Introduction And Nomenclature

Coordination compounds are compounds in which a central metal atom or ion is bound to a fixed number of ions or molecules, called ligands, through coordinate bonds, as in $[\mathrm{Co}(\mathrm{N}{\mathrm{H}}_3{)}_6]\mathrm{C}{\mathrm{l}}_3$ and ${\mathrm{K}}_4[\mathrm{Fe}(\mathrm{CN}{)}_6]$. This page covers Werner's theory, every key term (ligand, denticity, chelate, coordination number, oxidation state) and the IUPAC names of coordination compounds, using the latest NCERT forms such as chlorido and cyanido. Naming and formula questions from this topic appear every year in JEE Main and NEET.

1. What Are Coordination Compounds?

Transition metals have a strong tendency to hold neutral molecules or anions around themselves. The compounds formed this way are called coordination compounds, and they are everywhere. Chlorophyll, the green pigment of plants, is a complex of Mg. Haemoglobin, the oxygen carrier in blood, is a complex of Fe. Vitamin ${\mathrm{B}}_{12}$ is a complex of Co.

One of the earliest known coordination compounds, Prussian blue, was made by accident in 1704 by Diesbach, a colour maker in Berlin. Today these compounds are important in metallurgy, analytical chemistry, medicine and industrial catalysis (for example, polymerisation).

Figure 1 shows all the parts at once. The metal ion and the six ligands inside the square bracket move together as one particle. The three chloride ions outside the bracket are ordinary ions that separate in water.

1.1 Molecular (Addition) Compounds

When solutions containing two or more stable compounds in simple molar ratios are evaporated, crystals of new substances form. These are called molecular or addition compounds:

Addition compounds fall into two groups: those that lose their identity in solution (double salts) and those that keep it (coordination compounds). The last three compounds above are really $[\mathrm{Cu}(\mathrm{N}{\mathrm{H}}_3{)}_4]\mathrm{S}{\mathrm{O}}_4\cdot {\mathrm{H}}_2\mathrm{O}$, ${\mathrm{K}}_4[\mathrm{Fe}(\mathrm{CN}{)}_6]$ and $\mathrm{K}[\mathrm{Ag}(\mathrm{CN}{)}_2]$.

1.2 Double Salts vs Coordination Compounds

Double salts (lattice compounds) are stable only in the solid state. In water they break up completely into their simple ions. An aqueous solution of carnallite shows the properties of ${\mathrm{K}}^{+}$, $\mathrm{M}{\mathrm{g}}^{2+}$ and $\mathrm{C}{\mathrm{l}}^{-}$ ions; potash alum shows those of ${\mathrm{K}}^{+}$, $\mathrm{A}{\mathrm{l}}^{3+}$ and $\mathrm{S}{\mathrm{O}}_4^{2-}$ ions. Mohr's salt behaves the same way:

Coordination compounds behave very differently. In water they give complex ions that stay whole: copper(II) sulphate with ammonia gives the deep blue $[\mathrm{Cu}(\mathrm{N}{\mathrm{H}}_3{)}_4({\mathrm{H}}_2\mathrm{O}{)}_2{]}^{2+}$ ion, and potassium ferrocyanide gives the $[\mathrm{Fe}(\mathrm{CN}{)}_6{]}^{4-}$ ion. The ferrocyanide solution gives no test for $\mathrm{F}{\mathrm{e}}^{2+}$ or $\mathrm{C}{\mathrm{N}}^{-}$ ions.

2. Werner's Coordination Theory

Alfred Werner, a Swiss chemist often called the father of coordination chemistry, proposed his theory in 1893 after studying compounds such as $\mathrm{CoC}{\mathrm{l}}_3\cdot 6N{\mathrm{H}}_3$ and $\mathrm{CuS}{\mathrm{O}}_4\cdot 4N{\mathrm{H}}_3$. He received the Nobel Prize in 1913. His main postulates are:

