Structure and Nomenclature of Complexes

Coordination Compounds as Structured Particles

In this chapter, we focus on how coordination compounds (complexes) are built and how they are named. Their preparation, stability, bonding, and properties are dealt with in other chapters; here we only look at their architecture and systematic naming.

Basic Structural Elements of Complexes

A coordination compound is made up of a central atom/ion (usually a metal) and surrounding ligands that donate electron pairs to the metal.

Central Atom/Ion

• Typically a metal cation from the d-block or sometimes p-block or f-block.
• Written first in formulas of individual complex ions.
• The oxidation state of the central metal is not shown in the formula but is indicated in the name.

Examples of central ions:

• $ \ce{Fe^{2+}} $, $ \ce{Fe^{3+}} $
• $ \ce{Co^{2+}} $, $ \ce{Pt^{4+}} $

Ligands

Ligands are ions or neutral molecules that donate an electron pair to the central metal to form a coordinate bond (donor–acceptor bond).

Monodentate vs. Polydentate

• Monodentate ligands: bond to the metal through one donor atom.
• Typical examples:
• $ \ce{H2O} $ (aqua)
• $ \ce{NH3} $ (ammine)
• $ \ce{Cl^-} $ (chloro)
• $ \ce{CN^-} $ (cyano)
• Polydentate ligands: bind through two or more donor atoms (chelating ligands).
• Bidentate (2 donor atoms): e.g. ethane-1,2-diamine (en)
• Tridentate (3 donor atoms): e.g. diethylenetriamine (dien)
• Hexadentate (6 donor atoms): e.g. EDTA$^{4-}$

The number of donor atoms by which a ligand can bind is called its denticity.

Ambidentate Ligands

Some ligands can bind through different possible donor atoms, but usually only use one at a time.

Examples:

• $ \ce{NO2^-} $: binds via N (nitro) or O (nitrito).
• $ \ce{SCN^-} $: binds via S (thiocyanato-S or isothiocyanato) or N (thiocyanato-N).

Ambidentate character is important in structural formulas and names because it changes how the ligand is described.

Coordination Number and Coordination Sphere

Coordination Number

The coordination number of the central metal is the number of donor atoms directly bonded to it, not the number of ligands.

Examples:

• $ \ce{[Cu(NH3)4]^{2+}} $: 4 ammonia molecules, each monodentate → coordination number 4.
• $ \ce{[Co(en)3]^{3+}} $: 3 bidentate en ligands → $3 \times 2 = 6$ donor atoms → coordination number 6.
• $ \ce{[Cr(H2O)4Cl2]^+} $: 4 water + 2 chloride (all monodentate) → coordination number 6.

Typical coordination numbers for transition metals are 4 and 6, but other values occur as well.

Coordination Sphere

The coordination sphere includes the central metal plus all ligands directly bound to it. In formulas, it is written in square brackets.

• Example: $ \ce{[Co(NH3)6]Cl3} $
• Coordination sphere: $ \ce{[Co(NH3)6]^{3+}} $
• Counterions: $ \ce{3 Cl^-} $
• Species inside the brackets belong to the complex ion; species outside are counterions or uncoordinated molecules.

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Common Coordination Geometries

The spatial arrangement of ligands around the metal is the coordination geometry. For beginners, it is often enough to know the most common geometries for coordination numbers 4 and 6.

Coordination Number 4

Tetrahedral

• 4 ligands at the corners of a tetrahedron around the metal.
• Common for $ \ce{d^{10}} $ ions like $ \ce{Zn^{2+}} $, $ \ce{Cd^{2+}} $.

Example:

• $ \ce{[ZnCl4]^{2-}} $ – tetrahedral.

Square Planar

• 4 ligands at the corners of a square in one plane.
• Very common for $ \ce{d^8} $ metal ions like $ \ce{Pt^{2+}} $, $ \ce{Pd^{2+}} $.

Example:

• $ \ce{[PtCl4]^{2-}} $ – square planar.

For CN = 4 complexes, distinguishing between tetrahedral and square planar is often important for isomerism and reactivity (discussed elsewhere).

Coordination Number 6

Octahedral

• 6 ligands at the corners of an octahedron.
• The most common geometry for CN = 6.

