Coordination Compounds Notes

Chemistry 118

• Transition metals form complex compounds by sharing electrons with anions or neutral molecules.

• These are now called coordination compounds, an important area of inorganic chemistry.

• Coordination compounds contribute to our understanding of chemical bonding and molecular structure.

• They are vital in biological systems (chlorophyll, haemoglobin, vitamin B12).

• They are used in metallurgical processes, industrial catalysts, and analytical reagents.

• Applications: electroplating, textile dyeing, and medicinal chemistry.

Objectives

• Understand Werner’s theory of coordination compounds.

• Define coordination entity, central atom/ion, ligand, coordination number, coordination sphere, coordination polyhedron, oxidation number, homoleptic and heteroleptic.

• Learn nomenclature rules for coordination compounds.

• Write formulas and names of mononuclear coordination compounds.

• Define isomerism in coordination compounds.

• Understand bonding using Valence Bond and Crystal Field theories.

• Appreciate the importance and applications of coordination compounds.

Coordination Compounds

• Alfred Werner (1866-1919) was the first to formulate ideas about the structures of coordination compounds.

• He prepared and characterized compounds, studying their physical and chemical behavior.

• Werner proposed primary and secondary valences for metal ions.

• Binary compounds like $CrCl<em>3$, $CoCl</em>2$, and $PdCl_2$ have primary valences of 3, 2, and 2, respectively.

• In cobalt(III) chloride compounds with ammonia, some chloride ions precipitate as $AgCl$ with silver nitrate, while others remain in solution.

Werner’s Theory Observations

• $1 mol$ $CoCl<em>3 "." 6NH</em>3$ (Yellow) gave $3 mol$ $AgCl$.

• $1 mol$ $CoCl<em>3 "." 5NH</em>3$ (Purple) gave $2 mol$ $AgCl$.

• $1 mol$ $CoCl<em>3 "." 4NH</em>3$ (Green) gave $1 mol$ $AgCl$.

• $1 mol$ $CoCl<em>3 "." 4NH</em>3$ (Violet) gave $1 mol$ $AgCl$.

• Six groups (chloride ions or ammonia molecules) remain bonded to the cobalt ion during the reaction.

• Compounds are formulated with atoms within square brackets forming a single entity.

• Werner defined secondary valence as the number of groups bound directly to the metal ion (coordination number).

• Isomers: Compounds with identical empirical formulas but distinct properties (e.g., the two $CoCl<em>3 "." 4NH</em>3$ compounds).

Werner's Theory Postulates (1898)

1. Metals in coordination compounds show primary and secondary valences.

2. Primary valences are ionisable and satisfied by negative ions.

3. Secondary valences are non-ionisable, satisfied by neutral molecules or negative ions, and equal to the coordination number.

4. Ions/groups bound by secondary linkages have characteristic spatial arrangements (coordination polyhedra).

• Coordination entities or complexes are species within square brackets, and ions outside are counter ions.

• Octahedral, tetrahedral, and square planar shapes are common geometries.

• Examples: $[Co(NH<em>3)</em>6]^{3+}$, $[CoCl(NH<em>3)</em>5]^{2+}$, and $[CoCl<em>2(NH</em>3)_4]^+$ are octahedral.

• $[Ni(CO)<em>4]$ and $[PtCl</em>4]^{2-}$ are tetrahedral and square planar, respectively.

Example 5.1 Secondary Valences

Secondary valences based on observations with aqueous solutions:

1. $PdCl<em>2 "." 4NH</em>3$: Secondary valence 4 (2 moles of $AgCl$ precipitated per mole of compound)

2. $NiCl<em>2 "." 6H</em>2O$: Secondary valence 6 (2 moles of $AgCl$ precipitated per mole of compound)

3. $PtCl_4 "." 2HCl$: Secondary valence 6 (0 moles of $AgCl$ precipitated per mole of compound)

4. $CoCl<em>3 "." 4NH</em>3$: Secondary valence 6 (1 mole of $AgCl$ precipitated per mole of compound)

5. $PtCl<em>2 "." 2NH</em>3$: Secondary valence 4 (0 moles of $AgCl$ precipitated per mole of compound)

Difference Between Double Salt and Complex

• Both are formed by combining two or more stable compounds in stoichiometric ratios.