1. In a coordination compound the metal shows two kinds of valency: primary (ionisable) and secondary (non-ionisable).
2. Primary valency equals the oxidation state of the metal. It is satisfied only by anions, it can change from compound to compound and it is non-directional. Werner drew it as a dotted line.
3. Secondary valency equals the coordination number. It is fixed for a given metal ion and is satisfied by neutral molecules or anions that carry a lone pair. Werner drew it as a thick line.
4. Every metal ion tries to satisfy both valencies. An anion inside the coordination sphere can satisfy a primary and a secondary valency at the same time (like one $\mathrm{C}{\mathrm{l}}^{-}$ in $[\mathrm{CoCl}(\mathrm{N}{\mathrm{H}}_3{)}_5]\mathrm{C}{\mathrm{l}}_2$).
5. Ions attached by primary valency ionise in solution; groups attached by secondary valency do not.
6. Secondary valencies are directed in space: six point to the corners of an octahedron, four to the corners of a square or a tetrahedron. This fixed geometry is what makes stereoisomerism possible.

Werner tested his idea with silver nitrate, which precipitates only free chloride ions, and with conductivity, which counts the ions in solution. In $\mathrm{CoC}{\mathrm{l}}_3\cdot 6N{\mathrm{H}}_3$ all three chlorides precipitate at once, so all three are outside the coordination sphere and the compound is $[\mathrm{Co}(\mathrm{N}{\mathrm{H}}_3{)}_6]\mathrm{C}{\mathrm{l}}_3$:

In each compound, cobalt shows a primary valency of 3 and a secondary valency of 6. As ammonia molecules are removed, chloride ions move inside the bracket to keep the coordination number at 6, so fewer chlorides precipitate and fewer ions are present. Werner also noticed that $\mathrm{CoC}{\mathrm{l}}_3\cdot 4N{\mathrm{H}}_3$ exists in two forms (green and violet). He explained them as the trans and cis arrangements of the same octahedral ion, which strongly supported his octahedral model (see Isomerism).

3. Important Terms

3.1 Coordination Entity and Central Atom

The coordination entity is the central atom or ion together with the ligands bonded to it. It can be a cation ($[\mathrm{Co}(\mathrm{N}{\mathrm{H}}_3{)}_6{]}^{3+}$), an anion ($[\mathrm{PtC}{\mathrm{l}}_4{]}^{2-}$) or neutral ($[\mathrm{Ni}(\mathrm{CO}{)}_4]$, $[\mathrm{CoC}{\mathrm{l}}_3(\mathrm{N}{\mathrm{H}}_3{)}_3]$).

The central atom or ion is the atom to which a fixed number of ligands are attached in a definite geometry. It accepts electron pairs, so it acts as a Lewis acid and must have empty orbitals. This is why transition metals, with vacant d orbitals, form complexes so readily. In $[\mathrm{Ni}(\mathrm{N}{\mathrm{H}}_3{)}_6{]}^{2+}$ and $[\mathrm{Fe}(\mathrm{CN}{)}_6{]}^{3-}$ the central ions are $\mathrm{N}{\mathrm{i}}^{2+}$ and $\mathrm{F}{\mathrm{e}}^{3+}$; in $[\mathrm{Ni}(\mathrm{CO}{)}_4]$ it is a neutral Ni atom.

3.2 Ligands and Donor Atoms

A ligand is an ion or molecule that donates a pair of electrons to the central atom through a coordinate (dative) bond, so it acts as a Lewis base. Ligands are usually polar molecules such as $\mathrm{N}{\mathrm{H}}_3$ and ${\mathrm{H}}_2\mathrm{O}$, or anions such as $\mathrm{C}{\mathrm{l}}^{-}$, $\mathrm{O}{\mathrm{H}}^{-}$ and $\mathrm{C}{\mathrm{N}}^{-}$, each with at least one lone pair. The atom of the ligand that actually forms the bond is the donor atom (N in $\mathrm{N}{\mathrm{H}}_3$, O in ${\mathrm{H}}_2\mathrm{O}$, C in $\mathrm{C}{\mathrm{N}}^{-}$). A few cations can act as ligands if they still carry a lone pair, for example $\mathrm{N}{\mathrm{O}}^{+}$ and the hydrazinium ion $\mathrm{N}{\mathrm{H}}_2\mathrm{N}{\mathrm{H}}_3^{+}$.

