Common Ligands: Monodentate, Bidentate and Chelating

A metal ion on its own tells you only part of the story. Put Ni²⁺ in water and it is green; give it ammonia and it turns violet-blue; offer it dimethylglyoxime and a bright red solid drops out. The metal is the same each time. The difference is the ligand: the molecule or ion that donates a lone pair to the metal through a dative bond (see dative covalent bonds).

This list groups the ligands you are most likely to meet in A-level, IB and first-year university chemistry by denticity, the number of donor atoms one ligand uses to grip a single metal. For each, you get its donor atom, charge, usual abbreviation and the reason chemists care about it.

Quick reference table

Ligand	Formula / abbreviation	Charge	Donor atom(s)	Denticity
Water (aqua)	H₂O	0	O	1
Ammonia (ammine)	NH₃	0	N	1
Chloride (chlorido)	Cl⁻	−1	Cl	1
Hydroxide (hydroxido)	OH⁻	−1	O	1
Cyanide (cyanido)	CN⁻	−1	C	1
Carbon monoxide (carbonyl)	CO	0	C	1
Thiocyanate	SCN⁻	−1	S or N	1 (ambidentate)
Nitrite	NO₂⁻	−1	N or O	1 (ambidentate)
Pyridine	py, C₅H₅N	0	N	1
Triphenylphosphine	PPh₃	0	P	1
Ethane-1,2-diamine	en	0	N, N	2
Oxalate (ethanedioate)	ox, C₂O₄²⁻	−2	O, O	2
Acetylacetonate	acac⁻	−1	O, O	2
2,2′-Bipyridine	bipy (or bpy)	0	N, N	2
1,10-Phenanthroline	phen	0	N, N	2
Glycinate	gly⁻	−1	N, O	2
Dimethylglyoximate	dmgH⁻	−1	N, N	2
Diethylenetriamine	dien	0	N, N, N	3
Porphyrin (as in haem)	por²⁻	−2	4 × N	4
EDTA	EDTA⁴⁻	−4	2 × N, 4 × O	6

The names in brackets in the first column are the forms used inside complex names; see how to name coordination compounds.

Group 1: Monodentate ligands (one donor atom)

A monodentate ligand attaches through a single atom. These are the building blocks of most introductory complex chemistry.

1. Water, H₂O

Why it matters: almost every metal ion dissolved in water is really an aqua complex, such as [Fe(H₂O)₆]³⁺ or [Cu(H₂O)₆]²⁺. Water donates a lone pair from oxygen. Because the metal pulls electron density from the O–H bonds, coordinated water is more acidic than free water. That is why solutions of Fe³⁺ and Al³⁺ salts are noticeably acidic. Water is a mid-strength ligand in the spectrochemical series, so it is the usual reference point when comparing colours.

2. Ammonia, NH₃

Why it matters: ammonia is a slightly stronger ligand than water and is the classic substituting ligand in school chemistry. Excess ammonia turns pale blue copper(II) solutions deep blue, forming [Cu(NH₃)₄(H₂O)₂]²⁺. NH₃ and H₂O are similar in size and both neutral, so swapping one for the other usually keeps the coordination number at six.

3. Chloride, Cl⁻

Why it matters: chloride is larger than water and carries a negative charge, so fewer fit around a metal ion. Adding concentrated hydrochloric acid often drops the coordination number from six to four, as in yellow-green [CuCl₄]²⁻ or blue [CoCl₄]²⁻. Chloride is a weak-field ligand, so its complexes usually have small d-orbital splitting. The other halides (F⁻, Br⁻, I⁻) behave similarly, with iodide the weakest field of the common ligands.

4. Hydroxide, OH⁻

Why it matters: hydroxide rarely replaces all six waters in simple school reactions. Instead, it removes protons from coordinated water, leaving neutral hydroxides such as Cu(OH)₂ or Fe(OH)₃ that precipitate. Those coloured precipitates are the basis of qualitative tests for transition metal ions. With amphoteric metals, such as zinc and aluminium, excess OH⁻ does form soluble hydroxido complexes, for example [Zn(OH)₄]²⁻.