• Double salts dissociate into simple ions when dissolved in water (e.g., carnallite, Mohr’s salt, potash alum).

• Complex ions do not dissociate into simple ions (e.g., $[Fe(CN)<em>6]^{4-}$ of $K</em>4[Fe(CN)_6]$).

Alfred Werner Biography

• Born on December 12, 1866, in Mülhouse, Alsace.

• Studied chemistry in Karlsruhe (Germany) and Zurich (Switzerland).

• Doctoral thesis in 1890 explained isomerism in nitrogen-containing organic substances.

• Extended van't Hoff’s theory for nitrogen.

• Showed optical and electrical differences in complex compounds.

• First to discover optical activity in coordination compounds.

• Became a full professor at Technische Hochschule in Zurich in 1895 at age 29.

• Developed the theory of coordination compounds (1890-1893).

• Won the Nobel Prize in 1913 for work on linkage of atoms and coordination theory.

Important Terms

(a) Coordination Entity

• A central metal atom or ion bonded to a fixed number of ions or molecules.

• Examples: $[CoCl<em>3(NH</em>3)<em>3]$, $[Ni(CO)</em>4]$, $[PtCl<em>2(NH</em>3)<em>2]$, $[Fe(CN)</em>6]^{4-}$, $[Co(NH<em>3)</em>6]^{3+}$.

(b) Central Atom/Ion

• The atom/ion to which a fixed number of ions/groups are bound in a definite geometrical arrangement.

• Examples: $Ni^{2+}$ in $[NiCl<em>2(H</em>2O)<em>4]$, $Co^{3+}$ in $[CoCl(NH</em>3)<em>5]^{2+}$, $Fe^{3+}$ in $[Fe(CN)</em>6]^{3-}$.

• These are Lewis acids.

(c) Ligands

• Ions or molecules bound to the central atom/ion.

• Can be simple ions (e.g., $Cl^-$), small molecules (e.g., $H<em>2O$, $NH</em>3$), larger molecules (e.g., $H<em>2NCH</em>2CH<em>2NH</em>2$), or macromolecules (e.g., proteins).

• Unidentate: Ligand bound through a single donor atom (e.g., $Cl^-$, $H<em>2O$, $NH</em>3$).

• Didentate: Ligand bound through two donor atoms (e.g., $H<em>2NCH</em>2CH<em>2NH</em>2$ (ethane-1,2-diamine), $C<em>2O</em>4^{2-}$ (oxalate)).

• Polydentate: Ligand with several donor atoms (e.g., $N(CH<em>2CH</em>2NH<em>2)</em>3$).

• $EDTA^{4-}$: Important hexadentate ligand (two nitrogen and four oxygen atoms).

• Chelate ligand: A di- or polydentate ligand that uses two or more donor atoms simultaneously to bind a single metal ion.

• Chelate complexes are more stable than those with unidentate ligands.

• Ambidentate ligand: Ligand with two different donor atoms that can ligate (e.g., $NO_2^-$, $SCN^-$).

• $NO_2^-$ can coordinate through nitrogen or oxygen.

• $SCN^-$ can coordinate through sulfur or nitrogen.

(d) Coordination Number (CN)

• Number of ligand donor atoms to which the metal is directly bonded.

• Examples: $[PtCl<em>6]^{2-}$ (CN of Pt = 6), $[Ni(NH</em>3)_4]^{2+}$ (CN of Ni = 4).

• $[Fe(C<em>2O</em>4)<em>3]^{3-}$ and $[Co(en)</em>3]^{3+}$ (CN of Fe and Co = 6, because $C<em>2O</em>4^{2-}$ and en are didentate ligands).