3.3 Coordination Number

The coordination number (CN) is the number of donor atoms directly bonded to the central atom, which equals the number of coordinate bonds. Only $\sigma$ bonds are counted; any $\pi$ bonds between ligand and metal are ignored. A hexadentate ligand counts as six bonds, not one. The most common values are 2, 4 and 6; lighter transition metals usually show 4 or 6, while heavier metals can show 8 or more.

3.4 Coordination Sphere and Counter Ions

The central atom and the ligands directly attached to it form the coordination sphere, written inside square brackets. It behaves as a single unit, and the groups inside it do not ionise. The ions written outside the bracket form the ionisation sphere; they are called counter ions and they do ionise. In ${\mathrm{K}}_4[\mathrm{Fe}(\mathrm{CN}{)}_6]$ the coordination sphere is $[\mathrm{Fe}(\mathrm{CN}{)}_6{]}^{4-}$ and the counter ions are four ${\mathrm{K}}^{+}$ ions.

3.5 Coordination Polyhedron

The spatial arrangement of the donor atoms around the central atom is the coordination polyhedron. The common ones are octahedral ($[\mathrm{Co}(\mathrm{N}{\mathrm{H}}_3{)}_6{]}^{3+}$), square planar ($[\mathrm{Ni}(\mathrm{CN}{)}_4{]}^{2-}$), tetrahedral ($[\mathrm{Ni}(\mathrm{CO}{)}_4]$) and linear ($[\mathrm{Ag}(\mathrm{N}{\mathrm{H}}_3{)}_2{]}^{+}$). Why a complex takes a particular shape is explained in Bonding.

3.6 Oxidation Number of the Central Atom

The oxidation number is the charge the central atom would carry if all the ligands were removed together with the electron pairs they share with it. Neutral ligands (${\mathrm{H}}_2\mathrm{O}$, $\mathrm{N}{\mathrm{H}}_3$, $\mathrm{CO}$, en) count as 0. For $[\mathrm{Cr}({\mathrm{H}}_2\mathrm{O}{)}_4\mathrm{C}{\mathrm{l}}_2{]}^{+}$:

Similarly the oxidation number of Fe in $[\mathrm{Fe}(\mathrm{CN}{)}_6{]}^{4-}$ is +2, of Co in $[\mathrm{Co}(\mathrm{N}{\mathrm{H}}_3{)}_6{]}^{3+}$ is +3 and of Ni in $[\mathrm{Ni}(\mathrm{CO}{)}_4]$ is 0. The charge on a complex is the sum of the charges of its parts.

3.7 Homoleptic and Heteroleptic Complexes

A complex with only one kind of ligand is homoleptic ($[\mathrm{Co}(\mathrm{N}{\mathrm{H}}_3{)}_6{]}^{3+}$). A complex with more than one kind of ligand is heteroleptic ($[\mathrm{CoC}{\mathrm{l}}_2(\mathrm{N}{\mathrm{H}}_3{)}_4{]}^{+}$).

4. Types of Ligands

4.1 By Number of Donor Atoms (Denticity)

The denticity of a ligand is the number of donor atoms it uses to bind one metal ion. Ligands with one donor atom are unidentate; those with two or more are polydentate.

4.2 Chelating Ligands and the Chelate Effect

A polydentate ligand whose shape lets two or more of its donor atoms bind the same metal ion at once closes one or more rings. It is called a chelating ligand, the ring compound formed is a chelate, and the process is chelation. In $[\mathrm{Cu}(\mathrm{en}{)}_2{]}^{2+}$ each ethane-1,2-diamine closes a five-membered Cu-N-C-C-N ring.

Chelate complexes are much more stable than similar complexes of unidentate ligands. This extra stability is called the chelate effect. The main reason is entropy: when three en molecules replace six $\mathrm{N}{\mathrm{H}}_3$ molecules, the number of free particles rises from 4 to 7, so the change is strongly favoured:

• Ring size matters. Chelates with 5- or 6-membered rings are the most stable; smaller rings are strained.
• 5-membered rings are most stable with saturated ligands such as en.
• 6-membered rings are most stable with unsaturated ligands that have conjugated double bonds (such as acetylacetonate), because resonance involving the metal d orbitals and the ligand p orbitals adds stability.

