Co-Ordination Compounds MCQs with Answers – Part 1 (NEET Chemistry)

Explanation: A coordination entity consists of a central metal atom or ion surrounded by a definite number of directly bonded ligands. In \(\mathrm{[Co(NH_3)_6]Cl_3}\), cobalt is the central metal ion, and six ammonia molecules act as ligands. Each ammonia molecule contains a nitrogen atom with an available lone pair of electrons. Nitrogen donates this electron pair to the central cobalt ion, forming a coordinate bond. The ligand therefore acts as a Lewis base, while the central metal ion acts as a Lewis acid. All six ammonia molecules belong to the coordination sphere, represented by the square brackets. The three chloride ions outside the brackets are counter ions required for electrical neutrality. They are not the electron-pair donors responsible for the six metal–ammonia coordinate bonds. Thus, coordinate bonding is correctly represented by electron-pair donation from ammonia to cobalt.

Explanation: The electrical charge of a complete coordination compound must be distinguished from the charge of its coordination entity. In \(\mathrm{[Ni(CO)_4]}\), carbon monoxide is a neutral ligand, and the nickel atom has oxidation state zero. Consequently, the complete coordination entity has zero net charge. It is nevertheless a genuine coordination entity because four CO molecules are coordinated to the nickel atom. In \(\mathrm{[Co(NH_3)_6]Cl_3}\), the three chloride ions outside the brackets contribute a total charge of \(-3\). The bracketed entity must therefore have charge \(+3\) to make the entire compound electrically neutral. Ammonia is neutral and does not reduce that positive charge. Thus, \(\mathrm{[Co(NH_3)_6]^{3+}}\) is a complex cation, whereas \(\mathrm{[Ni(CO)_4]}\) is a neutral complex. The presence of coordination bonding does not require the coordination entity itself to be ionic.

Explanation: The square brackets identify the coordination sphere, which contains the iron centre and six coordinated cyanide ligands. Potassium ions occur outside the coordination sphere and function as counter ions. On dissolution, these potassium ions can separate from the complex anion without breaking the iron–cyanide coordination bonds. The ordinary ionic dissociation is therefore: \[ \mathrm{K_4[Fe(CN)_6]\longrightarrow 4K^+ + [Fe(CN)_6]^{4-}} \] The four potassium ions contribute a total charge of \(+4\). The intact complex anion carries charge \(-4\), maintaining electrical neutrality. Within the complex, the oxidation state of iron is \(+2\), because six cyanide ligands contribute a total charge of \(-6\). The oxidation state of iron must not be confused with the charge of the entire coordination entity. Although the formula contains six cyanide groups, they are not released as six independent cyanide ions during this ordinary dissociation. Thus, each formula unit produces five ionic species in the idealised dissociation model: four potassium ions and one intact complex anion.

Explanation: The Assertion is true because the six cyanide groups in \(\mathrm{K_4[Fe(CN)_6]}\) are directly coordinated to the central iron ion. These groups form part of the complex anion rather than existing as ordinary counter ions. During dissolution, potassium ions separate from the coordination entity, but the coordinated cyanide groups are not independently liberated by this process. The Reason is therefore also true. It identifies the essential distinction between dissociation of a coordination salt and decomposition of its complex ion. The Reason directly explains why the dissolved compound cannot be represented as free iron(II) and cyanide ions in the stated ratio. This behaviour distinguishes coordination compounds from double salts such as Mohr's salt. A double salt dissociates into its constituent simple ions when dissolved in water, whereas a coordination entity generally preserves its identity under comparable conditions. Thus, the behaviour of the bracketed entity provides the decisive explanation for the Assertion.

Explanation: \(\textbf{Charge of the coordination entity:}\) The complete compound is electrically neutral. The sulphate ion lies outside the coordination sphere and carries charge \(-2\). Therefore, the bracketed coordination entity must carry charge \(+2\). \[ \mathrm{[M(NH_3)_5Cl]^{2+}} \] \(\textbf{Oxidation-state relation:}\) Let the oxidation state of M be \(x\). Each ammonia molecule is neutral, while the coordinated chloride contributes \(-1\). \[ x+5(0)+(-1)=+2 \] \[ x-1=2 \] \[ x=+3 \] \(\textbf{Ionic dissociation:}\) The coordinated chloride remains inside the coordination sphere. The sulphate counter ion separates from the complex cation. \[ \mathrm{[M(NH_3)_5Cl]SO_4\longrightarrow [M(NH_3)_5Cl]^{2+}+SO_4^{2-}} \] One complex cation and one sulphate anion are produced per formula unit. The combined ionic charge is \(+2-2=0\), confirming electrical consistency. \(\textbf{Final answer:}\) \[ \boxed{x=+3,\quad N_{\mathrm{ions}}=2} \]