(e) Coordination Sphere

• The central atom/ion and the ligands attached to it, enclosed in square brackets.

• Ionisable groups written outside the bracket are counter ions.

• Example: $K<em>4[Fe(CN)</em>6]$; coordination sphere is $[Fe(CN)_6]^{4-}$, counter ion is $K^+$.

(f) Coordination Polyhedron

• Spatial arrangement of ligand atoms directly attached to the central atom/ion.

• Common polyhedra: octahedral, square planar, and tetrahedral.

• Examples: $[Co(NH<em>3)</em>6]^{3+}$ (octahedral), $[Ni(CO)<em>4]$ (tetrahedral), $[PtCl</em>4]^{2-}$ (square planar).

(g) Oxidation Number of Central Atom

• Charge the central atom would carry if all ligands are removed with shared electron pairs.

• Represented by Roman numeral in parenthesis after the name of the coordination entity.

• Example: Copper in $[Cu(CN)_4]^{3-}$ is +1, written as Cu(I).

(h) Homoleptic and Heteroleptic Complexes

• Homoleptic: Metal bound to only one kind of donor group (e.g., $[Co(NH<em>3)</em>6]^{3+}$).

• Heteroleptic: Metal bound to more than one kind of donor group (e.g., $[Co(NH<em>3)</em>4Cl_2]^+$).

Nomenclature

• Important for describing formulas and systematic names unambiguously, especially with isomers.

• Based on IUPAC recommendations.

5.3.1 Formulas of Mononuclear Coordination Entities

• Shorthand tool to provide basic information. Rules:

  1. Central atom listed first.

  2. Ligands listed alphabetically.

  3. Polydentate ligands also listed alphabetically.

  4. Formula for the entire coordination entity enclosed in square brackets.

  5. No space between ligands and metal within a coordination sphere.

  6. Charge indicated outside the square brackets as a right superscript.

  7. Charge of cation(s) balanced by the charge of the anion(s).

5.3.2 Naming of Mononuclear Coordination Compounds

Rules:

1. Cation named first.

2. Ligands named alphabetically before the central atom/ion (reversed from writing formula).

3. Anionic ligands end in –o; neutral and cationic ligands are the same (aqua for $H<em>2O$, ammine for $NH</em>3$, carbonyl for CO, nitrosyl for NO).

4. Prefixes mono, di, tri, etc., indicate the number of individual ligands.

5. bis, tris, tetrakis are used when ligand names include a numerical prefix; the ligand is placed in parentheses.

   • Example: $[NiCl<em>2(PPh</em>3)_2]$ is dichloridobis(triphenylphosphine)nickel(II).

6. Oxidation state of the metal indicated by Roman numeral in parenthesis.

7. If the complex ion is a cation, the metal is named as the element (e.g., cobalt, platinum).

8. If the complex ion is an anion, the name of the metal ends with the suffix –ate (e.g., cobaltate, ferrate).

Examples

1. $[Cr(NH<em>3)</em>3(H<em>2O)</em>3]Cl_3$: triamminetriaquachromium(III) chloride

   • Complex ion inside the square bracket is a cation.

   • Amine ligands named before aqua ligands alphabetically.

   • Oxidation number of chromium is +3.

2. $[Co(H<em>2NCH</em>2CH<em>2NH</em>2)<em>3]</em>2(SO<em>4)</em>3$: tris(ethane-1,2–diamine)cobalt(III) sulphate

   • Sulphate is the counter anion.

   • Oxidation number of cobalt is +3.