4.3 Flexidentate Ligands

A polydentate ligand does not always use all its donor atoms. Such ligands are called flexidentate. EDTA is normally hexadentate but can act as a penta- or tetradentate ligand with some metal ions. Sulphate and carbonate can bind through one O atom (unidentate) or two (bidentate).

4.4 Ambidentate Ligands

An ambidentate ligand has two different donor atoms but binds the metal through only one of them at a time (Figure 7, bottom row). The donor atom is shown in the name:

Because the same ligand can bind in two ways, ambidentate ligands give linkage isomers (see Isomerism). Do not confuse them with bidentate ligands: en uses both its N atoms together, while $\mathrm{SC}{\mathrm{N}}^{-}$ uses either S or N.

5. Writing the Formula of a Complex

The formula is a short-hand description of the composition. IUPAC (and NCERT) use these rules:

1. Write the central atom first.
2. Then list the ligands in alphabetical order of their formulas or abbreviations; the charge of a ligand does not affect its position ($\mathrm{Cl}$ before ${\mathrm{H}}_2\mathrm{O}$ before $\mathrm{N}{\mathrm{H}}_3$).
3. Put polyatomic ligands and abbreviations in round brackets: ($\mathrm{N}{\mathrm{H}}_3$), (en), (CN).
4. Enclose the whole coordination entity in square brackets, with no spaces between the ligands and the metal.
5. Write the charge of a complex ion as a right superscript outside the bracket: $[\mathrm{Co}(\mathrm{CN}{)}_6{]}^{3-}$.
6. In a salt, write the cation first: ${\mathrm{K}}_3[\mathrm{Fe}(\mathrm{CN}{)}_6]$, $[\mathrm{Co}(\mathrm{N}{\mathrm{H}}_3{)}_6]\mathrm{C}{\mathrm{l}}_3$.

Example: tetraamminedichloridocobalt(III) ion is written $[\mathrm{CoC}{\mathrm{l}}_2(\mathrm{N}{\mathrm{H}}_3{)}_4{]}^{+}$. The order of ligands in a formula is sometimes relaxed in textbooks; it never changes the compound. In a name, however, the alphabetical order is strict.

6. IUPAC Nomenclature of Coordination Compounds

Many complexes still have trivial names (ferrocyanide, Zeise's salt, cisplatin), so IUPAC set up one system that works for all of them. The names below follow the current NCERT book, which uses the 2005 IUPAC rules (chlorido, cyanido). Older forms (chloro, cyano) are shown alongside because they still appear in exam options.

6.1 Rule 1: Cation Before Anion

Whether the complex is the cation or the anion, always name the cation first and then the anion, as in simple salts (sodium chloride).

• $[\mathrm{Co}(\mathrm{N}{\mathrm{H}}_3{)}_6]\mathrm{C}{\mathrm{l}}_3$: hexaamminecobalt(III) chloride (complex cation)
• ${\mathrm{K}}_4[\mathrm{Fe}(\mathrm{CN}{)}_6]$: potassium hexacyanidoferrate(II) (complex anion)
• $[\mathrm{Pt}(\mathrm{N}{\mathrm{H}}_3{)}_4][\mathrm{PtC}{\mathrm{l}}_4]$: tetraammineplatinum(II) tetrachloridoplatinate(II) (both complex)
• $[\mathrm{Ni}(\mathrm{CO}{)}_4]$: tetracarbonylnickel(0) (neutral complex, written as one word)

6.2 Rule 2: Ligands First, Then the Metal

Inside the coordination entity, name the ligands first, in alphabetical order, and then the central atom, all as one word. Ligand names follow three patterns.

(a) Anionic ligands end in -o: -ide becomes -ido, -ate becomes -ato and -ite becomes -ito.

(b) Neutral ligands keep the name of the molecule. The exceptions (in bold) are aqua, ammine, carbonyl and nitrosyl, plus thiocarbonyl for $\mathrm{CS}$.