Explanation: Werner proposed that a metal in a coordination compound exhibits primary and secondary valences. Primary valence broadly corresponds to the oxidation state of the metal, whereas secondary valence corresponds to its coordination number. In \(\mathrm{[Co(NH_3)_4Cl_2]Cl}\), the chloride ion outside the coordination sphere has charge \(-1\). Therefore, the coordination entity carries charge \(+1\). Since the four ammonia molecules are neutral and two coordinated chloride ions contribute \(-2\), cobalt has oxidation state \(+3\). Thus, the primary valence is three. The coordination sphere contains four ammonia molecules and two chloride ions, giving six directly coordinated groups. Therefore, the secondary valence is six. The two coordinated chloride ions participate in satisfying the secondary valence while also contributing towards satisfying the primary valence. The remaining primary-valence requirement is satisfied by the chloride counter ion outside the brackets. Hence, it is incorrect to assume that only external anions satisfy primary valence in Werner's formulation.

Explanation: \(\textbf{Interpretation of precipitation:}\) The compound contains three chloride ions per cobalt atom. Only chloride ions outside the coordination sphere are precipitated as \(\mathrm{AgCl}\) under the stated conditions. Two moles of \(\mathrm{AgCl}\) from one mole of compound indicate two ionisable chloride ions. The remaining chloride ion must be coordinated to cobalt. \(\textbf{Coordination-sphere construction:}\) Five ammonia molecules and one coordinated chloride provide six directly attached groups. This agrees with Werner's secondary valence of six. The coordination formula is: \[ \mathrm{[Co(NH_3)_5Cl]Cl_2} \] Cobalt has oxidation state \(+3\), and the coordinated chloride contributes \(-1\). Consequently, the bracketed entity has charge \(+2\). \(\textbf{Dissociation behaviour:}\) \[ \mathrm{[Co(NH_3)_5Cl]Cl_2\longrightarrow [Co(NH_3)_5Cl]^{2+}+2Cl^-} \] One complex cation and two chloride counter ions are obtained. Thus, the ideal dissociation produces three ionic species per formula unit. The precipitation result, coordination number and ionic charge are mutually consistent.

Explanation: \(\textbf{Initial coordination sphere:}\) The original compound is: \[ \mathrm{[Co(NH_3)_5Cl]Cl_2} \] Five neutral ammonia ligands and one chloride ligand surround cobalt. With cobalt in oxidation state \(+3\), the bracketed entity carries charge \(+2\). Two chloride ions remain outside the coordination sphere. Therefore, one mole of the original compound produces two moles of \(\mathrm{AgCl}\) under the specified precipitation conditions. Its ideal dissociation produces one complex cation and two chloride ions, giving three ions. \(\textbf{Effect of ligand substitution:}\) A neutral ammonia molecule replaces the negatively charged coordinated chloride. The cobalt oxidation state remains unchanged, but the negative charge within the coordination sphere decreases by one unit. Thus, the charge on the coordination entity increases from \(+2\) to \(+3\). The resulting compound is: \[ \mathrm{[Co(NH_3)_6]Cl_3} \] \(\textbf{New precipitation behaviour:}\) All three chloride ions now occur outside the coordination sphere. \[ \mathrm{[Co(NH_3)_6]Cl_3\longrightarrow [Co(NH_3)_6]^{3+}+3Cl^-} \] One mole of the new compound therefore gives three moles of \(\mathrm{AgCl}\) with excess silver nitrate. Its ideal dissociation produces four ionic species per formula unit. Both changes arise from replacing one coordinated anionic ligand with a neutral ligand without changing the cobalt oxidation state. \(\textbf{Final answer:}\) The \(\mathrm{AgCl}\) yield increases from \(2\) to \(3\) moles per mole, and the ideal ionic-species count increases from \(3\) to \(4\).