5. Cyanide, CN⁻

Why it matters: cyanide sits near the top of the spectrochemical series. It binds through carbon, accepts electron density back from the metal into empty π* orbitals, and produces very large d-orbital splitting. Complexes such as [Fe(CN)₆]⁴⁻ are low-spin and extremely stable, which is why the cyanide in hexacyanidoferrates and in the pigment Prussian blue stays locked to the iron rather than behaving like free cyanide. Cyanide’s grip on the iron in cytochrome oxidase is also what makes it a poison.

6. Carbon monoxide, CO

Why it matters: CO is the strongest-field common ligand and the ligand of organometallic chemistry. It donates a lone pair from carbon (σ donation) and accepts d-electron density into its π* orbitals (π back-donation). This two-way bonding stabilises metals in very low oxidation states, giving neutral carbonyls such as [Ni(CO)₄], [Fe(CO)₅] and [Cr(CO)₆], where the metal’s oxidation state is zero. CO binds haemoglobin’s iron far more tightly than O₂ does, which is the chemistry behind carbon monoxide poisoning (see haemoglobin).

7. Thiocyanate, SCN⁻

Why it matters: thiocyanate is ambidentate: it can bond through sulfur or through nitrogen. It is famous in the lab for the blood-red colour it gives with Fe³⁺, a sensitive test for iron(III). Which end binds depends on the metal. As a rough guide, harder metal ions such as Fe³⁺ prefer N, while softer ones such as Hg²⁺ and Pt²⁺ prefer S.

8. Nitrite, NO₂⁻

Why it matters: another ambidentate ligand, and historically important. Pentaamminecobalt(III) complexes with N-bonded and O-bonded nitrite can be isolated as two distinct compounds of different colour. This linkage isomerism gave early coordination chemists a clear example of isomers that differ only in which atom touches the metal.

9. Pyridine, C₅H₅N (py)

Why it matters: pyridine is the simplest aromatic N-donor. It is a slightly stronger field ligand than ammonia and is useful for dissolving metal complexes in organic solvents. More importantly, it is the “half” from which the bidentate ligands bipy and phen are built.

10. Triphenylphosphine, P(C₆H₅)₃ (PPh₃)

Why it matters: phosphines are soft, bulky donors that stabilise metals in low oxidation states, especially the platinum-group metals. Wilkinson’s catalyst, [RhCl(PPh₃)₃], uses them for the catalytic hydrogenation of alkenes. The size of a phosphine can be tuned by changing its substituents, which lets chemists control how many ligands fit around a metal and how reactive it is.

Group 2: Bidentate ligands (two donor atoms)

A bidentate ligand holds the metal with two donor atoms, forming a ring. Rings of five atoms (including the metal) are especially stable because they have little strain. Complexes containing such rings are called chelates.

11. Ethane-1,2-diamine, H₂NCH₂CH₂NH₂ (en)

Why it matters: en is the standard bidentate ligand. Each nitrogen donates a lone pair, and the M–N–C–C–N ring has five members. Replacing six ammonia molecules with three en molecules releases three extra particles into solution, an entropy gain that makes [Ni(en)₃]²⁺ far more stable than [Ni(NH₃)₆]²⁺. This is the chelate effect in its simplest form. Tris(en) complexes are also chiral, giving a clean example of optical isomerism in octahedral complexes.

12. Oxalate (ethanedioate), C₂O₄²⁻

Why it matters: oxalate binds through one oxygen on each carboxylate group, again forming a five-membered ring. It gives well-defined anionic complexes such as [Fe(C₂O₄)₃]³⁻, which is light-sensitive and is used in actinometry, a way of measuring light intensity chemically. Oxalate is also the ligand behind calcium oxalate, the main mineral in many kidney stones.

13. Acetylacetonate, CH₃COCHCOCH₃⁻ (acac⁻)

Why it matters: acac⁻ forms a six-membered chelate ring through two oxygens. With M³⁺ ions, three acac⁻ ligands give neutral complexes such as [Cr(acac)₃] and [Fe(acac)₃]. Being neutral and wrapped in hydrocarbon groups, these dissolve in organic solvents and can often be sublimed, which makes them useful as metal sources in materials chemistry.

14. 2,2′-Bipyridine (bipy)

Why it matters: two pyridine rings joined side by side, with both nitrogens pointing the same way. Bipy is a strong-field ligand. Its ruthenium(II) complex, [Ru(bipy)₃]²⁺, absorbs visible light and has a long-lived excited state, so it is a workhorse in photochemistry and studies of light-driven catalysis.

