8.1 Nomenclature and structure of coordination compounds

Coordination Compounds

Components of Coordination Compounds

A coordination compound consists of a central metal atom or ion bonded to surrounding molecules or ions called ligands. These bonds are coordinate covalent bonds, meaning the ligand donates both electrons in the shared pair to the metal center.

• Central metal atom: Typically a transition metal or its ion (Fe, Co, Ni, Cu, etc.). The metal acts as a Lewis acid, accepting electron pairs from ligands.
• Ligands: Molecules or ions that donate electron pairs to the metal (Lewis bases). They can be neutral ($NH_3$, $H_2O$) or charged ($Cl^-$, $CN^-$).
• Coordination number: The number of coordinate bonds formed between the metal and its ligands. Common coordination numbers are 2, 4, and 6.

Ligands are classified by how many donor atoms they use to bond to the metal:

• Monodentate: One donor atom, forming a single bond to the metal (e.g., $Cl^-$, $NH_3$)
• Bidentate: Two donor atoms, forming two bonds to the metal (e.g., ethylenediamine, abbreviated "en")
• Polydentate: Three or more donor atoms (e.g., EDTA, which bonds through six donor atoms)

The coordination number is determined by counting donor atoms, not ligands. For example, if three bidentate ligands bind to a metal, the coordination number is 6, not 3.

IUPAC Naming for Coordination Compounds

Naming coordination compounds follows a specific set of rules. Here's the process broken into steps:

1. Identify the cation and anion. The cation is always named first, regardless of which one is the complex ion.

2. Name the ligands in alphabetical order (alphabetical by ligand name, ignoring prefixes like di- or tri-).

   • Anionic ligands: change the -ide ending to -o ($Cl^-$ = chlorido, $CN^-$ = cyanido, $OH^-$ = hydroxido). Note: older naming uses "chloro," "cyano," etc., and many textbooks still use these forms.
   • Neutral ligands: generally use the molecule's name, with specific exceptions: $H_2O$ = aqua, $NH_3$ = ammine, CO = carbonyl, NO = nitrosyl.

3. Use Greek prefixes to indicate the number of each ligand: di-, tri-, tetra-, penta-, hexa-. If the ligand name already contains a Greek prefix (like ethylenediamine), use bis-, tris-, tetrakis- instead, and enclose the ligand name in parentheses.

4. Name the metal.

   • If the complex is a cation or neutral, use the normal metal name.
   • If the complex is an anion, add the suffix -ate to the metal name. Some metals use their Latin root: Fe → ferrate, Cu → cuprate, Ag → argentate.

5. Indicate the metal's oxidation state with a Roman numeral in parentheses immediately after the metal name.

Examples:

• $[Co(NH_3)_6]Cl_3$: The complex ion is $[Co(NH_3)_6]^{3+}$. Six ammine ligands, cobalt in +3 oxidation state. Name: hexaamminecobalt(III) chloride.
• $K_4[Fe(CN)_6]$: The complex ion is $[Fe(CN)_6]^{4-}$. Six cyanido ligands, iron in +2 oxidation state, and the complex is an anion. Name: potassium hexacyanidoferrate(II) (or potassium hexacyanoferrate(II) in older convention).
• $[CoCl_2(en)_2]Cl$: Two chlorido ligands and two ethylenediamine ligands, cobalt in +3 oxidation state. Name: dichlorobis(ethylenediamine)cobalt(III) chloride. (Note "bis" because ethylenediamine already contains "di-".)

Geometry of Coordination Compounds

The coordination number largely determines the shape of the complex. Ligand type and the metal's electron configuration also play a role.

Coordination number 2:

• Linear geometry (bond angle 180°)
• Relatively rare; most common with $d^{10}$ metal ions like $Ag^+$ and $Cu^+$
• Example: $[Ag(NH_3)_2]^+$

Coordination number 4:

• Two possible geometries:
  • Tetrahedral: Bond angles of approximately 109.5°. Common with metals that have a full or nearly full d-subshell, or with small, non-bulky ligands. Example: $[CoCl_4]^{2-}$
  • Square planar: Bond angles of 90°. Strongly favored by $d^8$ metal ions ($Pt^{2+}$, $Pd^{2+}$, $Ni^{2+}$ with strong-field ligands, $Au^{3+}$). Example: $[PtCl_4]^{2-}$
• The distinction between tetrahedral and square planar matters because it affects the compound's properties, including whether geometric isomers can exist (square planar complexes can have cis/trans isomers; tetrahedral ones generally cannot).

Coordination number 6:

• Octahedral geometry (bond angles of 90°). This is the most common geometry for coordination compounds.
• Example: $[Co(NH_3)_6]^{3+}$

Bidentate and polydentate ligands constrain the geometry because their donor atoms are connected by a molecular backbone. This limits the positions those donor atoms can occupy around the metal.

Chelation and Compound Stability

Chelation occurs when a single polydentate ligand bonds to a metal through two or more donor atoms, forming a ring that includes the metal. The term comes from the Greek word for "claw."

For example, ethylenediamine (en) bonds to a metal through both of its nitrogen atoms, forming a five-membered ring: M–N–C–C–N.

The chelate effect describes why complexes with chelating ligands are significantly more stable than comparable complexes with monodentate ligands. Consider this substitution reaction:

$[Ni(H_2O)_6]^{2+} + 3\,en \rightarrow [Ni(en)_3]^{2+} + 6\,H_2O$

Both complexes have six Ni–N/O bonds, so the enthalpy change is relatively small. But the reaction goes from 4 particles on the left to 7 on the right, which is a large increase in entropy ($\Delta S > 0$). This favorable entropy change drives the reaction forward, making the chelated complex more stable.

Factors that influence chelate stability:

1. Ring size: Five- and six-membered chelate rings are most stable because they have minimal angle strain. Rings with four or fewer members are too strained, while very large rings lose the entropic advantage.

2. Donor atom type: Ligands with nitrogen and oxygen donor atoms tend to form strong bonds with hard metal ions (like $Fe^{3+}$, $Co^{3+}$). Sulfur and phosphorus donors bond more strongly to soft metal ions (like $Pt^{2+}$, $Hg^{2+}$). This is related to Hard-Soft Acid-Base (HSAB) theory.

3. Metal-ligand size match: The chelate ring must accommodate the metal ion without excessive strain. A mismatch between the metal ion's radius and the natural "bite size" of the ligand reduces stability.