Examples:

• $ \ce{[Fe(H2O)6]^{2+}} $
• $ \ce{[Co(NH3)6]^{3+}} $
• $ \ce{[Cr(H2O)4Cl2]^+} $

For octahedral complexes, the relative positions of ligands give rise to cis/trans and other isomers, which are closely linked to the structural description.

Structural Formulas and Representation

Formulas of complexes convey both composition and often aspects of structure.

Order of Writing Formulas

Within the coordination sphere:

1. Central metal symbol first.
2. Ligands are written after the metal.
3. Ligands are usually grouped:
• First: neutral ligands (often in alphabetical order by their names, not symbols).
• Then: anionic ligands.
4. The overall charge of the complex is indicated as a superscript outside the bracket.

Examples:

• $ \ce{[Co(NH3)6]Cl3} $
• $ \ce{K4[Fe(CN)6]} $
• $ \ce{[Cr(H2O)4Cl2]Cl} $

Showing Chelate Ligands

Polydentate ligands are usually written as a unit, often repeated with a numerical subscript:

• $ \ce{[Co(en)3]^{3+}} $ (en = ethane-1,2-diamine)
• $ \ce{[Ni(edta)]^{2-}} $ (edta = ethylenediaminetetraacetate)

In line structures, multiple bonds to the same ligand may be indicated by showing the ligand name once with an index, but chemically the important feature is how many donor atoms bind.

Systematic Nomenclature of Coordination Compounds

The aim is unambiguous naming: from the name, you should be able to write the formula, and vice versa. Here we follow commonly used IUPAC-style rules at an introductory level.

We proceed in three steps:

1. Identify and name the complex ion (coordination sphere).
2. Identify and name counterions if present.
3. Combine the parts into the full name.

1. Naming the Complex Part

1.1. Ligands First, then Metal

Within the complex (the species inside the brackets):

1. Name ligands (in alphabetical order).
2. Then name the metal with its oxidation state in Roman numerals in parentheses.

The overall pattern is:

$ \text{[ligand names] + [metal name] + (oxidation state)} $

Example:

• $ \ce{[Co(NH3)6]^{3+}} $ → hexaamminecobalt(III)

1.2. Names of Common Ligands

For systematic names, many familiar molecules have special ligand names:

• Neutral ligands:
• $ \ce{H2O} $ → aqua
• $ \ce{NH3} $ → ammine (note: ammine with double m)
• $ \ce{CO} $ → carbonyl
• $ \ce{NO} $ → nitrosyl
• Anionic ligands (ending usually in -o):
• $ \ce{Cl^-} $ → chloro
• $ \ce{Br^-} $ → bromo
• $ \ce{I^-} $ → iodo
• $ \ce{OH^-} $ → hydroxo
• $ \ce{CN^-} $ → cyano
• $ \ce{O^{2-}} $ → oxo
• $ \ce{SO4^{2-}} $ → sulfato
• $ \ce{NO3^-} $ → nitrato
• $ \ce{CO3^{2-}} $ → carbonato

Polydentate ligands:

• Ethane-1,2-diamine (often abbreviated en) → ethan-1,2-diamine (ethane-1,2-diamine).
• EDTA$^{4-}$ → ethylenediaminetetraacetato.

Important: In naming, ligands that are commonly written with abbreviations (en, EDTA) are usually retained as such after they have been defined, but the systematic name still exists.

1.3. Multiplicity Prefixes

When there is more than one of the same ligand, numerical prefixes are used:

• 2: di–
• 3: tri–
• 4: tetra–
• 5: penta–
• 6: hexa–

These prefixes are attached directly to the ligand names:

• $ \ce{[CoCl4]^{2-}} $ → tetrachlorocobaltate(II)
• $ \ce{[Cr(H2O)4Cl2]^+} $ → tetraaquadichlorochromium(III)

Special Case: Polydentate or Complex Ligand Names

If the ligand name itself already contains a numerical prefix (for example “ethane-1,2-diamine”), another set of prefixes is used to avoid confusion:

• 2: bis–
• 3: tris–
• 4: tetrakis–
• 5: pentakis–
• 6: hexakis–

These are written outside a bracketed ligand name, if necessary.