3. $[Ag(NH<em>3)</em>2][Ag(CN)_2]$: diamminesilver(I)dicyanidoargentate(I)

Examples 5.2 and 5.3

Example 5.2: Write the formulas for the following coordination compounds:

(a) tetraammineaquachloridocobalt(III) chloride

(b) potassium tetrahydroxidozincate(II)

(c) potassium trioxalatoaluminate(III)

(d) dichloridobis(ethane-1,2-diamine)cobalt(III)

(e) tetracarbonylnickel(0)

Solutions:

(a) $[Co(NH<em>3)</em>4(H<em>2O)Cl]Cl</em>2$

(b) $K<em>2[Zn(OH)</em>4]$

(c) $K<em>3[Al(C</em>2O<em>4)</em>3]$

(d) $[CoCl<em>2(en)</em>2]^+$

(e) $[Ni(CO)_4]$

Example 5.3: Write the IUPAC names of the following coordination compounds:

(a) $[Pt(NH<em>3)</em>2Cl(NO_2)]$

(b) $K<em>3[Cr(C</em>2O<em>4)</em>3]$

(c) $[CoCl<em>2(en)</em>2]Cl$

(d) $[Co(NH<em>3)</em>5(CO_3)]Cl$

(e) $Hg[Co(SCN)_4]$

Solutions:

(a) diamminechloridonitrito-N-platinum(II)

(b) potassium trioxalatochromate(III)

(c) dichloridobis(ethane-1,2-diamine)cobalt(III) chloride

(d) pentaamminecarbonatocobalt(III) chloride

(e) mercury (I) tetrathiocyanato-S-cobaltate(III)

Isomerism

• Isomers: Compounds with the same chemical formula but different arrangements of atoms.

• Differ in physical or chemical properties.

• Two principal types:

  • Stereoisomerism

    1. Geometrical isomerism

    2. Optical isomerism

  • Structural isomerism

    1. Linkage isomerism

    2. Coordination isomerism

    3. Ionisation isomerism

    4. Solvate isomerism

• Stereoisomers: Same chemical formula and bonds but different spatial arrangements.

• Structural isomers: Different bonds.

5.4.1 Geometric Isomerism

• Arises in heteroleptic complexes due to different geometric arrangements of ligands.

• Important examples with coordination numbers 4 and 6.

• Square planar complex $[MX<em>2L</em>2]$ (X and L are unidentate):

  • cis isomer: Two ligands X are adjacent.

  • trans isomer: Two ligands X are opposite.

• Square planar complex MABXL (A, B, X, L are unidentates): three isomers (two cis, one trans).

• Not possible for tetrahedral geometry.

• Octahedral complexes of formula $[MX<em>2L</em>4]$: ligands X can be cis or trans to each other.

• Didentate ligands L – L [e.g., $NH<em>2CH</em>2CH<em>2NH</em>2$ (en)] in complexes of formula $[MX<em>2(L–L)</em>2]$.

• Octahedral coordination entities of the type $[Ma<em>3b</em>3]$ like $[Co(NH<em>3)</em>3(NO<em>2)</em>3]$.

  • facial (fac) isomer: Three donor atoms of the same ligands occupy adjacent positions.

  • meridional (mer) isomer: Positions are around the meridian of the octahedron.

Example 5.4

• Question: Why is geometrical isomerism not possible in tetrahedral complexes having two different types of unidentate ligands coordinated with the central metal ion?

• Answer: Tetrahedral complexes do not show geometrical isomerism because the relative positions of the unidentate ligands attached to the central metal atom are the same with respect to each other.

5.4.2 Optical Isomerism

• Optical isomers are mirror images that cannot be superimposed (enantiomers).

• Molecules or ions that cannot be superimposed are called chiral.

• Two forms: dextro (d) and laevo (l), depending on the direction they rotate plane-polarized light.

• Common in octahedral complexes involving didentate ligands.

• Coordination entity of the type $[PtCl<em>2(en)</em>2]^{2+}$, only the cis-isomer shows optical activity.

5.4.3 Linkage Isomerism

• Arises in coordination compounds containing ambidentate ligands.

• Example: Complexes containing the thiocyanate ligand, $NCS^-$, which may bind through nitrogen (M–NCS) or sulfur (M–SCN).

