Coordination Compounds: common doubts, answered
The questions students ask most often about Coordination Compounds, each with a short answer. For the full chapter, read the Coordination Compounds notes.
Werner's theory
Read this section in the notes →How did Werner use silver nitrate to work out the structure of cobalt ammine chlorides?
Silver nitrate precipitates only chloride ions that are free in solution, not those bonded directly to cobalt. CoCl₃·6NH₃ gave 3 moles of AgCl per mole, so all three chlorides were ionic: [Co(NH₃)₆]Cl₃. CoCl₃·5NH₃ gave 2 moles, so one chloride sat inside the sphere: [CoCl(NH₃)₅]Cl₂. Conductivity matched, showing the number of ions in each case.
What is the difference between primary and secondary valence in Werner's theory?
Primary valence is the metal's oxidation state; it is ionisable and satisfied by negative ions. Secondary valence is the coordination number; it is non-ionisable, fixed for a given metal, and satisfied by ligands, which may be neutral or negative. Secondary valences also point in definite directions, which is why complexes have particular shapes such as octahedral or square planar.
What is the difference between a double salt and a complex?
A double salt such as Mohr's salt, FeSO₄·(NH₄)₂SO₄·6H₂O, breaks up completely in water into its simple ions, so it gives every test for Fe²⁺, NH₄⁺ and SO₄²⁻. A complex keeps its coordination entity intact in solution. K₄[Fe(CN)₆] gives no test for Fe²⁺ or CN⁻, because those are held together in the [Fe(CN)₆]⁴⁻ ion.
Ligands and denticity
Read this section in the notes →What is an ambidentate ligand?
An ambidentate ligand has two different donor atoms but uses only one of them at a time to bind the metal. Nitrite can attach through nitrogen, as nitrito-N, or through oxygen, as nitrito-O. Thiocyanate can bind through sulphur or through nitrogen. This choice of donor atom is what gives rise to linkage isomerism.
What is a chelate and why are chelate complexes more stable?
A chelate forms when a ligand with two or more donor atoms grips the same metal ion, making a ring, as ethane-1,2-diamine and oxalate do. Chelates are more stable than comparable complexes of monodentate ligands, mainly because one chelating ligand replaces several separate ligands or water molecules, raising entropy; this extra stability is the chelate effect. EDTA⁴⁻, a hexadentate ligand, forms very stable chelates.
Coordination number and oxidation number
Read this section in the notes →Why is the coordination number of [Co(en)₃]³⁺ six and not three?
Coordination number counts the donor atoms bonded to the metal, not the number of ligands. Ethane-1,2-diamine is bidentate, binding through two nitrogen atoms, so three of them provide six donor atoms. The same rule gives a coordination number of 6 for [Fe(C₂O₄)₃]³⁻ and for [Co(EDTA)]⁻, where a single EDTA ligand supplies all six.
How do you find the oxidation state of the metal in a complex?
Set the sum of all charges equal to the overall charge of the coordination entity. Neutral ligands such as NH₃, H₂O and CO count as zero; anionic ligands carry their own charge, Cl⁻ as −1 and CN⁻ as −1. In [Fe(CN)₆]³⁻, x + 6(−1) = −3, so iron is +3. In [Ni(CO)₄] nickel is 0, because carbonyl is neutral.
Formulas and names
Read this section in the notes →When does the metal name end in -ate in a complex's name?
Only when the coordination entity is an anion. Then the metal takes the -ate ending, often from its Latin name: ferrate for iron, cuprate for copper, argentate for silver and plumbate for lead, while cobalt becomes cobaltate. So K₃[Fe(CN)₆] is potassium hexacyanidoferrate(III). In a cationic or neutral complex the ordinary name is kept, as in hexaamminecobalt(III) chloride.
In what order are ligands named in a coordination compound?
Ligands are listed in alphabetical order before the metal, ignoring numerical prefixes such as di, tri or tetra. For complicated ligands bis, tris and tetrakis are used, with the ligand name in brackets. So [CoCl₂(NH₃)₄]⁺ is tetraamminedichloridocobalt(III), since ammine is placed before chlorido. In the formula itself, the metal symbol is written first.
Geometrical and optical isomerism
Read this section in the notes →Why don't tetrahedral complexes show geometrical isomerism?
In a tetrahedron every corner is adjacent to every other, so no two positions can be described as across from or beside each other. Any arrangement of two different ligands therefore gives the same molecule. Geometrical isomers arise in square planar complexes such as [Pt(NH₃)₂Cl₂], which has cis and trans forms, and in octahedral ones such as [Co(NH₃)₄Cl₂]⁺.
What are fac and mer isomers?
They are the two geometrical isomers of an octahedral complex of type MA₃B₃, such as [Co(NH₃)₃(NO₂)₃]. In the fac isomer the three identical ligands occupy one triangular face of the octahedron. In the mer isomer they lie around a meridian, an arc passing through the metal, so two of them are trans to each other.