15. 1,10-Phenanthroline (phen)

Why it matters: phen is like bipy but with the two rings locked flat by a third ring. With Fe²⁺ it forms the intensely red-orange [Fe(phen)₃]²⁺ ion, known as ferroin. Ferroin is used as a redox indicator and as the basis of a sensitive colorimetric method for measuring iron in water (see colorimetry).

16. Glycinate, H₂NCH₂COO⁻

Why it matters: the simplest amino acid anion binds through its amino nitrogen and one carboxylate oxygen. It shows that biological molecules make natural chelating ligands, and it explains why amino acids and proteins bind metal ions such as copper and zinc so readily.

17. Dimethylglyoximate (dmgH⁻)

Why it matters: two dmgH⁻ ligands wrap around Ni²⁺ to give a flat, square planar complex that precipitates as a bright red solid. The reaction is selective and sensitive enough to be a classic test and gravimetric method for nickel. Hydrogen bonds between the two ligands help hold the complex rigid.

Group 3: Polydentate ligands (three or more donor atoms)

18. Diethylenetriamine (dien)

Why it matters: dien is en with an extra –CH₂CH₂NH– unit, giving three nitrogen donors in a row. Two dien ligands can fill all six positions of an octahedron. It is a useful stepping stone for seeing how stability rises as denticity increases.

19. Porphyrins (tetradentate)

Why it matters: a porphyrin is a large, flat ring with four nitrogen atoms pointing inward. After losing two protons it becomes a 2− ligand that holds a metal ion in the centre of the ring. Haem (iron in a porphyrin) carries oxygen in haemoglobin, chlorophyll holds magnesium in a closely related ring (see chlorophyll), and vitamin B₁₂ uses a similar ring called a corrin around cobalt. The rigid ring leaves the positions above and below the metal free, which is exactly where O₂ binds in haemoglobin.

20. EDTA⁴⁻ (hexadentate)

Why it matters: ethylenediaminetetraacetate uses two nitrogens and four carboxylate oxygens to occupy all six octahedral positions of a single metal ion. One EDTA replacing six waters increases the number of free particles from 2 to 7, a large entropy gain. EDTA forms 1:1 complexes with almost every metal ion, which makes it the standard titrant for measuring water hardness (see complexometric titration with EDTA and the water hardness calculator). It is also used to treat lead poisoning and to stop metal ions spoiling foods and cosmetics.

Bonus: where the ligands sit in the spectrochemical series

Ligands differ in how strongly they split the metal’s d orbitals. A commonly quoted order, from weak field to strong field, is:

I⁻ < Br⁻ < Cl⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < py < NH₃ < en < bipy < phen < NO₂⁻ < PPh₃ < CN⁻ < CO

Halides are at the weak end, oxygen donors in the middle, nitrogen donors above them, and ligands that accept π back-donation (CN⁻, CO) at the top. This order decides colour and whether a complex is high-spin or low-spin, covered in crystal field theory.

Common mistakes

• Counting ligands instead of donor atoms. [Co(en)₃]³⁺ has three ligands but a coordination number of six.
• Assuming every atom with a lone pair donates. In en, only the two nitrogens bond. In EDTA, only the carboxylate oxygens that point towards the metal bond, not the C=O oxygens.
• Thinking charged ligands are always stronger. CO is neutral yet sits at the top of the spectrochemical series, while negatively charged I⁻ sits at the bottom.
• Mixing up bidentate and bridging. A bidentate ligand grips one metal with two atoms; a bridging ligand links two different metals.

Key takeaways

• Ligands are Lewis bases; what matters is the donor atom, the charge and the number of donor atoms (denticity).
• Monodentate ligands (H₂O, NH₃, Cl⁻, CN⁻, CO) are the basics; ambidentate ones (SCN⁻, NO₂⁻) can bind through either end.
• Bidentate ligands (en, oxalate, acac⁻, bipy, phen) form chelate rings, most stable when five-membered.
• Polydentate ligands (porphyrins, EDTA) make the most stable complexes, thanks largely to the chelate effect.
• A ligand’s place in the spectrochemical series controls the colour and spin state of its complexes.