Explanation: \(\textbf{Common coordination entity:}\) Both compounds contain the same complex cation: \[ \mathrm{[Co(NH_3)_6]^{3+}} \] The cobalt coordination sphere therefore remains unchanged. However, chloride and sulphate differ in ionic charge, so the numbers of counter ions required for neutrality differ. \(\textbf{Chloride solution:}\) Each formula unit of \(\mathrm{[Co(NH_3)_6]Cl_3}\) contains one cobalt atom. \[ n_{\mathrm{salt}}=0.20\,\mathrm{mol} \] Its ideal dissociation is: \[ \mathrm{[Co(NH_3)_6]Cl_3\longrightarrow [Co(NH_3)_6]^{3+}+3Cl^-} \] Each formula unit gives four ionic species. \[ n_{\mathrm{ions,Cl}}=(0.20)(4)=0.80\,\mathrm{mol} \] \(\textbf{Sulphate solution:}\) Each formula unit of \(\mathrm{[Co(NH_3)_6]_2(SO_4)_3}\) contains two cobalt atoms. Hence, the amount of sulphate salt required to supply \(0.20\) mol of cobalt is: \[ n_{\mathrm{salt}}=\frac{0.20}{2}=0.10\,\mathrm{mol} \] Its ideal dissociation is: \[ \mathrm{[Co(NH_3)_6]_2(SO_4)_3\longrightarrow 2[Co(NH_3)_6]^{3+}+3SO_4^{2-}} \] Each formula unit produces five ionic species. \[ n_{\mathrm{ions,SO_4}}=(0.10)(5)=0.50\,\mathrm{mol} \] \(\textbf{Comparison:}\) \[ \frac{n_{\mathrm{ions,Cl}}}{n_{\mathrm{ions,SO_4}}} =\frac{0.80}{0.50}=\frac{8}{5} \] The ratio differs from the ratio of ionic-species counts per formula unit because the sulphate salt contains two cobalt atoms per formula unit. \(\textbf{Final answer:}\) \[ \boxed{8:5} \]

Explanation: \(\textbf{Information from the coordination formulae:}\) Compound P contains two chloride counter ions and therefore yields two moles of \(\mathrm{AgCl}\) per mole. Compound Q contains one chloride counter ion and yields one mole of \(\mathrm{AgCl}\) per mole. Let the amounts of P and Q be \(x\) and \(y\) mol, respectively. \(\textbf{Total amount relation:}\) \[ x+y=0.30 \] \(\textbf{Precipitation relation:}\) Each mole of P contributes two moles of \(\mathrm{AgCl}\), while each mole of Q contributes one mole. \[ 2x+y=0.48 \] Subtracting the first equation from the second gives: \[ x=0.18\,\mathrm{mol} \] Hence: \[ y=0.30-0.18=0.12\,\mathrm{mol} \] \(\textbf{Ionic species from P:}\) Compound P dissociates into one complex cation and two chloride ions. \[ \mathrm{[Co(NH_3)_5Cl]Cl_2\longrightarrow [Co(NH_3)_5Cl]^{2+}+2Cl^-} \] Thus, three ionic species are obtained per formula unit. \[ n_{\mathrm{ions,P}}=3(0.18)=0.54\,\mathrm{mol} \] \(\textbf{Ionic species from Q:}\) Compound Q dissociates into one complex cation and one chloride ion. \[ \mathrm{[Co(NH_3)_4Cl_2]Cl\longrightarrow [Co(NH_3)_4Cl_2]^++Cl^-} \] Thus, two ionic species are obtained per formula unit. \[ n_{\mathrm{ions,Q}}=2(0.12)=0.24\,\mathrm{mol} \] \(\textbf{Total ionic-species amount:}\) \[ n_{\mathrm{total}}=0.54+0.24=0.78\,\mathrm{mol} \] As a consistency check, each formula unit supplies one complex cation, giving \(0.30\) mol of complex cations in total. The mixture also supplies \(0.48\) mol of ionisable chloride ions, consistent with the observed precipitation. Their sum is again \(0.78\) mol. \(\textbf{Final answer:}\) \[ \boxed{0.78\,\mathrm{mol}} \]

Explanation: \(\textbf{Spatial relationship of the ligands:}\) In an octahedral coordination entity, six ligand positions surround the central metal along three mutually perpendicular axes. Two ligands occupying adjacent positions form an angle of \(90^\circ\) at the central metal. Two ligands occupying opposite positions form an angle of \(180^\circ\). In form P, the chloride ligands occupy the positive \(x\) and positive \(y\) directions. These directions are perpendicular, so P has the cis arrangement. In form Q, the chloride ligands occupy the positive and negative \(z\) directions. These positions are opposite, so Q has the trans arrangement. \(\textbf{Relationship with Werner's evidence:}\) Both structures contain four coordinated ammonia molecules and two coordinated chloride ions. Each therefore satisfies a secondary valence of six. The different ligand positions change the geometrical arrangement without changing the coordination formula. Both compounds also contain one chloride ion outside the brackets. Consequently, each gives one mole of \(\mathrm{AgCl}\) per mole of compound with excess silver nitrate under the specified conditions. Their ideal dissociation patterns are also identical, producing one complex cation and one chloride counter ion per formula unit. Thus, precipitation and simple ionic-species counting do not distinguish these geometrical isomers. The decisive difference is the relative spatial orientation of the two coordinated chloride ligands.

