Electron Sharing: The Covalent Bond

Introduction

Non-metal atoms often join without forming the separate positive and negative ions used in a simple salt model. Instead, electrons are shared between them. Covalent bonding explains molecules such as hydrogen, water and methane, as well as extended structures such as diamond. A shared-pair drawing is the first step towards describing this electron distribution.

Core explanation

In an introductory covalent model, a single bond consists of one shared pair of electrons. Electron density between the atoms is attracted by both positively charged nuclei. A stable bond reflects the balance of attractions and repulsions at an energetically favourable separation. “Sharing” names the distribution; attraction explains its bonding effect.

Hydrogen provides the simplest example. Each neutral atom supplies one electron, and the two electrons form the shared pair in H₂. Each hydrogen counts that pair towards a two-electron first shell. The molecule contains two electrons altogether, not four: the local shell counts overlap because both refer to the same pair.

For many familiar main-group atoms, shared pairs and lone pairs together give an octet around each atom. Chlorine in Cl₂ has one shared pair and three lone pairs around each atom. Oxygen in water has two shared pairs and two lone pairs. These counts are useful, but the octet pattern is not a universal law for every molecule.

Sharing need not be equal. Different atoms can attract bonding electrons with different strengths, producing a polar covalent bond. Even then, a covalent drawing does not usually show fully transferred electrons and separate monatomic ions. Those are different idealised descriptions.

Covalent bonds do not always stop at small molecules. In diamond they continue throughout a giant network. Consequently, the presence of covalent bonds alone cannot establish low melting temperature, volatility or poor mechanical strength. To connect bonding with properties, determine whether the bonds occur within separate molecules or extend through the whole solid.

Step-by-step reasoning

1. Count the valence electrons available from the participating atoms. 2. Place two electrons in each proposed single bonding region. 3. Count shared pairs around both atoms while counting each electron only once in the molecular total. 4. Add remaining lone pairs and check the appropriate duet or octet pattern without ignoring known exceptions.

Visual explanation

Draw two overlapping hydrogen shell circles. Put one dot and one cross in the overlap and label it “one shared pair.” Under the complete drawing write “two electrons total”; beside each hydrogen write “local duet.”

Real-world analogy

One shared bridge provides a connection for people on both riverbanks, while remaining a single bridge. A bonding pair similarly contributes to the local electron count of both atoms. The analogy concerns shared access, not a literal solid bridge made from stationary electrons.

Real-world example

Water contains two O–H covalent bonds within each molecule. When water evaporates under ordinary conditions, those molecules generally remain intact. The change mainly concerns their separation and intermolecular interactions, demonstrating why intramolecular bonding must be distinguished from bulk physical changes.

Why?

Why does sharing help hold nuclei together even though the nuclei repel? The negatively charged electron distribution attracts both nuclei. The overall energy includes these attractions and the repulsions; at the bond's preferred separation the bonded state can be lower in energy than separated atoms.

Common misconception

“Each atom owns one electron forever after the bond forms.” Dot-and-cross origin labels are bookkeeping. Electrons are indistinguishable, and the bonded electron distribution cannot be described as permanently tagged particles circulating exclusively around their original atoms.

Worked example

Two chlorine atoms supply seven valence electrons each, fourteen total. Put one shared pair between them, using one electron from each atom. Twelve electrons remain, arranged as three lone pairs on each chlorine. Each chlorine locally counts six lone-pair electrons plus two bonding electrons, giving eight; the whole molecule still contains fourteen valence electrons.

Quick check

1. How many electrons are represented by a single covalent bond in the shared-pair model? Answer: Two electrons, forming one shared pair between the atoms.

Exam focus

Define the bond with shared electrons and attraction to both nuclei. Avoid saying that every covalent substance consists of small molecules or that all sharing is equal.

Advanced insight

Orbital models describe covalent bonding in terms of quantum states and electron density rather than literal electron paths. Some bonding cannot be captured by a single two-centre pair, but the shared-pair model remains effective for many introductory Lewis structures.

Summary

Covalent bonding involves shared electron density attracted to neighbouring nuclei. A single bond is represented by one shared pair, counted around both atoms but once in the total inventory. Covalent structures may be molecular or extended, and sharing can be unequal.

Practice questions

1. Why does H₂ have two electrons rather than four despite both hydrogens counting a duet? Answer: Both local counts include the same shared pair. 2. Does covalent bonding always produce a small molecule? Answer: No. Diamond is an extended covalent network. 3. What is missing from the statement “a covalent bond forms because atoms share”? Answer: An explanation of the attraction between the shared electron density and both nuclei, together with the resulting favourable energy balance.