Why is cis-[PtCl₂(en)₂]²⁺ optically active but the trans form is not?
The cis isomer lacks any plane of symmetry, so it and its mirror image cannot be laid over each other and exist as a pair of enantiomers that rotate polarised light in opposite directions. In the trans isomer the two chlorides sit opposite each other, creating a plane of symmetry through the two en rings, so its mirror image is identical and it is optically inactive.
Structural isomerism
Read this section in the notes →What is the difference between ionisation isomers and hydrate isomers?
Ionisation isomers swap a ligand inside the coordination sphere with an ion outside, so they release different ions in water; [Co(NH₃)₅(SO₄)]Br and [Co(NH₃)₅Br]SO₄ are an example. Hydrate, or solvate, isomers differ in whether water is a ligand or lattice water: [Cr(H₂O)₆]Cl₃ is violet, while [CrCl(H₂O)₅]Cl₂·H₂O is grey-green.
Valence bond theory
Read this section in the notes →Why is [Ni(CN)₄]²⁻ square planar while [NiCl₄]²⁻ is tetrahedral?
Ni²⁺ is d⁸. Cyanide is a strong-field ligand that pairs the eight electrons into four d orbitals, leaving one 3d orbital empty for dsp² hybridisation, which gives a square planar, diamagnetic complex. Chloride is weak and cannot cause pairing, so the two unpaired electrons remain and nickel uses sp³ hybridisation, giving a tetrahedral, paramagnetic complex.
Magnetic properties and limits of VBT
Read this section in the notes →What is the difference between inner orbital and outer orbital complexes?
An inner orbital, or low-spin, complex uses inner (n−1)d orbitals in d²sp³ hybridisation, as [Co(NH₃)₆]³⁺ does, after strong-field ligands pair up the electrons. An outer orbital, or high-spin, complex uses outer nd orbitals in sp³d² hybridisation, as [CoF₆]³⁻ does, with no pairing. The outer orbital complex has more unpaired electrons and is more strongly paramagnetic.
What are the limitations of valence bond theory for complexes?
Valence bond theory explains shapes and magnetism through hybridisation, but it cannot explain why complexes are coloured. It gives no quantitative account of magnetic data, cannot reliably predict whether a four-coordinate complex will be tetrahedral or square planar, and makes no distinction between weak and strong ligands. These gaps are filled by crystal field theory.
Crystal field theory
Read this section in the notes →Why do the d orbitals split in an octahedral field?
In an octahedral complex six ligands approach along the x, y and z axes. The d(x²−y²) and d(z²) orbitals point straight at them and are repelled more, so they rise in energy as the eg set. The other three point between the axes and form the lower t₂g set. The energy gap between the sets is the crystal field splitting, Δo.
How does the strength of a ligand decide between high spin and low spin?
It depends on whether Δo is larger or smaller than the pairing energy P. Weak-field ligands such as halides give a small Δo, so the fourth electron enters an eg orbital, giving a high-spin complex. Strong-field ligands such as CN⁻ and CO give Δo greater than P, so electrons pair in t₂g first, giving a low-spin complex with fewer unpaired electrons.
Why are tetrahedral complexes almost always high spin?
The splitting in a tetrahedral field is only about four-ninths of that in an octahedral field, Δt = (4/9)Δo, because there are fewer ligands and none point directly at the orbitals. This splitting is almost always smaller than the pairing energy, so electrons spread into the upper set before they pair. The tetrahedral pattern is also inverted, with the e set below the t₂ set.
Colour in coordination compounds
Read this section in the notes →Why does [Ti(H₂O)₆]³⁺ look violet?
Ti³⁺ has one d electron, and absorbing light lifts it from t₂g to eg, a d-d transition. The complex absorbs most strongly near 498 nm, in the blue-green region. The eye sees the complementary colour of the light that remains, which is violet. The ligands set the size of the gap, so changing them changes the colour: white anhydrous CuSO₄ turns blue when water molecules bind to Cu²⁺.
Metal carbonyls
Read this section in the notes →Why is the metal-carbon bond in metal carbonyls so strong?
Bonding works in both directions, known as synergic bonding. Carbon monoxide donates its lone pair into an empty metal orbital, forming a σ bond. The metal then pushes electron density from a filled d orbital back into an empty antibonding π* orbital of CO. Each part strengthens the other, so the M–C bond is strong while the C–O bond is slightly weakened.
Importance and applications
Read this section in the notes →What are the important uses of coordination compounds?
EDTA estimates water hardness by binding Ca²⁺ and Mg²⁺, and it treats lead poisoning. Cis-platin inhibits tumour growth. Wilkinson's catalyst, [(Ph₃P)₃RhCl], hydrogenates alkenes. Silver and gold are extracted as cyanide complexes, and photographic fixing dissolves unused silver bromide as a thiosulphate complex. In living systems, chlorophyll contains magnesium, haemoglobin iron and vitamin B₁₂ cobalt.
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