Explanation: Coordination number specifies how many ligand donor atoms are directly attached to the central metal. It does not, by itself, uniquely determine the geometry for every coordination number. In \(\mathrm{[Ni(CO)_4]}\), four carbon monoxide ligands coordinate to nickel, and the entity has tetrahedral geometry. The four donor atoms occupy positions directed towards the corners of a tetrahedron. In \(\mathrm{[PtCl_4]^{2-}}\), four chloride ligands coordinate to platinum, but the entity has square-planar geometry. Its four donor atoms occupy positions around platinum in the same plane, approximately \(90^\circ\) apart. Both entities therefore possess coordination number four despite their different spatial arrangements. Werner recognised characteristic geometrical arrangements of secondary valences and included tetrahedral and square-planar structures among common coordination geometries. The example shows why counting coordinated ligands is not sufficient to choose between these two geometries. Additional information about the metal and its bonding environment is needed to determine which arrangement occurs.

Explanation: \(\textbf{Palladium compound:}\) The empirical composition contains four ammonia molecules and two chloride ions per palladium atom. Its secondary valence is four, so the four ammonia molecules can occupy all four coordination positions. The observation of two moles of \(\mathrm{AgCl}\) per mole of compound confirms that both chlorides are outside the coordination sphere. The appropriate coordination formula is: \[ \mathrm{[Pd(NH_3)_4]Cl_2} \] The bracketed entity has charge \(+2\), while the two external chloride ions maintain electrical neutrality. \(\textbf{Platinum compound:}\) The empirical composition contains two ammonia molecules and two chloride ions. Its secondary valence is also four. Because no chloride is precipitated as \(\mathrm{AgCl}\) under the specified conditions, neither chloride behaves as an external counter ion. Both chloride ions must therefore occupy coordination positions along with the two ammonia molecules. The appropriate formula is: \[ \mathrm{[Pt(NH_3)_2Cl_2]} \] Platinum has oxidation state \(+2\), and the two coordinated chloride ions contribute a total charge of \(-2\). The complete coordination entity is therefore neutral. \(\textbf{Combined inference:}\) Both metals have primary valence two and secondary valence four. However, the palladium compound has a complex cation and external chloride counter ions, while the platinum compound has a neutral coordination entity. This demonstrates that identical primary and secondary valence values do not require identical positions of anionic and neutral ligands within the coordination formula. The precipitation observations distinguish the two arrangements of coordination and counter-ion bonding.

Explanation: \(\textbf{Construction of the coordination sphere:}\) The empirical composition contains one platinum atom and six chloride ions. All six chloride ions occupy coordination positions around platinum. The coordination entity is therefore: \[ \mathrm{[PtCl_6]^{q}} \] \(\textbf{Charge calculation:}\) Platinum has oxidation state \(+4\). Each coordinated chloride contributes \(-1\). \[ q=(+4)+6(-1) \] \[ q=4-6=-2 \] Thus, the coordination entity is \(\mathrm{[PtCl_6]^{2-}}\). \(\textbf{Counter-ion requirement:}\) The empirical composition also provides two hydrogen ions as counter ions. The coordination formulation is: \[ \mathrm{H_2[PtCl_6]} \] The two hydrogen ions balance the \(-2\) charge of the coordination entity. \(\textbf{Ideal dissociation:}\) \[ \mathrm{H_2[PtCl_6]\longrightarrow 2H^+ + [PtCl_6]^{2-}} \] One formula unit produces two hydrogen ions and one complex anion. Therefore, three ionic species are obtained per formula unit. All six chloride ions remain within the coordination entity under the stated precipitation conditions. The absence of free chloride is therefore compatible with the compound producing ions on dissolution. This example demonstrates that a coordination compound can contain an anionic coordination entity with positively charged counter ions. \(\textbf{Final answer:}\) \[ \boxed{q=-2,\quad N_{\mathrm{ions}}=3} \]