• Discovered by Jørgensen in $[Co(NH<em>3)</em>5(NO<em>2)]Cl</em>2$:

  • red form: nitrite ligand bound through oxygen (–ONO)

  • yellow form: nitrite ligand bound through nitrogen (–NO2)

5.4.4 Coordination Isomerism

• Arises from the interchange of ligands between cationic and anionic entities of different metal ions.

• Example: $[Co(NH<em>3)</em>6][Cr(CN)<em>6]$; coordination isomer is $[Cr(NH</em>3)<em>6][Co(CN)</em>6]$.

5.4.5 Ionisation Isomerism

• Arises when the counter ion in a complex salt is itself a potential ligand and can displace a ligand.

• Example: ionisation isomers $[Co(NH<em>3)</em>5(SO<em>4)]Br$ and $[Co(NH</em>3)<em>5Br]SO</em>4$.

5.4.6 Solvate Isomerism

• Hydrate isomerism when water is involved as a solvent.

• Solvate isomers differ by whether a solvent molecule is directly bonded to the metal ion or merely present as free solvent molecules in the crystal lattice.

• Example: aqua complex $[Cr(H<em>2O)</em>6]Cl<em>3$ (violet) and its solvate isomer $[Cr(H</em>2O)<em>5Cl]Cl</em>2 "." H_2O$ (grey-green).

• Werner described bonding features, his theory couldn't answer:

  1. Why certain elements form coordination compounds?

  2. Why bonds have directional properties?

  3. Why compounds have magnetic and optical properties?

Bonding in Coordination Compounds

• Approaches to explain bonding: Valence Bond Theory (VBT), Crystal Field Theory (CFT), Ligand Field Theory (LFT), and Molecular Orbital Theory (MOT).

• Focus on VBT and CFT.

5.5.1 Valence Bond Theory

• Metal atom/ion under ligand influence uses $(n-1)d$, $ns$, $np$ orbitals for hybridisation.

• Yields equivalent orbitals of definite geometry (octahedral, tetrahedral, square planar).

• Hybridised orbitals overlap with ligand orbitals that can donate electron pairs for bonding.

Diamagnetic Octahedral Complex, $\left[Co(NH<em>3)</em>6\right]^{3+}$

• Cobalt ion is in +3 oxidation state and has the electronic configuration $3d^6$.

• Hybridisation scheme: use inner d orbital (3d) in hybridisation, the complex called inner orbital or low spin or spin paired complex.

• Six pairs of electrons, one from each $NH_3$ molecule, occupy the six hybrid orbitals, octahedral geometry. Diamagnetic because of the absence of unpaired electron.

The paramagnetic octahedral complex$\left[CoF_6\right]^{3-}$

• Uses outer orbital (4d ) in hybridisation (sp 3d2). It is thus called outer orbital or high spin or spin free complex

The tetrahedral complexes

$[NiCl_4]^{2-}$

• One s and three p orbitals are hybridised to form four equivalent orbitals oriented tetrahedrally. compound is paramagnetic since it contains two unpaired electrons.

*Coordination Compounds

5.5.2 Magnetic Properties of Coordination Compounds

• Magnetic moment measured by magnetic susceptibility experiments.

• Results used to obtain information about the number of unpaired electrons and structures adopted by metal complexes.

*Critical study of the magnetic data of coordination compounds of metals of the first transition series reveals some complications

Tetrahedral complexes

$\left[Ni(CO)_4\right]$

• Has tetrahedral geometry but is diamagnetic since nickel is in zero oxidation state and contains no unpaired electron

Square Planar Complexes

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

• The hybridisation involved is dsp2

• Compound is diamagnetic as evident from the absence of unpaired electron.

5.5.3 Limitations of Valence Bond Theory

1. Involves a number of assumptions.

2. Does not give quantitative interpretation of magnetic data.

3. Does not explain the color exhibited by coordination compounds.

4. Does not give a quantitative interpretation of the thermodynamic or kinetic stabilities of coordination compounds.

5. Does not make exact predictions regarding the tetrahedral and square planar structures of 4-coordinate complexes.

6. Does not distinguish between weak and strong ligands.

5.5.4 Crystal Field Theory

• Electrostatic model that considers the metal-ligand bond to be ionic.

