What is the fundamental difference in electron contribution between a standard covalent bond and a coordinate bond?

In a standard covalent bond, each atom contributes one electron to the shared pair. In a coordinate bond, one atom (the donor) contributes both electrons to the shared pair.

How do the physical properties (length and strength) of a coordinate bond compare to a standard covalent bond once formed?

They are identical. Once the bond is formed, the electrons are shared equally between the nuclei, and there is no physical way to distinguish a coordinate bond from a standard covalent bond.

When comparing a donor and an acceptor, which one must have an empty orbital?

The acceptor must have an empty orbital to accommodate the incoming pair of electrons. The donor must have a lone pair of electrons to provide.

What is a common error when drawing the arrow representation for a dative covalent bond?

Drawing the arrow in the wrong direction. The arrow must always point from the donor atom (the source of the lone pair) toward the acceptor atom.

Why is it incorrect to assume that a coordinate bond is a type of ionic bond?

Because the electrons are shared between the two nuclei rather than being completely transferred. It retains the characteristics of a covalent bond, such as directional orbital overlap.

What mistake do students often make regarding the octet rule in coordinate bonding?

They may fail to realize that coordinate bonding is often the method by which an 'electron-deficient' atom (with fewer than 8 electrons) finally completes its octet.

Define 'Electron-Deficient' in the context of coordinate bonding.

An atom or molecule is electron-deficient if it has an unfilled valence shell (fewer than 8 electrons) and possesses an empty orbital available for bonding.

What is a 'Lone Pair' of electrons?

A lone pair is a pair of valence electrons that are not shared with another atom in a covalent bond and are available for donation in a coordinate bond.

What symbol is used to denote a coordinate bond in a structural formula?

An arrow ($ ightarrow$) pointing from the donor atom to the acceptor atom.

Why does a hydrogen ion ($H^+$) frequently participate in coordinate bonding?

A hydrogen ion has lost its only electron, leaving it with an empty $1s$ orbital. It is highly electron-deficient and seeks a lone pair to achieve a stable 'duet' configuration.

Coordinate Bonding | Cambridge International Examinations AS-Level Chemistry

Revision Notes

AS-Level

Cambridge International Examinations

Chemistry

3. Chemical Bonding

Coordinate Bonding

Summary

Coordinate bonding, also known as dative covalent bonding, is a variation of covalent bonding where the shared pair of electrons originates from a single atom. This mechanism allows electron-deficient species to achieve stable electronic configurations by accepting lone pairs from donor atoms.

1. Definition & Core Concepts

A coordinate bond is a type of covalent bond in which both electrons in the shared pair are provided by the same atom. This differs from a standard covalent bond where each participating atom typically contributes one electron to the shared pair.

The atom that provides the electron pair is referred to as the donor, while the atom that receives the pair into its empty orbital is known as the acceptor.

For a coordinate bond to form, the donor must possess at least one lone pair of electrons (a pair not already involved in bonding), and the acceptor must be electron-deficient, meaning it has an unfilled outer orbital.

Diagram showing a donor atom with a lone pair and an acceptor atom with an empty orbital, with an arrow indicating the coordinate bond formation.

2. Underlying Principles

The formation of a coordinate bond is driven by the drive for atoms to reach a more stable, lower-energy state, often by completing their octet (eight electrons in the valence shell).

The bond involves the overlap of orbitals: the filled orbital of the donor atom overlaps with the empty orbital of the acceptor atom, creating a new molecular orbital that contains both electrons.

Once the bond is established, the electrons are held by the electrostatic attraction of both nuclei, making the bond indistinguishable from a standard covalent bond in terms of strength and length.

3. Methods & Representation

In Lewis structures and displayed formulas, a coordinate bond is represented by an arrow ( $ightarrow$ ). The arrow starts at the donor atom and points toward the acceptor atom.

When calculating formal charges, the donor atom effectively 'loses' one electron to the acceptor's side of the bond, often resulting in a positive formal charge for the donor and a negative formal charge for the acceptor.

In complex ions, the entire species is often enclosed in square brackets with the overall charge indicated outside, reflecting the sum of all formal charges within the molecule.

4. Key Distinctions

5. Exam Strategy & Tips

6. Common Pitfalls & Misconceptions

Coordinate Bonding

Summary

1. Definition & Core Concepts

The atom that provides the electron pair is referred to as the donor, while the atom that receives the pair into its empty orbital is known as the acceptor.

Diagram showing a donor atom with a lone pair and an acceptor atom with an empty orbital, with an arrow indicating the coordinate bond formation.

2. Underlying Principles

The formation of a coordinate bond is driven by the drive for atoms to reach a more stable, lower-energy state, often by completing their octet (eight electrons in the valence shell).

Once the bond is established, the electrons are held by the electrostatic attraction of both nuclei, making the bond indistinguishable from a standard covalent bond in terms of strength and length.

3. Methods & Representation

In Lewis structures and displayed formulas, a coordinate bond is represented by an arrow ( $ightarrow$ ). The arrow starts at the donor atom and points toward the acceptor atom.

In complex ions, the entire species is often enclosed in square brackets with the overall charge indicated outside, reflecting the sum of all formal charges within the molecule.

4. Key Distinctions

Feature	Standard Covalent Bond	Coordinate (Dative) Bond
Electron Source	One electron from each atom	Both electrons from one atom
Initial State	Two unpaired electrons	One lone pair + one empty orbital
Representation	Solid line (—)	Arrow ($
ightarrow$)
Post-formation	Identical properties	Identical properties

It is critical to recognize that the 'arrow' notation is merely a bookkeeping tool to show the origin of the electrons; it does not imply that the bond behaves differently once it exists.

5. Exam Strategy & Tips

Identify Lone Pairs: Always check the central atom of a molecule (like Nitrogen in $NH_3$ or Oxygen in $H_2O$ ) for lone pairs that could be donated to form a dative bond.
Spot Electron Deficiency: Look for Group 13 elements (like Boron or Aluminium) or hydrogen ions ( $H^+$ ) which often have empty orbitals and act as acceptors.
Arrow Direction: Ensure the arrow always points from the 'rich' (donor) to the 'poor' (acceptor). Reversing this is a common error that loses marks.
Dimerization: Be aware that some molecules, like $AlCl_3$ , form dimers ( $Al_2Cl_6$ ) specifically through coordinate bonding to achieve stability at lower temperatures.

6. Common Pitfalls & Misconceptions

A common misconception is that coordinate bonds are weaker than standard covalent bonds. In reality, because the resulting electronic environment is the same, the bond energies are comparable.

Students often forget to account for the formal charge changes. When a neutral molecule like ammonia donates to a proton, the resulting ion carries a $+1$ charge because the nitrogen has effectively shared its 'private' pair.

Another error is confusing coordinate bonding with ionic bonding. While there is a transfer of 'ownership' of the electron pair's origin, the electrons remain shared between the two nuclei, maintaining a covalent character.