Explanation: Statement I is correct because Werner distinguished primary valences from secondary valences based on their chemical roles. Primary valences broadly correspond to the oxidation state of the central metal and are normally ionisable. Secondary valences correspond to the number of directly coordinated ligand groups and are non-ionisable in Werner's formulation. Statement II is also correct because Werner proposed definite spatial arrangements of the groups satisfying secondary valences. Such arrangements allow compounds with identical compositions to possess different spatial structures. This provides a structural explanation for geometrical isomerism, as illustrated by different arrangements of coordinated chloride ligands in octahedral cobalt compounds. Statement III is incorrect because Werner's theory does not describe the detailed electronic distribution responsible for magnetic behaviour. Quantitative spin-only magnetic moments require information about unpaired electrons, which is addressed using electronic bonding approaches. Statement IV is incorrect because splitting of d-orbital energies is a feature of crystal-field theory rather than Werner's original structural theory. Werner's model explains coordination formulae, ionisable and coordinated groups, and characteristic geometrical arrangements. It does not provide a detailed electronic account of metal–ligand bonding, magnetic properties or optical transitions. Therefore, only Statements I and II are correct.

Explanation: \(\textbf{Secondary-valence requirement:}\) Cobalt has secondary valence six, so six monodentate groups must be directly coordinated to the central ion. The empirical formula supplies three ammonia molecules per cobalt atom. These three neutral molecules occupy three coordination positions. The remaining three positions must therefore be occupied by chloride ions. \(\textbf{Coordination formula:}\) The required coordination entity is: \[ \mathrm{[Co(NH_3)_3Cl_3]} \] All three chloride ions are inside the coordination sphere. The compound consequently has no external chloride counter ions. \(\textbf{Charge consistency:}\) Cobalt has oxidation state \(+3\). The three ammonia ligands contribute zero charge. The three coordinated chloride ions contribute a total charge of \(-3\). \[ q_{\mathrm{entity}}=(+3)+3(0)+3(-1) \] \[ q_{\mathrm{entity}}=0 \] The coordination entity is electrically neutral. Its neutrality is consistent with the absence of counter ions. \(\textbf{Precipitation inference:}\) Under the specified Werner precipitation conditions, only ionisable chloride outside the coordination sphere contributes directly to \(\mathrm{AgCl}\) formation. Since no chloride ion occurs outside the brackets, the expected precipitation is: \[ n_{\mathrm{AgCl}}=0\,\mathrm{mol} \] per mole of the compound. \(\textbf{Verification of the alternatives:}\) A structure containing only two coordinated chlorides would provide five coordination positions from the stated three ammonia molecules and two chloride ligands, contradicting secondary valence six. Likewise, structures with one or zero coordinated chlorides would leave still more coordination positions unoccupied. The proposed neutral entity alone simultaneously satisfies the empirical composition, the cobalt oxidation state, electrical neutrality, the fixed secondary valence and the precipitation behaviour. This illustrates how Werner's valence requirements can be used together to reconstruct a coordination formula rather than merely identify the number of counter ions.

Explanation: \(\textbf{Dissociation of Mohr's salt:}\) Mohr's salt is a double salt consisting of ferrous sulphate and ammonium sulphate in a definite stoichiometric proportion. Double salts dissociate into their constituent simple ions when dissolved in water. Ignoring the neutral water molecules of crystallisation, its dissociation can be represented as: \[ \mathrm{FeSO_4\cdot(NH_4)_2SO_4\longrightarrow Fe^{2+}+2NH_4^++2SO_4^{2-}} \] The number of ionic species per formula unit is: \[ N_1=1+2+2=5 \] The iron is present as an independently dissociated \(\mathrm{Fe^{2+}}\) ion. \(\textbf{Dissociation of the coordination salt:}\) Potassium hexacyanidoferrate(II) contains the coordination entity \(\mathrm{[Fe(CN)_6]^{4-}}\). The six cyanide ligands remain coordinated to iron during ordinary dissolution. \[ \mathrm{K_4[Fe(CN)_6]\longrightarrow 4K^++[Fe(CN)_6]^{4-}} \] The number of ionic species per formula unit is: \[ N_2=4+1=5 \] Although iron has oxidation state \(+2\) in this coordination entity, it is not present as an independently dissociated \(\mathrm{Fe^{2+}}\) ion. \(\textbf{Comparison:}\) Both salts produce the same ideal number of ionic species per formula unit. However, the chemical identities of those ions are different. Mohr's salt releases simple iron(II) ions, whereas potassium hexacyanidoferrate(II) retains iron within a complex anion. Equal ionic-species counts therefore do not imply identical solution chemistry. This distinction is fundamental to differentiating double salts from coordination compounds.

