Covalent Bonding: Sharing Electrons in Molecules

Sharing instead of transferring

When two nonmetal atoms meet, neither can overpower the other the way sodium overpowers chlorine. Their electronegativities are too close, so instead of transferring electrons they compromise: each atom contributes one or more electrons to a shared pair that belongs to both nuclei at once.

Hydrogen shows the mechanism cleanly. Each H atom brings a single 1s electron; the two electrons pair up in the region between the nuclei, and both protons pull on that concentrated negative charge. The attraction between the shared pair and the two nuclei outweighs the nucleus-nucleus repulsion, so the joined atoms sit at a lower energy than the separated ones — by 436 kJ/mol, the H−H bond energy.

The atoms settle at an equilibrium distance of 74 pm, where attraction and repulsion balance. Push them closer and nuclear repulsion dominates; pull them apart and the shared pair can no longer hold both. Every covalent bond is this kind of compromise distance, and every molecule is a collection of such compromises frozen into a definite shape.

Single, double, and triple bonds

Atoms can share more than one pair. A single bond shares one pair, a double bond shares two, and a triple bond shares three — and each added pair shortens and strengthens the connection. Carbon-carbon bonds make the trend concrete:

• C−C single bond: 347 kJ/mol, 154 pm
• C=C double bond: 614 kJ/mol, 134 pm
• C≡C triple bond: 839 kJ/mol, 120 pm

Notice that a double bond is not twice a single bond. The first shared pair forms a sigma bond along the axis between nuclei; the second forms a pi bond from sideways orbital overlap, which is weaker and more exposed. That is why double bonds are reactive sites: the pi electrons are the first to be attacked in addition reactions.

Bond order — the number of shared pairs — also shapes molecules. Triple-bonded N2 is so tightly held (945 kJ/mol) that nitrogen gas barely reacts at room temperature, which is why the atmosphere's nitrogen sits idle while oxygen does the chemical work.

Coordinate covalent bonds

Sometimes both electrons in a shared pair come from the same atom. In ammonia, nitrogen donates its lone pair to a hydrogen ion (H+), forming the ammonium ion NH4+. The resulting N−H bond is called a coordinate covalent or dative bond.

Once formed, a coordinate bond is indistinguishable from any other covalent bond — in NH4+ all four N−H bonds are identical in length and strength. The label only records the bond's history, not its nature.

Carbon monoxide is the classic neutral example: carbon and oxygen share three pairs, one of them donated by oxygen's lone pair. Coordinate bonding matters far beyond curiosities — it is how metal ions grip ligands in complexes, how hemoglobin holds iron, and how many catalysts latch onto their substrates. Whenever an electron-rich atom meets an electron-hungry one, a dative bond is the likely outcome.

Network covalent solids: giant molecules

Not all covalent bonding stops at small molecules. In network covalent solids, sharing extends through the entire crystal: every atom is covalently bonded to its neighbors in a continuous three-dimensional web.

Diamond is pure carbon with each atom bonded tetrahedrally to four others; there are no discrete molecules, only one giant interconnected structure. Quartz (SiO2) is a similar network of silicon and oxygen. These materials inherit covalent bond strength at the macroscopic scale: diamond is the hardest known substance, and quartz melts at 1713 °C.

Contrast molecular solids like ice, sugar, or paraffin wax. Their molecules are internally covalent, but the molecules merely rest against each other through weak intermolecular forces — so they melt easily. The practical rule: ask whether covalent bonds run between the molecules or only within them. Between means a network solid with extreme properties; within means a molecular solid that melts and vaporizes readily.

Molecules across the three levels

Covalent substances are easiest to grasp by shuttling between chemistry's three levels. Macroscopically, you see water flowing, sugar dissolving sweetly, and plastic bending — behaviors of bulk matter. At the particulate level, you picture discrete molecules: two hydrogen atoms locked to one oxygen by shared pairs, long polymer chains tangled like spaghetti.

The symbolic level compresses all of this into formulas and structures. H2O tells you the atom ratio; a Lewis drawing shows the shared pairs; N≡N records a triple bond in four characters. Chemists read these symbols the way musicians read scores — as instructions for reconstructing the particulate reality.

Each level checks the others. If a symbolic structure predicts a polar molecule but the macroscopic substance refuses to mix with water, something is wrong with the structure. Fluency in covalent chemistry means translating freely: observation to particles to symbols and back.

Covalent compounds all around you

Covalent compounds dominate both nature and industry. The air is mostly N2 and O2; natural gas is CH4; the fuels in gasoline are covalent hydrocarbons. Every plastic — polyethylene bags, PVC pipes, nylon fibers — is a forest of long covalent chains.

Living things are covalent machinery. Proteins, DNA, sugars, and fats are all covalent molecules, and the reactions of life rearrange covalent bonds with enzyme guidance. Medicines are designed as covalent molecules shaped to fit protein pockets.

Even the materials of technology lean covalent: silicon chips are a covalent network, glass is a disordered silicon-oxygen network, and carbon fiber's strength comes from covalent sheets. From breath to smartphones, shared electron pairs do the structural work of the modern world.

Common mistakes

Real-world applications

• Burning natural gas (CH4) for heat and electricity

• Building plastics and fibers from long covalent chains

• Cutting and drilling with diamond's network lattice

• Designing medicines as precisely shaped molecules