excess of AgNO3AgNO_3 solution is added to one molar solution of [CrCl(H2O)5]Cl2[CrCl(H_2O)_5]Cl_2?

Which isomerism is shown by [Co(NH3)6]Cl3[Co(NH_{3})_{6}]Cl_3. Write name?

Which one is more stable in the following pairs of complexes? Give reasons.i) [Co(NO2)6]4−[Co(NO_{2})_{6}]^{4-} and [Co(NO2)6]3−[Co(NO_{2})_{6}]^{3-}ii) K4[Fe(CN)6]K_{4}[Fe(CN)_{6}] and K3[Fe(CN)6]K_{3}[Fe(CN)_{6}].(iii)[CO(H2O)6]2+[CO(H_{2}O)_{6}]^{2+} and [CO(NH3)6]2+[CO(NH_{3})_{6}]^{2+}

The hexaaquomanganese (II) ion contains five unpaired electrons while the hexacyano ion contains only one unpaired electron. Explain using Crystal Field Theory.

Which of the following is not correctly matched?Complex ionCFSEa[Co(NH3)6]3+[Co(NH_3)_6]^{3+}24 Dq24 \ Dqb[Cr(NH3)6]3+[Cr(NH_3)_6]^{3+}12 Dq12 \ Dqc[FeF6]3−[FeF_6]^{3-}4 Dq4 \ Dqd[Fe(CN)6]3−[Fe(CN)_6]^{3-}20 Dq20 \ Dq

If crystal field stabilization energy of [ML6]+n[ML_6]^{+n} is −0.8△0-0.8 \triangle_0. Find minimum number of electrons in t2gt_{2g} orbitals of metal ion?

On the basis of crystal field theory write the electronic configuration for d4d^{4} ion, if Δ0<P\Delta_{0} < P.

What is crystal field stabilization energy?On the basic of crystal field theory, write the electronic configuration for d4d^{4} ion if Δo<P\Delta_{o}<P

Write what are the limitation of crystal field theory.

How would you account for the following:The d1{d}^{1} configuration is very unstable in ions.

List and explain the factors which affect the stability of coordination complexes.

Draw an energy level diagram to show the lifting of the degeneracy of the 3d orbitals in a tetrahedral ligand field.

Explain, why Mn3+Mn^{3+} is less stable than Mn2+Mn^{2+} and Mn4+Mn^{4+} ions?Explain, why Mn3+Mn^{3+} disproportionate into Mn2+Mn^{2+} and Mn4+Mn^{4+} ions?

Explain why [Cu(en)2]2+[Cu(en)_2]^{2+} is less stable than [Fe(EDTA)]−[Fe(EDTA)]^-?

Why is the complex [Co(en)3]3+[Co(en)_{3}]^{3+} more stable than the complex [CoF6]3−[CoF_{6}]^{3-}?

What is spectrochemical series?

Which of the following complex is more stable and why ? [Co(NH3)6]3+[Co(NH_{3})_{6}]^{3+} and [Co(en)3]3+[Co(en)_{3}]^{3+}

Why do compounds having similar geometry have different magnetic moment?

Arrange following complex ions in increasing order of crystal field splitting energy (Δ0)(\Delta_0):[Cr(Cl)6]3−,[Cr(CN)6]3−,[Cr(NH3)6]3+[Cr(Cl)_6]^{3-}, [Cr(CN)_6]^{3-}, [Cr(NH_3)_6]^{3+}

Which complex is more stable in each pair? (a) [Ni(CN)4]2−[Ni(CN)_4]^{2-} , [Pt(CN)4]2−[Pt(CN)_4]^{2-}.

What is crystal field splitting energy? How does the magnitude of Δ0\Delta_{0} decide the actual configuration of d orbitals in a coordination entity?

What is meant by stability of a coordination compound in solution? State the factors which govern stability of complexes.

Find the crystal field stabilization energy (CFSE) (in kJ/mol) for complex, [Ti(H2O)6]3+{\left[ {Ti{{\left( {{H_2}O} \right)}_6}} \right]^{3 + }}. According to CFT, the first absorption maximum is obtained at 20,3000cm−120,3000\, cm^{-1} for the transition

Explain hydrate (solvate) isomerism and linkage isomerism with suitable examples.

Class 12 Medical Chemistry Extra Questions

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On the basis of crystal field theory write the electronic configuration for ${d}^{4}$ if ${\Delta}_{0}< P$ .

Which of the following is more stable complex and why? ${ \left[ Co{ \left( N{ H }_{ 3 } \right) }_{ 6 } \right] }^{ 3+ }$ and ${ \left[ Co{ \left( en \right) }_{ 3 } \right] }^{ 3+ }$

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

Which isomerism is shown by $[Co(NH_{3})_{6}]Cl_3$ . Write name?

Which one is more stable in the following pairs of complexes? Give reasons.
i) $[Co(NO_{2})_{6}]^{4-}$ and $[Co(NO_{2})_{6}]^{3-}$
ii) $K_{4}[Fe(CN)_{6}]$ and $K_{3}[Fe(CN)_{6}]$ .
(iii) $[CO(H_{2}O)_{6}]^{2+}$ and $[CO(NH_{3})_{6}]^{2+}$

Which of the following is not correctly matched?

Complex ion CFSE
a $[Co(NH_3)_6]^{3+}$ $24 \ Dq$
b $[Cr(NH_3)_6]^{3+}$ $12 \ Dq$
c $[FeF_6]^{3-}$ $4 \ Dq$
d $[Fe(CN)_6]^{3-}$ $20 \ Dq$

If crystal field stabilization energy of $[ML_6]^{+n}$ is $-0.8 \triangle_0$ . Find minimum number of electrons in $t_{2g}$ orbitals of metal ion?

On the basis of crystal field theory write the electronic configuration for $d^{4}$ ion, if $\Delta_{0} < P$ .

What is crystal field stabilization energy?
On the basic of crystal field theory, write the electronic configuration for $d^{4}$ ion if $\Delta_{o}<P$

How would you account for the following:
The ${d}^{1}$ configuration is very unstable in ions.

Explain, why $Mn^{3+}$ is less stable than $Mn^{2+}$ and $Mn^{4+}$ ions?
Explain, why $Mn^{3+}$ disproportionate into $Mn^{2+}$ and $Mn^{4+}$ ions?

Explain why $[Cu(en)_2]^{2+}$ is less stable than $[Fe(EDTA)]^-$ ?

Why is the complex $[Co(en)_{3}]^{3+}$ more stable than the complex $[CoF_{6}]^{3-}$ ?

Which of the following complex is more stable and why ?
$[Co(NH_{3})_{6}]^{3+}$ and $[Co(en)_{3}]^{3+}$

Arrange following complex ions in increasing order of crystal field splitting energy $(\Delta_0)$ :
$[Cr(Cl)_6]^{3-}, [Cr(CN)_6]^{3-}, [Cr(NH_3)_6]^{3+}$

Which complex is more stable in each pair?

(a) $[Ni(CN)_4]^{2-}$ , $[Pt(CN)_4]^{2-}$ .

What is crystal field splitting energy? How does the magnitude of $\Delta_{0}$ decide the actual configuration of d orbitals in a coordination entity?

Find the crystal field stabilization energy (CFSE) (in kJ/mol) for complex, ${\left[ {Ti{{\left( {{H_2}O} \right)}_6}} \right]^{3 + }}$ . According to CFT, the first absorption maximum is obtained at $20,3000\, cm^{-1}$ for the transition