Covalent Bonding — Print · Grade 11 Science

Covalent Bonding

Learning goals

• I can explain how nonmetal atoms share electron pairs to form single, double, and triple covalent bonds and complete their octets.
• I can draw simple Lewis structures and identify bonding pairs and lone pairs.

Readings for this unit

• Why three solids need three different explanations

A covalent bond involves shared electron density in an interacting atomic arrangement. A Lewis line represents a shared pair for bookkeeping, not a rigid connection or atoms consciously reaching a compromise. Covalent bonding can produce discrete molecules, such as H₂O, or extended network structures, such as diamond. The octet rule is a useful pattern for many main-group examples, with known exceptions; it does not itself supply a physical cause or a complete energy account.

Picture two hydrogen atoms drifting together. Each brings a single electron, and rather than one stealing from the other, they pool their two electrons into one shared bonding pair that both nuclei attract. That shared pair — one pair, two electrons — is a single covalent bond, often drawn as a simple line: H–H. Not every electron gets shared, though. In water, the oxygen atom shares one pair with each hydrogen but keeps two pairs entirely to itself; electrons that aren't shared are called lone pairs, and, as you'll see in the next lesson, they quietly shape the whole molecule.

A double bond is represented by two shared pairs in a simple Lewis model, and a triple bond by three. Comparable atom pairs often have shorter, stronger bonds at higher order, but this is not a universal comparison across unrelated species or a claim that bond energy is an exact multiple of pair count. Nitrogen’s strong N≡N bond helps explain some kinetic barriers, while actual reactivity still depends on pathway and conditions.

Chemists keep track of all these shared and unshared electrons with Lewis structures — simple maps that show every valence electron as a dot and every bonding pair as a line. To draw one, you add up the valence electrons from all the atoms, connect the atoms with single bonds, and then hand out the leftover electrons as lone pairs until each atom (except hydrogen, which is content with just two) is surrounded by eight. If there aren't enough electrons to give everyone an octet with single bonds, you pull lone pairs in to make double or triple bonds. The Lewis structure of water, for instance, shows two O–H bonding pairs and two lone pairs resting on the oxygen.

In many molecular substances, ordinary melting or boiling changes arrangements and intermolecular interactions while the molecules’ internal covalent connections remain intact. Extended covalent networks require another structural account. Melting, boiling and conductivity depend on size, shape, phase and electronic/ionic carriers, so low boiling or poor conduction is not a definition of every covalent substance.

So covalently bonded atoms build an astonishing range of substances through varied bonding arrangements — the O₂ you breathe, the CO₂ you exhale, the water that fills the oceans, and the enormous molecules of DNA and protein that run every living cell. Understanding *how many* pairs atoms share, and *where* the lone pairs sit, is the launch pad for the final piece of the bonding story: the three-dimensional shapes these molecules fold into, and why those shapes matter so much.

Model check: Sharing need not be equal; polar covalent bonds share electron density unevenly. Three shared pairs give six bonding electrons, not three.

Total Valence Electrons (for a Lewis structure)

total valence e⁻ = sum of every atom's valence electrons (add 1 for each unit of negative charge, subtract 1 for each unit of positive charge)

the first step in drawing a Lewis structure. Example: H₂O = 6 (from O) + 1 + 1 (from the two H) = 8 electrons = 4 pairs.

Electrons Shared in a Bond

shared electrons = 2 × (number of shared pairs)

a single bond shares 1 pair = 2 electrons, a double bond shares 2 pairs = 4 electrons, and a triple bond shares 3 pairs = 6 electrons.

Worked examples

Avoid these mistakes

• Sharing need not be equal; polar covalent bonds share electron density unevenly.

Words to know

covalent bond

Bonding involving shared electron density between atoms; Lewis lines represent shared pairs in a simplified model.

molecule

A neutral group of atoms held together by covalent bonds.

bonding pair

A pair of electrons shared between two atoms, forming a covalent bond.

lone pair

A pair of valence electrons on an atom that is not shared in a bond.

Lewis structure

A diagram that shows an atom's or molecule's valence electrons as dots and each bonding pair as a line.

double bond

A covalent bond in which two pairs of electrons (four electrons) are shared.

triple bond

A covalent bond in which three pairs of electrons (six electrons) are shared.

diatomic molecule

A molecule made of exactly two atoms, such as O₂ or N₂.

The chemist who invented the dot diagram

The shared-pair picture you use today came from Gilbert N. Lewis, who in 1916 proposed that atoms bond by sharing pairs of electrons — years before physicists fully understood what an electron even was. He first sketched electrons as dots on the corners of a cube, and that humble idea grew into the Lewis structures chemists still draw a century later. Lewis was nominated for the Nobel Prize dozens of times but never won it, one of the most famous oversights in the history of chemistry.

When the octet rule bends

The octet rule is a superb guide, not an unbreakable law. Hydrogen is satisfied with just two electrons (a 'duet'), not eight. Boron happily settles for only six electrons around it in a molecule like BF₃, staying electron-hungry. And larger atoms from the third row down, like the sulfur in SF₆, can hold more than eight in an 'expanded octet.' Knowing when the rule bends is part of thinking like a chemist rather than memorizing a slogan.

The covalent solids that refuse to melt

You just learned that molecular substances melt at low temperatures — so how can diamond, pure covalently bonded carbon, be the hardest natural material, melting only above 3500°C? The trick is that diamond isn't built from separate little molecules. Every carbon is bonded to four others in one continuous covalent network that runs through the whole crystal, so melting it means shattering countless strong bonds at once. Quartz (SiO₂) and the silicon carbide in sandpaper are network solids too — a giant exception well worth remembering.

The strongest bond in the air you breathe

About four-fifths of every breath is nitrogen gas, N₂, locked by a triple bond so strong (roughly 945 kilojoules per mole) that the gas is almost completely unreactive. That's mostly a blessing — but it also means plants can't use the nitrogen floating right above them. Only lightning and certain bacteria can crack the triple bond naturally, and in the early 1900 s the Haber–Bosch process learned to do it industrially under crushing heat and pressure. The fertilizer it makes now helps feed roughly half the people on Earth.

Why carbon builds a whole branch of chemistry

One covalent atom is so versatile that an entire science is named after it. Carbon has four valence electrons, so it forms four strong covalent bonds and, uniquely, bonds happily to itself — chaining into rings, sheets, and branches of almost unlimited length. That single talent, called catenation, is why there are millions of carbon compounds, from the methane in a stove to the DNA in your cells, and why organic chemistry exists as its own enormous field.

In the real world

The oxygen you breathe (O₂, held by a double bond) and the nitrogen that makes up most of the air (N₂, held by a very strong triple bond) are both covalent molecules. That tough nitrogen triple bond is why the gas is so unreactive — and why turning it into fertilizer to feed the world takes enormous energy.

Exit check

How many electrons are in a triple bond?