JEE Chemistry Coordination Compounds Notes Set 03

Basic Terms of Coordination Chemistry

1. A coordination compound contains a central metal atom or ion bonded to a fixed number of ions or molecules called ligands. The metal accepts electron pairs and each ligand donates a pair, so every metal-ligand link is a coordinate (dative) bond.
2. The metal together with its ligands, written inside square brackets, is the coordination sphere, for example \( \ce{[Co(NH3)6]^{3+}} \). Species outside the brackets are counter ions and are the only ones that ionise freely in water.
3. Coordination number is the number of donor atoms attached to the metal. Common values: 6 for \( \ce{Fe^{2+}, Fe^{3+}, Co^{3+}, Cr^{3+}, Pt^{4+}} \); 4 for \( \ce{Pt^{2+}, Ni^{2+}, Cu^{2+}, Zn^{2+}} \); 2 for \( \ce{Ag+} \).
4. Oxidation state of the metal = charge on the complex minus the sum of the charges on the ligands. In \( \ce{K4[Fe(CN)6]} \) it is +2, in \( \ce{[Co(NH3)6]^{3+}} \) it is +3 and in \( \ce{Ni(CO)4} \) it is 0.
5. A double salt (Mohr's salt, alum, carnallite) breaks up completely into simple ions in water, while a complex keeps its identity. Stable complexes such as \( \ce{[Fe(CN)6]^{4-}} \) do not give the tests of Fe(II) or cyanide.

Types of Ligands

Polydentate ligands that close five- or six-membered rings are called chelating ligands. A chelate is more stable than the corresponding complex of unidentate ligands (the chelate effect). Ligands such as CO, \( \ce{CN-} \) and \( \ce{NO+} \) also accept electron density back from the metal into empty antibonding orbitals; they are pi-acceptor (pi-acid) ligands.

IUPAC Naming: Rules in Order

1. Name the cation first and then the anion, as for ordinary salts.
2. Inside the complex, name the ligands first in alphabetical order (numerical prefixes are ignored), and then the metal with its oxidation state in Roman numerals.
3. Neutral ligands keep their molecule names, except \( \ce{H2O} \) (aqua), \( \ce{NH3} \) (ammine), CO (carbonyl) and NO (nitrosyl). Anionic ligands ending in -ide change to -ido, -ate to -ato and -ite to -ito: chlorido, cyanido, hydroxido, oxalato, sulphato, nitro (N-bonded) and nitrito-O.
4. Use di, tri, tetra for simple ligands and bis, tris, tetrakis (with brackets) for ligands whose own names contain a prefix, such as ethane-1,2-diamine.
5. If the complex ion is a cation or neutral, the metal name is unchanged. If it is an anion, the metal name ends in -ate (ferrate, cuprate, platinate, argentate, aurate, with the Latin stem for some metals).
6. A ligand bridging two metals is shown with the Greek letter mu before its name.

Examples: \( \ce{K4[Fe(CN)6]} \) is potassium hexacyanidoferrate(II); \( \ce{[Co(NH3)6]Cl3} \) is hexaamminecobalt(III) chloride; \( \ce{[Cr(H2O)4Cl2]Cl} \) is tetraaquadichloridochromium(III) chloride; \( \ce{[Pt(NH3)2Cl4]} \) is diamminetetrachloridoplatinum(IV).

Werner's Theory and Its Evidence

1. A metal shows two kinds of valency. The primary valency equals the oxidation state, is satisfied by anions and is ionisable. The secondary valency equals the coordination number, is satisfied by anions or neutral molecules and is non-ionisable.
2. Secondary valencies are directed in space: six give an octahedral arrangement, four give a tetrahedral or square planar arrangement.
3. Evidence 1: silver nitrate precipitates only the chloride held by primary valency. \( \ce{CoCl3.6NH3} \) gives 3 AgCl, \( \ce{CoCl3.5NH3} \) gives 2, \( \ce{CoCl3.4NH3} \) gives 1.
4. Evidence 2: molar conductance follows the number of ions produced.

Effective Atomic Number Rule

E.A.N. = (atomic number of metal - oxidation state) + 2 x (number of ligands). Stable complexes often reach the atomic number of the next noble gas (36, 54 or 86). Examples: \( \ce{[Co(NH3)6]^{3+}} \) gives 24 + 12 = 36 and \( \ce{[Fe(CN)6]^{4-}} \) gives 24 + 12 = 36; \( \ce{[Cu(NH3)4]^{2+}} \) gives 27 + 8 = 35 and \( \ce{[Cr(H2O)6]^{3+}} \) gives 21 + 12 = 33, so the rule has exceptions. For neutral carbonyls the formula follows directly: \( \ce{Cr(CO)6} \), \( \ce{Fe(CO)5} \), \( \ce{Ni(CO)4} \), \( \ce{Ru(CO)5} \), and \( \ce{Co2(CO)8} \) with a Co-Co bond.

