Dative (Coordinate) Covalent Bonds Explained

In an ordinary covalent bond, each atom contributes one electron to the shared pair. But sometimes both electrons come from the same atom. That’s a dative covalent bond, also called a coordinate bond. Once it has formed, it’s just like any other covalent bond — but knowing where the electrons came from explains a wide range of chemistry, from why ammonia is a base to how haemoglobin carries oxygen and why transition metal compounds are so colourful.

Definition

A dative covalent bond is a covalent bond in which both electrons in the shared pair come from one atom.

For it to form, two things are needed:

1. A donor atom with a lone pair of electrons.
2. An acceptor atom with an empty orbital that can take the pair.

How to draw it

In displayed formulas, a dative bond is often shown as an arrow pointing from the donor atom to the acceptor atom:

H₃N → BF₃

The arrow shows where the electron pair came from. In dot-and-cross diagrams, both electrons in the dative bond are drawn with the same symbol (both dots or both crosses) (see how to draw dot-and-cross diagrams). Some books draw an ordinary line instead, with formal charges added — both are correct.

Example 1: the ammonium ion, NH₄⁺

Ammonia, NH₃, has a lone pair on nitrogen. A hydrogen ion, H⁺, has no electrons at all — just a proton, with an empty 1s orbital. When ammonia reacts with an acid, nitrogen donates its lone pair to H⁺:

NH₃ + H⁺ → NH₄⁺

The new N–H bond is dative. But here’s the key point: once formed, all four N–H bonds in NH₄⁺ are identical. They have the same length and energy, and the ion is a perfect tetrahedron. You can’t tell which bond was the dative one. The positive charge is spread over the whole ion (formally it sits on nitrogen — see formal charge).

This reaction is exactly what makes ammonia a base: it accepts a proton by donating a lone pair.

Example 2: the oxonium ion, H₃O⁺

When an acid dissolves in water, the H⁺ doesn’t float around on its own. It bonds to a water molecule, which donates one of oxygen’s two lone pairs:

H₂O + H⁺ → H₃O⁺

The oxonium ion (also called hydronium) has one lone pair left on oxygen and a trigonal pyramidal shape, like ammonia. Whenever you see H⁺(aq) in an equation, it’s really H₃O⁺ (and more complex clusters of water molecules around the proton) (see the pH scale explained).

Example 3: ammonia–boron trifluoride, NH₃BF₃

Boron trifluoride is a molecule with an incomplete octet: boron has only six electrons around it, leaving an empty p orbital (see Lewis structures practice). Ammonia can donate its lone pair into that empty orbital:

H₃N: + BF₃ → H₃N→BF₃

The product is a stable white solid. Both nitrogen and boron now have full octets. The shape changes, too: BF₃ is flat (trigonal planar), but in the product boron is tetrahedral.

This is the textbook example of a Lewis acid–base reaction: the Lewis base (NH₃) donates an electron pair; the Lewis acid (BF₃) accepts it (see Lewis acids and bases). Every dative bond forms in this way.

Example 4: aluminium chloride dimers, Al₂Cl₆

Aluminium chloride, AlCl₃, is like BF₃: aluminium has only six electrons around it. In the vapour at moderate temperatures, and in the solid’s structure, two AlCl₃ units join together. A chlorine atom on each molecule donates one of its lone pairs to the aluminium of the other, forming two dative bonds:

2AlCl₃ ⇌ Al₂Cl₆

Each aluminium ends up surrounded by four chlorines in a roughly tetrahedral arrangement, and two chlorine atoms bridge the two aluminium atoms. At higher temperatures, the dimer breaks apart again into AlCl₃ molecules.

Aluminium chloride’s appetite for electron pairs makes it an important catalyst in organic chemistry (for example, in Friedel–Crafts reactions on benzene).

Example 5: carbon monoxide

In the best Lewis structure of CO, carbon and oxygen share three pairs of electrons: C≡O. Counting electrons shows that oxygen contributes four of the six bonding electrons and carbon only two. So one of the three bonds can be thought of as dative, from oxygen to carbon. This leaves carbon with a formal charge of −1 and oxygen +1, and helps explain why CO bonds to metals through carbon — including the iron in haemoglobin, which is why carbon monoxide is so toxic (see carbon monoxide poisoning).

Example 6: metal complexes

Dative bonds are at the heart of coordination chemistry. Transition metal ions have empty orbitals that can accept electron pairs from molecules or ions called ligands. Each ligand donates a lone pair to the metal, forming a coordinate bond. The result is a complex ion:

Complex	Metal ion	Ligands	Colour
[Cu(H₂O)₆]²⁺	Cu²⁺	6 water molecules	pale blue
[Cu(NH₃)₄(H₂O)₂]²⁺	Cu²⁺	4 ammonia, 2 water	deep blue
[Fe(H₂O)₆]³⁺	Fe³⁺	6 water	yellow-brown (in solution)
[Ag(NH₃)₂]⁺	Ag⁺	2 ammonia	colourless
[Fe(CN)₆]⁴⁻	Fe²⁺	6 cyanide ions	pale yellow

The number of dative bonds to the metal is the coordination number — 6, 4 and 2 are common, giving octahedral, tetrahedral or square planar, and linear shapes.

Some ligands can form more than one dative bond each. EDTA, for example, wraps around a metal ion and forms six dative bonds, gripping it very tightly. This is used to measure water hardness and to treat heavy-metal poisoning (see complexometric titration with EDTA).

Dative bonds to metals are everywhere in biology:

• Haemoglobin: an Fe²⁺ ion is held by four nitrogen atoms in the haem ring and one from a histidine in the protein; oxygen binds to the sixth position (see haemoglobin).
• Chlorophyll: a Mg²⁺ ion held by four nitrogen atoms (see chlorophyll).
• Vitamin B₁₂: a cobalt ion at the centre of a ring.

How to spot a dative bond

Look for:

1. An atom with a lone pair (N in NH₃, O in H₂O, halide ions, CO).
2. An atom or ion with an empty orbital or an incomplete octet (H⁺, B in BF₃, Al in AlCl₃, transition metal ions).
3. An atom with more bonds than its usual valency: nitrogen with four bonds (NH₄⁺), oxygen with three (H₃O⁺), or boron with four ([BF₄]⁻).

Is a dative bond different once formed?

Chemically, no. A dative bond is a covalent bond like any other: a shared pair of electrons attracted to two nuclei. The label only describes how it formed. In NH₄⁺, H₃O⁺ and [BF₄]⁻, all the bonds become equivalent. In metal complexes, the metal–ligand bonds often have significant ionic character too.

Common misconceptions

• “Dative bonds are weaker than normal covalent bonds.” Not necessarily; strength depends on the atoms involved, not on where the electrons came from.
• “You can tell which bond in NH₄⁺ is dative.” All four N–H bonds are identical.
• “Only metals can accept electron pairs.” H⁺, boron and aluminium compounds are classic non-metal and main-group acceptors.
• “The arrow means electrons flow continually.” It just records the origin of the pair.

Key takeaways

• A dative (coordinate) covalent bond is a shared pair where both electrons come from one atom.
• It needs a lone-pair donor and an acceptor with an empty orbital; drawn as an arrow from donor to acceptor.
• Examples: NH₄⁺, H₃O⁺, NH₃BF₃, Al₂Cl₆, CO and metal complexes.
• Once formed, a dative bond is identical to other covalent bonds.
• Every dative bond is a Lewis acid–base interaction.