Examples:

• $ \ce{[Co(en)3]^{3+}} $ → tris(ethane-1,2-diamine)cobalt(III)
• $ \ce{[Ni(edta)]^{2-}} $ → ethylenediaminetetraacetatonickelate(II)
  (only one EDTA ligand → no need for bis/tris prefix)

1.4. Alphabetical Order

Ligands in the name are listed in alphabetical order of their ligand names, ignoring multiplicity prefixes.

Examples:

• $ \ce{[Cr(H2O)4Cl2]^+} $ → ligands: aqua (4), chloro (2)
• Alphabetical: aqua before chloro → tetraaquadichlorochromium(III).
• $ \ce{[CoCl2(NH3)4]^+} $ → ammine (4), chloro (2)
• Alphabetical: ammine before chloro → tetraammine­dichlorocobalt(III) (oxidation state must be computed).

2. Naming the Metal: Cationic vs Anionic Complexes

Whether the complex ion as a whole is positively or negatively charged affects the metal’s name.

2.1. Cationic and Neutral Complexes

If the complex ion is:

• Cationic (overall positive charge), or
• Neutral,

then the metal is named by its element name as in usual inorganic nomenclature.

Examples:

• $ \ce{[Co(NH3)6]^{3+}} $ → hexaamminecobalt(III)
• $ \ce{[PtCl2(NH3)2]} $ (neutral) → diamminedichloroplatinum(II)

2.2. Anionic Complexes

If the complex ion is anionic (overall negative charge), the metal name usually takes the suffix -ate. Sometimes the Latin name of the element is used.

Typical examples:

• Iron → ferrate
• Copper → cuprate
• Lead → plumbate
• Silver → argentate
• Gold → aurate
• Tin → stannate
• Cobalt → cobaltate (Latin: cobaltum, but modern practice often just uses English root + ate)
• Nickel → nickelate
• Chromium → chromate

Examples:

• $ \ce{[Fe(CN)6]^{4-}} $ → hexacyanidoferrate(II)
• $ \ce{[Fe(CN)6]^{3-}} $ → hexacyanidoferrate(III)
• $ \ce{[NiCl4]^{2-}} $ → tetrachloronickelate(II)

Note on spelling: Some IUPAC recommendations may use -ido endings (e.g. cyanido, chlorido), but beginners often first encounter the simpler -o forms (cyano, chloro), which we use consistently here.

Determining the Oxidation State of the Metal

To complete the name, the oxidation state of the central metal (in Roman numerals) must be determined from the formula.

Method:

1. Assign formal charges to each ligand (e.g. $ \ce{Cl^-} = -1 $, $ \ce{NH3} = 0 $, $ \ce{CN^-} = -1 $, $ \ce{H2O} = 0 $).
2. Let the oxidation state of the metal be $ x $.
3. Sum of oxidation state of metal + sum of ligand charges = overall charge of complex ion.

$$ x + \sum (\text{ligand charges}) = \text{charge of complex} $$

Solve for $ x $.

Example 1: $ \ce{[Co(NH3)6]Cl3} $

• Counterions: 3 $ \ce{Cl^-} $ → total $ -3 $.
• Therefore complex cation is $ +3 $: $ \ce{[Co(NH3)6]^{3+}} $.
• Ligands: $ \ce{NH3} $ all neutral.
• Equation: $ x + 6 \cdot 0 = +3 $ → $ x = +3 $.

Name: hexaamminecobalt(III) chloride.

Example 2: $ \ce{K4[Fe(CN)6]} $

• Counterions: 4 $ \ce{K^+} $ → total $ +4 $.
• Therefore complex anion is $ -4 $: $ \ce{[Fe(CN)6]^{4-}} $.
• Each $ \ce{CN^-} $ ligand has charge $ -1 $.
• Equation: $ x + 6 \cdot (-1) = -4 $ → $ x - 6 = -4 $ → $ x = +2 $.

Name of complex anion: hexacyanidoferrate(II); full compound: potassium hexacyanidoferrate(II).

Example 3: $ \ce{[Cr(H2O)4Cl2]Cl} $

• Outside the bracket: $ \ce{Cl^-} $ → complex cation is $ +1 $: $ \ce{[Cr(H2O)4Cl2]^+} $.
• Ligand charges: 4 $ \ce{H2O} $ (0), 2 $ \ce{Cl^-} $ ($ -1 $ each).
• Equation: $ x + 4 \cdot 0 + 2 \cdot (-1) = +1 $ → $ x - 2 = +1 $ → $ x = +3 $.

