3.1.3.2 - Covalent and Dative Bonding

Shared pairs in covalent bonds

A covalent bond forms when two atoms share a pair of electrons. The most important phrase is "shared pair". If you keep that idea in mind, the diagrams and definitions all become much easier to understand.

In a normal single covalent bond, one electron usually comes from each atom. Those two electrons are then shared between the two nuclei. Electrostatic attraction between the shared pair and both positive nuclei holds the atoms together.

When chemists stop drawing every electron individually, they replace each shared pair with a line. So a line is not just a convenient mark on the page: it stands for one shared pair of electrons.

• $\ce{H-H}$ shows one shared pair, so it is a single bond.
• $\ce{O}=\ce{O}$ shows two shared pairs, so it is a double bond.
• $\ce{N#N}$ shows three shared pairs, so it is a triple bond.

That means multiple bonds contain multiple shared pairs of electrons. A double bond is two shared pairs. A triple bond is three shared pairs. In a dot-and-cross diagram you can see the electrons themselves; in a displayed formula the lines are a quicker shorthand for exactly the same bonding idea.

The diagram below compares the same bonding idea in two forms, so you can see that the number of lines in a displayed formula matches the number of shared pairs in the dot-and-cross version.

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A common mistake is to say "a covalent bond is a line". That is not precise enough. The accurate statement is that a covalent bond is a shared pair of electrons, and we often represent that shared pair with a line.

How a dative covalent bond forms

Most covalent bonds form with one electron contributed by each atom. A co-ordinate bond is different only in how the shared pair is formed. Both electrons in the shared pair come from the same atom.

For this to happen, one atom must have a lone pair available to donate. A lone pair is a pair of outer-shell electrons that is not currently being used in bonding. The atom or ion receiving that pair must be able to accept it. This is often an electron-deficient species.

Ammonia is a useful example. In $\ce{NH3}$, the nitrogen atom has a lone pair. A hydrogen ion, $\ce{H^+}$, has no electrons at all, so it can accept a pair of electrons from nitrogen. Nitrogen donates both electrons, and a new shared pair forms between N and H.

The diagram below shows the lone pair on nitrogen in ammonia being donated to $\ce{H^+}$, and then the ammonium ion, $\ce{NH4^+}$, with the co-ordinate bond drawn as $\mathrm{N \rightarrow H}$.

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This new bond is shown using an arrow from the donor to the acceptor: $\mathrm{N \rightarrow H}$. The arrow points in the direction the electron pair was donated. It does not mean electrons are permanently moving along the bond. It simply shows how that shared pair was formed.

After the bond has formed, it is still a shared pair of electrons between two atoms. So a dative bond is still a covalent bond. The special feature is only the source of the two electrons in the shared pair. Once ammonium has formed, its four N–H bonds are equivalent; the arrow records how the bond formed, not a permanently different bond strength.

Counting and representing shared pairs

A dot-and-cross diagram keeps track of where electrons came from, while a displayed formula shows bonds as lines. Dots and crosses do not represent different kinds of electron: use one symbol consistently for the electrons supplied by one atom and a second symbol for the other atom.

For $\ce{O2}$, each oxygen starts with six outer electrons. Two shared pairs place four electrons between the atoms; each oxygen also has two lone pairs. Count twelve outer electrons in the whole diagram and eight around each oxygen when shared electrons are counted for both atoms. Replace the two shared pairs with two lines to obtain $\ce{O=O}$. Do not turn the lone pairs into bond lines.

For $\ce{N2}$, each nitrogen contributes five outer electrons. Three shared pairs and one lone pair on each nitrogen account for all ten electrons. The displayed formula is $\ce{N#N}$.

For a co-ordinate bond, draw an arrow from the atom supplying a lone pair to the atom accepting it. In forming $\ce{NH4+}$, this is N→H; place the complete ion in square brackets with an overall + charge outside. A line or arrow is a bond representation, not an electron trajectory. Reversing the donor arrow would claim that the electron-free hydrogen ion supplied the pair.