• Arises from electrostatic interactions between the metal ion and the ligand.

• Ligands are treated as point charges (anions) or point dipoles (neutral molecules).

• Five d orbitals in an isolated gaseous metal atom/ion have the same energy (degenerate).

• Degeneracy is maintained if a spherically symmetrical field of negative charges surrounds the metal atom/ion.

• Asymmetrical field due to ligands lifts the degeneracy of the d orbitals, resulting in splitting.

• Splitting pattern depends on the nature of the crystal field.

• In an octahedral coordination entity with six ligands surrounding the metal atom/ion, there will be repulsion between the electrons in metal d orbitals and the electrons (or negative charges) of the ligands.

• such a repulsion is more when the metal d orbital is directed towards the ligand than when it is away from the ligand.

*Thus, the $d<em>{x^2-y^2}$ and $d</em>z^2$ orbitals which point towards the axes along the direction of the ligand will experience more repulsion and will be raised in energy

• the $d<em>{xy}$, $d</em>{yz}$ and $d_{xz}$ orbitals which are directed between the axes will be lowered in energy relative to the average energy in the spherical crystal field.

Crystal Field Splitting

• Degeneracy of the d orbitals is removed due to ligand electron-metal electron repulsions.

• Yields three orbitals of lower energy, $t<em>{2g}$ set, and two orbitals of higher energy, $e</em>g$ set.

• The energy separation is denoted by $\Delta_o$ (subscript o for octahedral).

• Energy of the two $e<em>g$ orbitals increases by $(3/5) \Delta</em>o$, and that of the three $t<em>{2g}$ decreases by $(2/5)\Delta</em>o$.

• Crystal field splitting, $\Delta_o$, depends on the field produced by the ligand and charge on the metal ion.

Spectrochemical Series

• Ligands arranged in order of increasing field strength:

• I^- < Br^- < SCN^- < Cl^- < S^{2-} < F^- < OH^- < C2O4^{2-} < H2O < NCS^- < edta^{4-} < NH3 < en < CN^- < CO

• Experimentally determined series based on absorption of light.
  Crystal Field Splitting and Electron Distribution

1. d1: Single d electron occupies one of the lower energy $t_{2g}$ orbitals.

2. d2 and d3: d electrons occupy the $t_{2g}$ orbitals singly in accordance with Hund’s rule.

3. d4: Two possible patterns of electron distribution arise:

   1. Fourth electron enters the $t_{2g}$ level and pairs with an existing electron.

   2. It avoids paying the price of the pairing energy by occupying the $e_g$ level.

Crystal Field Splitting, Δo and the Pairing Energy, P.

*Which of these possibilities occurs, depends on the relative magnitude of the crystal field splitting, Δo and the pairing energy, P