Name of complex cation: tetraaquadichlorochromium(III); full compound: tetraaquadichlorochromium(III) chloride.

Naming Entire Coordination Compounds

To name the full compound (possibly containing both complex ions and simple ions):

1. Name the cation first, then the anion, as in standard salt nomenclature.
2. Within each ionic species, use the coordination naming rules described above.

Examples:

1. $ \ce{[Co(NH3)6]Cl3} $
• Cation: $ \ce{[Co(NH3)6]^{3+}} $ → hexaamminecobalt(III)
• Anion: $ \ce{3 Cl^-} $ → chloride
• Full name: hexaamminecobalt(III) chloride.
2. $ \ce{K4[Fe(CN)6]} $
• Cation: $ \ce{4 K^+} $ → potassium
• Anion: $ \ce{[Fe(CN)6]^{4-}} $ → hexacyanidoferrate(II)
• Full name: potassium hexacyanidoferrate(II).
3. $ \ce{[Pt(NH3)2Cl4]} $
• Neutral complex (no counterions).
• Ligands: $ \ce{NH3} $ (ammine, neutral), $ \ce{Cl^-} $ (chloro).
• Oxidation state: $ x + 2\cdot 0 + 4\cdot(-1) = 0 \Rightarrow x = +4 $.
• Name: diamminetetrachloroplatinum(IV).

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Stereochemical Descriptors in Names (Basic Level)

Some complexes with the same formula can have different arrangements of ligands (isomers). At a structural level, the simplest stereochemical descriptors you need to recognize are:

Cis / Trans (for Square Planar and Octahedral)

• cis: like ligands adjacent (90° apart).
• trans: like ligands opposite each other (180° apart).

Examples (octahedral):

• $ \ce{[CoCl2(NH3)4]^+} $
• cis isomer: the two chlorido ligands adjacent → cis-tetraammine­dichlorocobalt(III) ion.
• trans isomer: the two chlorido ligands opposite → trans-tetraammine­dichlorocobalt(III) ion.

These prefixes are placed before the entire name of the complex ion:

• cis-tetraaquadichlorochromium(III)
• trans-diamminedichloroplatinum(II)

fac / mer (for Octahedral Complexes with 3 Identical Ligands)

When three identical ligands occupy positions on:

• One face of an octahedron (each adjacent to the others) → fac (facial).
• A meridian going through the metal (two trans to each other, one cis) → mer (meridional).

These become important when three identical monodentate ligands are present, e.g. $ \ce{[CoCl3(NH3)3]} $. Naming examples:

• fac-triamminetrichlorocobalt(III)
• mer-triamminetrichlorocobalt(III)

(Details of geometric and optical isomerism are treated more fully elsewhere; here they are introduced only as naming descriptors that reflect structural differences.)

From Name to Formula and Back: Summary of Procedure

From Formula to Name

1. Identify coordination sphere: look for square brackets.
2. Determine charge of the complex ion (from counterions).
3. Determine oxidation state of the metal.
4. List ligands:
• Determine each ligand’s name and charge.
• Count how many of each.
5. Arrange ligand names in alphabetical order (ignore di-, tri-, etc.).
6. Construct the complex ion name:
• [numerical prefix + ligand name] + [metal name (with -ate if anion)] + (oxidation state).
7. Name counterions as usual.
8. Combine parts: cation name first, then anion.

From Name to Formula

1. Identify cation and anion from word order.
2. In the complex ion name:
• Identify each ligand and the number of each from prefixes.
• Identify the metal and its oxidation state.
3. Assign charges to each ligand; use the metal oxidation state to determine the complex ion charge.
4. Combine complex ion(s) and counterion(s) so that the total charge is zero for neutral compounds.
5. Write complex ion in square brackets, with the metal first and ligands after it.

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This chapter has focused on how complexes are structured (metal, ligands, coordination sphere, geometry) and how these structural features are encoded in systematic names. How these structures arise during synthesis, how they affect stability, and how we interpret the bonding will be discussed in later chapters on synthesis, stability, and bonding in complexes.