Valence Bond Theory

1. The metal provides empty orbitals that hybridise and accept the ligand lone pairs. Six ligands: \( d^2sp^3 \) (inner orbital, low spin, uses (n-1)d) or \( sp^3d^2 \) (outer orbital, high spin, uses nd). Four ligands: \( sp^3 \) (tetrahedral) or \( dsp^2 \) (square planar).
2. Magnetic moment (spin only) is \( \mu = \sqrt{n(n+2)} \) B.M., where n is the number of unpaired electrons: 1 gives 1.73, 2 gives 2.83, 3 gives 3.87, 4 gives 4.90 and 5 gives 5.92.
3. \( \ce{[Fe(CN)6]^{4-}} \): Fe(II) d6, strong ligand, \( d^2sp^3 \), diamagnetic. \( \ce{[FeF6]^{3-}} \): Fe(III) d5, weak ligand, \( sp^3d^2 \), five unpaired electrons. \( \ce{[Ni(CN)4]^{2-}} \): \( dsp^2 \), square planar, diamagnetic. \( \ce{[NiCl4]^{2-}} \): \( sp^3 \), tetrahedral, two unpaired electrons. \( \ce{Ni(CO)4} \): \( sp^3 \), diamagnetic.

Crystal Field Theory

1. Ligands are treated as point charges. In an octahedral field the d orbitals split into a lower \( t_{2g} \) set (three orbitals, energy -0.4 \( \Delta_o \)) and a higher \( e_g \) set (two orbitals, +0.6 \( \Delta_o \)). In a tetrahedral field the order reverses and \( \Delta_t = \dfrac{4}{9}\Delta_o \).
2. If \( \Delta_o \) is smaller than the pairing energy P the complex is high spin, otherwise low spin. For d4 to d7 octahedral ions both are possible; d1 to d3 and d8 to d10 have one arrangement only.
3. Spectrochemical series (weak to strong): \( \ce{I-} < \ce{Br-} < \ce{Cl-} < \ce{F-} < \ce{OH-} < \ce{C2O4^{2-}} < \ce{H2O} < \ce{NH3} < en < \ce{NO2-} < \ce{CN-} < CO \).
4. \( \Delta \) rises with the oxidation state of the metal and from 3d to 4d to 5d.
5. CFSE for octahedral complexes = [-0.4 x (electrons in \( t_{2g} \)) + 0.6 x (electrons in \( e_g \))] \( \Delta_o \) plus pairing energy terms. High spin d4 gives -0.6 \( \Delta_o \); low spin d6 gives -2.4 \( \Delta_o \) + 2P.
6. Colour arises from d-d transitions. The complex absorbs the colour whose energy equals \( \Delta \) and shows the complementary colour; larger \( \Delta \) means shorter wavelength absorbed. d0 and d10 ions (\( \ce{[Al(OH)4]^-} \), \( \ce{Zn^{2+}} \)) are colourless; permanganate and chromate are coloured by charge transfer.

Isomerism in Complexes

Counting rules for common octahedral types: Ma4b2 has 2 geometrical forms, Ma3b3 has 2 (fac, mer), M(AA)2b2 has 3 stereoisomers (trans plus a cis pair), M(AA)3 has 2 (a pair of enantiomers). Square planar Ma2b2 has 2 geometrical forms; tetrahedral complexes show no geometrical isomerism. A square planar complex of type Ma2b2 has no mirror-image problem because the molecule lies in a plane, so it is never optically active.

Organometallic Compounds and Metal Carbonyls

1. An organometallic compound has at least one direct metal-carbon bond. Examples: Zeise's salt \( \ce{K[Pt(\eta^2-C2H4)Cl3]} \), ferrocene \( \ce{Fe(\eta^5-C5H5)2} \), dibenzenechromium, Grignard reagents, trialkylaluminium (Ziegler-Natta catalyst with titanium).
2. Metal carbonyls bond by sigma donation from carbon to metal and pi back-donation from filled metal d orbitals into the empty antibonding pi orbital of CO. The back donation lengthens the C-O bond and lowers its stretching frequency.
3. Shapes: \( \ce{Ni(CO)4} \) tetrahedral; \( \ce{Fe(CO)5} \) trigonal bipyramidal; \( \ce{Cr(CO)6} \) octahedral; \( \ce{Fe2(CO)9} \) has three bridging CO groups; \( \ce{Mn2(CO)10} \) and \( \ce{Co2(CO)8} \) contain M-M bonds.

Stability and Applications

1. Stability rises with the charge on the metal ion, with a smaller metal ion, with stronger ligands and when chelate rings form.
2. EDTA titrations measure the hardness of water. Dimethylglyoxime gives a red precipitate with Ni(II). Excess ammonia with Cu(II) gives deep blue \( \ce{[Cu(NH3)4]^{2+}} \). Silver halides dissolve in hypo as \( \ce{[Ag(S2O3)2]^{3-}} \) (photographic fixing) and in ammonia or cyanide.
3. Chlorophyll is a magnesium complex, haemoglobin an iron complex and vitamin B12 a cobalt complex. Cis-platin is an anticancer drug.
4. Brown ring test: \( \ce{[Fe(H2O)5(NO)]^{2+}} \), with iron formally +1 (NO as nitrosonium). Nitroprusside gives violet colour with sulphide.

Worked Examples

1. Find the spin-only magnetic moment of \( \ce{[Mn(H2O)6]^{2+}} \). Mn(II) is d5, water is a weak ligand, so five unpaired electrons: \( \mu = \sqrt{5 \times 7} = 5.92 \) B.M.
2. How many ions and how many moles of AgCl does 1 mol of \( \ce{[Co(NH3)4Cl2]Cl} \) give? It gives 2 ions, \( \ce{[Co(NH3)4Cl2]+} \) and \( \ce{Cl-} \), and 1 mol of AgCl.
3. Predict the formula of the binary carbonyl of ruthenium. Ru is 44, noble gas needs 54, so 10 electrons from 5 CO: \( \ce{Ru(CO)5} \).