1. If \Deltao < P the fourth electron enters one of the eg$ orbitals giving the configuration t{2g}^3 eg^1$. Ligands for which$\Delta_o < P$are known as weak field ligands and form high spin complexes.</p></li><li><p>If$\Deltao > P$it becomes more energetically favourable for the fourth electron to occupy a$t{2g}$orbital with configuration$t{2g}^4 eg^0$. Ligands which produce this effect are known as strong field ligands and form low spin complexes.</p></li></ol><p>*(Calculations show that d4 to d7 coordination entities are more stable for strong field as compared to weak field cases).</p><h5 id="0d30b378-2237-4af1-8c11-733721fe7dfb" data-toc-id="0d30b378-2237-4af1-8c11-733721fe7dfb" collapsed="false" seolevelmigrated="true">5.5.5 Colour in Coordination Compounds</h5><table style="min-width: 100px"><colgroup><col style="min-width: 25px"><col style="min-width: 25px"><col style="min-width: 25px"><col style="min-width: 25px"></colgroup><tbody><tr><th colspan="1" rowspan="1"><p>Coordination entity</p></th><th colspan="1" rowspan="1"><p>Wavelength of light absorbed (nm)</p></th><th colspan="1" rowspan="1"><p>Colour of light absorbed</p></th><th colspan="1" rowspan="1"><p>Colour of coordination entity</p></th></tr><tr><td colspan="1" rowspan="1"><p>[CoCl(NH3)5]2+</p></td><td colspan="1" rowspan="1"><p>535</p></td><td colspan="1" rowspan="1"><p>Yellow</p></td><td colspan="1" rowspan="1"><p>Violet</p></td></tr><tr><td colspan="1" rowspan="1"><p>[Co(NH3)5(H2O)]3+</p></td><td colspan="1" rowspan="1"><p>500</p></td><td colspan="1" rowspan="1"><p>Blue Green</p></td><td colspan="1" rowspan="1"><p>Red</p></td></tr><tr><td colspan="1" rowspan="1"><p>[Co(NH3)6]3+</p></td><td colspan="1" rowspan="1"><p>475</p></td><td colspan="1" rowspan="1"><p>Blue</p></td><td colspan="1" rowspan="1"><p>Yellow Orange</p></td></tr><tr><td colspan="1" rowspan="1"><p>[Co(CN)6]3–</p></td><td colspan="1" rowspan="1"><p>310</p></td><td colspan="1" rowspan="1"><p>Ultraviolet</p></td><td colspan="1" rowspan="1"><p>Pale Yellow</p></td></tr><tr><td colspan="1" rowspan="1"><p>[Cu(H2O)4]2+</p></td><td colspan="1" rowspan="1"><p>600</p></td><td colspan="1" rowspan="1"><p>Red</p></td><td colspan="1" rowspan="1"><p>Blue</p></td></tr><tr><td colspan="1" rowspan="1"><p>[Ti(H2O)6]3+</p></td><td colspan="1" rowspan="1"><p>498</p></td><td colspan="1" rowspan="1"><p>Blue Green</p></td><td colspan="1" rowspan="1"><p>Violet</p></td></tr></tbody></table><p>Crystal Field Theory Explanation</p><ul><li><p>Consider the complex$[Ti(H2O)6]^{3+}$, which is violet in color. This is an octahedral complex where the single electron ($Ti^{3+}$is a$3d^1$system) in the metal d orbital is in the$t_{2g}$level in the ground state of the complex.</p></li><li><p>The next higher state available for the electron is the empty$eg$level. If light corresponding to the energy of blue-green region is absorbed by the complex, it would excite the electron from$t{2g}$level to the$eg$level ($t{2g}^1 eg^0 \rightarrow t{2g}^0 e_g^1$). Consequently, the complex appears violet in colour.</p></li></ul><p>The Crystal field theory attributes the colour of the coordination compounds to d-d transition of the electron.</p><p>Crystal field splitting in tetrahedral coordination entities</p><p>*D orbital splitting is inverted and is smaller as compared to the octahedral field splitting<br>*For the same metal, the same ligands and metal-ligand distances, it can be shown that$\Deltat = (4/9) \Delta0$.</p><p>*Consequently, the orbital splitting energies are not sufficiently large for forcing pairing and, therefore, low spin configurations are rarely observed.</p><p>Examples:</p><p>*Anhydrous$CuSO4$is white, but$CuSO4.5H_2O$is blue in colour</p><p>Aqueous solutions of complexes of nickel(II) with an increasing number of ethane-1, 2-diamine ligands.</p><p>Nickel(II)Chloride is Dissolved in water.