A coordinate bond is formed,When the shared electron pair is provided by one of the combining atoms.

The atom, which provides the electron pair is termed as the donor atom, while

 the other atom, which accepts it, is termed as the acceptor atom.

 Such a bond is also known as dative bond.

An arrow (-->) pointing from donor towards the acceptor atom represents a coordinate bond. When a one sided sharing of electrons takes place, the coordinate bond so formed cannot be distinguished from a normal covalent bond.

Formation of Coordinate Bond

The formation of such bonds is illustrated through some examples given below.

Formation of ammonium (NH⁺4) ion

During the formation of ammonium ion, nitrogen is the donor atom, while H⁺ is the acceptor ion as shown below:

Formation of ozone (O3)Molecule

A molecule of oxygen contains two oxygen atoms joined by a double covalent bond (An oxygen atom  is two electron short of neon configuration). Thus, the two atoms of oxygen share two electrons each . If an atom of oxygen having six electrons comes closer to the oxygen molecule, the new atom may share a lone pair of electrons from either of these two oxygen atoms, which donates to the third oxygen atom without sharing any of the electrons of the third oxygen atom. As a result, a coordinate bond is formed between one of the oxygen atoms of the oxygen molecule, and the third atom of oxygen. This is shown below:

Formation of a coordinate bond between two molecules

Sometimes, two or more stable molecules combine to form a molecular complex. In a complex molecule, the constituent molecules are held together by a 'coordinate bond'. One typical example involves the molecules of NH3 and BF3. The electron dot structures of these molecules are:

The nitrogen atom has a complete octet around it, but boron atom has only six electrons around it. The nitrogen atom therefore donates its lone pair of electrons to boron so that its atoms also acquire the octet. This one-sided sharing between N and B atoms gives rise to a coordinate bond. Coordinate bonds are involved in the formation of transition metal complexes known as coordination compounds.

Lewis structures of some typical co-ordinate covalent compounds

Step by Step instructions are designed to produce a correct Lewis Structure for any molecular assembly (neutral molecule or polyatomic ion) which contains only Main Group elements joined by two-atom shared-pair covalent bonds.

 

1. VSE - Count valence shell electrons

2. Connectivity - Arrange bonded atoms

3. BP - Assign bond pairs

4. PLP - Assign lone pairs to peripheral atoms

5. CLP - Assign lone pairs to central atoms

6. Rearrange - Find best Lewis Structure(s) [formal charge, resonance]

 

These rules are not appropriate for free radicals or molecular assemblies which contain multicenter bonds or transition metals.

 

Step 1: VSE - Count the total number of Valence Shell Electrons; divide these VSE into pairs.

1. Sum the number of valence shell electrons of each atom;
2. Subtract the charge on the assembly
3. The total should be even (these rules do no apply to assemblies with an odd number of electrons)

Examples:

HCN = 1+4+5-0 = 10 VSE = 5 VSE pairs

C₂H₆O = 2(4)+6(1)+6-0 = 20 VSE = 10 VSE pairs

CO₃^2- = 4+3(6)-(-2) = 24 VSE = 12 VSE pairs

PO₄^3- = 5+4(6)-(-3) = 32 VSE = 16 VSE pairs

H₃O⁺ = 3(1)+6-(+1) = 8 VSE = 4 VSE pairs

Step 2: Connectivity  -Arrange the atomic symbols so that covalently bonded atoms are contiguous.

Some general rules and definitions:

1.  The number of covalent bonds an atom forms is called its valency.
2. Some atoms have fixed valence. For example:

 C = 4, F = 1.

3. Some atoms have variable valence. For example:

O = 2 (sometimes 3), B, N = 3 (sometimes 4).

4. An atom bonded to only one other atom is peripheral (monovalent atoms such as H and F are always peripheral).
5. An atom bonded to two or more other atoms is central.

Step 3: Assign BP - Place one VSE pair of electrons between each bonded pair of atoms.

e.g.H3O+ = 4 VSE pairs - 3 BP = 1 pair remaining

Step 4: Assign Peripheral LP - Place up to three VSE pairs on each peripheral atom.

e.g. H3O⁺ : 1 pair, not assigned (since all peripheral atoms are H); 1 VSE pair remains.

Step 5: Assign Central LP - Place any remaining VSE pairs as Lone Pairs on central atom(s) according to the Rule of Orbitals.

 

The Rule of Orbitals: the total number of lone pairs and bond pairs (LP+BP) associated with an atom cannot exceed the number of Valence Shell Orbitals (VSO = n², where n is the row of the Periodic Table in which that atom resides).

n = 1 (H): maximum VSE pairs (LP+BP) = VSO = 1;

n = 2 (B, C, N, O, F): maximum VSE pairs (LP+BP) = VSO = 4 ("octet rule")

n = 3 ((Al, Si, P, S, Cl): maximum VSE pairs (LP+BP) = VSO = 9; etc.

Step 6: Rearrange VSE Pairs - If necessary, push electron pairs according to the Rule of Orbitals and the Principle of Electro-neutrality.

 

Principle of Electroneutrality: each atom in a covalent molecular assembly has a formal charge close to zero.

Formal Charge: FC = (Group Number) - (Bond Pairs) - 2(Lone Pairs)

Electron Pushing: formally changing a lone pair into a bond pair, or vice versa, while retaining association with the atom.

Examples

HCN

Original Lewis Structure

H: FC = 1-1-2(0) = 0;

H: rule of orbitals satisfied (1 orbital, 1 VSE pair);

C: FC = 4-2-2(0) = +2;

C: rule of orbitals not satisfied (4 orbitals, only 2 VSE pairs; it must be associated with 4 VSE pairs);

N: FC = 5-1-2(3) = -2;

N: rule of orbitals satisfied (4 orbitals, 4 VSE pairs)

Rearranged Lewis Structure

Push two lone pairs on N into C-N bonding position, creating a C-N triple covalent bond.

H: FC = 0;

C: FC = 4-4-2(0) = 0;

N: FC = 5-3-2(1) = 0;

All atoms: rule of orbitals satisfied.