Explanation: Statement I is correct because carnallite is a double salt containing potassium chloride and magnesium chloride. On ideal dissolution, its constituent ions are one potassium ion, one magnesium ion and three chloride ions. Consequently, it produces two cations and three anions per formula unit. Statement II is incorrect because potash alum produces one potassium ion, one aluminium ion and two sulphate ions. It therefore gives two cations and two anions, not three cations and one anion. Statement III is correct because both carnallite and potash alum are standard examples of double salts. Their constituent simple ions become available in aqueous solution rather than remaining together as a characteristic coordination entity. This contrasts with the behaviour of the complex anion in potassium hexacyanidoferrate(II). Statement IV is also correct because water of crystallisation does not contribute additional charged species merely by being released into the solvent. Counting these water molecules as ions would incorrectly inflate the predicted dissociation count. The distinction between hydrated composition and ionic dissociation must therefore be maintained. The correct combination is Statements I, III and IV only.

Explanation: \(\textbf{Oxidation states in the two coordination entities:}\) In \(\mathrm{[Fe(CN)_6]^{4-}}\), each cyanide ligand has charge \(-1\). Let the oxidation state of iron be \(x\). \[ x+6(-1)=-4 \] \[ x=+2 \] In \(\mathrm{[Fe(CN)_6]^{3-}}\), the corresponding relation is: \[ x+6(-1)=-3 \] \[ x=+3 \] Thus, iron has oxidation state \(+2\) in the first compound and \(+3\) in the second. The number of coordinated cyanide ligands remains six in both entities. \(\textbf{First solution:}\) \[ \mathrm{K_4[Fe(CN)_6]\longrightarrow4K^++[Fe(CN)_6]^{4-}} \] For \(0.15\) mol of compound: \[ n(\mathrm{K^+})_1=4(0.15)=0.60\,\mathrm{mol} \] \[ n(\mathrm{ions})_1=5(0.15)=0.75\,\mathrm{mol} \] \(\textbf{Second solution:}\) \[ \mathrm{K_3[Fe(CN)_6]\longrightarrow3K^++[Fe(CN)_6]^{3-}} \] For \(0.15\) mol of compound: \[ n(\mathrm{K^+})_2=3(0.15)=0.45\,\mathrm{mol} \] \[ n(\mathrm{ions})_2=4(0.15)=0.60\,\mathrm{mol} \] \(\textbf{Changes between solutions:}\) \[ \Delta n(\mathrm{K^+})=0.45-0.60=-0.15\,\mathrm{mol} \] \[ \Delta n(\mathrm{ions})=0.60-0.75=-0.15\,\mathrm{mol} \] The higher oxidation state of iron makes the second complex anion less negatively charged. Therefore, fewer potassium counter ions are required to maintain electrical neutrality in the corresponding salt. The six metal–cyanide coordination attachments remain unchanged. \(\textbf{Final answer:}\) \[ \boxed{\Delta n(\mathrm{K^+})=-0.15\,\mathrm{mol},\quad \Delta n(\mathrm{ions})=-0.15\,\mathrm{mol}} \]

Explanation: The coordination entity contains two neutral ammonia ligands and two negatively charged chloride ligands. Since the complete entity is neutral, the metal must have oxidation state \(+2\). Each ligand is monodentate, so the four ligands provide four direct donor attachments to the central metal. Werner's secondary valence is therefore four. A coordination number of four can correspond to either tetrahedral or square-planar geometry. However, the observed geometrical isomerism provides additional structural evidence. In a tetrahedral arrangement containing two ligands of each type, all ligand positions are equivalent with respect to the relative arrangement of the two identical ligands. Consequently, the tetrahedral \(\mathrm{MA_2B_2}\) arrangement does not produce distinct cis and trans geometrical isomers. In a square-planar arrangement, the two chloride ligands may occupy adjacent positions or opposite positions. These arrangements correspond to cis and trans geometrical isomers, respectively. The observation of two geometrical isomers therefore supports square-planar geometry. This reasoning combines charge balance, secondary-valence counting and the geometrical consequences of coordination-polyhedron shape.