<br>[Ni(H2O)6]2+ (aq) + en (aq) = [Ni(H2O)4(en)]2+(aq) + 2H2O<br>green pale blue<br>[Ni(H2O)4 (en)]2+(aq) + en (aq) = [Ni(H2O)2(en)2]2+(aq) + 2H2O<br>blue/purple<br>[Ni(H2O)2(en)2]2+(aq) + en (aq) = [Ni(en)3]2+(aq) + 2H2O<br>violet<br>The colours produced by electronic transitions within the d orbitals of a transition metal ion occur frequently in everyday life<br>Examples<br>The absorption bands seen in the ruby shift to longer wavelength, namely yellow-red and blue, causing emerald to transmit light in the green region</p><h5 id="de47d183-053b-4520-a042-c35210bc4039" data-toc-id="de47d183-053b-4520-a042-c35210bc4039" collapsed="false" seolevelmigrated="true">5.5.6 Limitations of Crystal Field Theory</h5><p>*It does not take into account the covalent character of bonding between the ligand and the central atoms<br>The metal–carbon bond in metal carbonyls possesses both σ and π character. The ligand to metal is σ bond and metal to ligand is π bond. This unique synergic bonding provides stability to metal carbonyls.</p><p>The chemistry of coordination compounds is an important and challenging area of modern inorganic chemistry</p><p>Coordination Compounds are of great importance. These compounds provide critical insights into the functioning and structures of vital components of biological systems</p><h5 id="d693d4c8-cdeb-4e15-b824-df60d94305c9" data-toc-id="d693d4c8-cdeb-4e15-b824-df60d94305c9" collapsed="false" seolevelmigrated="true">5.6 Bonding in Metal Carbonyls</h5><p>Metal-carbon bond in metal carbonyls possesses both σ and π character. The M–C σ bond is formed by the donation of lone pair of electrons on the carbonyl carbon into a vacant orbital of the metal. The M–C π bond is formed by the donation of a pair of electrons from a filled d orbital of metal into the vacant antibonding π* orbital of carbon monoxide. The metal to ligand bonding creates a synergic effect which strengthens the bond between CO and the metal.Coordination compounds find use in many qualitative and quantitative chemical analysis</p><p>Article can be electroplated with silver and gold much more smoothly and evenly from solutions of the complexes,$[Ag(CN)2]^−$and$[Au(CN)2]^−$than from a solution of simple metal ions.The coordination compounds are of great importance. These compounds are widely present in the mineral, plant and animal worlds and are known to play many important functions in the area of analytical chemistry, metallurgy, biological systems, industry and medicine.The first systematic attempt at explaining the formation, reactions, structure and bonding of a coordination compound was made by A. Werner. His theory postulated the use of two types of linkages (primary and secondary) by a metal atom/ion in a coordination compound. In the modern language of chemistry these linkages are recognised as the ionisable (ionic) and non-ionisable (covalent) bonds, respectively. Using the property of isomerism, Werner predicted the geometrical shapes of a large number of coordination entities.The Valence Bond Theory (VBT) explains with reasonable success, the formation, magnetic behaviour and geometrical shapes of coordination compounds. It, however, fails to provide a quantitative interpretation of magnetic behaviour and has nothing to say about the optical properties of these compounds</p><p>*For example, removal of water from$[Ti(H2O)6]Cl_3$on heating renders it colourless.</p><h4 id="5c267d82-70ef-4174-9ea2-33ce41ee7e81" data-toc-id="5c267d82-70ef-4174-9ea2-33ce41ee7e81" collapsed="false" seolevelmigrated="true">Other properties of coordination Compounds</h4><p>*Similarly, anhydrous$CuSO4$is white, but$CuSO4.5H2O$is blue in colour. *The influence of the ligand on the colour of a complex may be illustrated by considering the$[Ni(H2O)_6]^{2+}$complex, which forms when nickel(II) chloride is dissolved in water. <br>*If the didentate ligand, ethane-1,2-diamine(en) is progressively added in the molar ratios en:Ni, 1:1, 2:1, 3:1, the following series of reactions and their associated colour changes occur:</p><p>The Crystal Field Theory (CFT) to coordination compounds is based on the effect of different crystal fields (provided by the ligands taken as point charges), For example, impure nickel is converted to$[Ni(CO)_4]$$, which is decomposed to yield pure nickel.