A covalent bond forms when two atoms share one or more pairs of electrons. Instead of one atom stealing an electron from another (which is how ionic bonds work), both atoms hold onto the shared electrons simultaneously, and the attraction between each atom’s nucleus and the shared electrons is what glues the molecule together. This sharing is the foundation of nearly all the molecules you encounter daily, from the water you drink to the oxygen you breathe to the proteins in your muscles.
Why Atoms Share Electrons
Every atom “wants” a full outer shell of electrons, because that arrangement is especially stable. For most elements, a full shell means eight electrons, a guideline chemists call the octet rule. Hydrogen is the exception: it only needs two. Noble gases like neon and argon already have full shells, which is why they almost never react with anything.
Atoms that are short a few electrons can solve the problem by sharing. A hydrogen atom has one electron but needs two. When two hydrogen atoms each contribute their single electron to a shared pair, both atoms effectively “see” two electrons in their outer shell. The resulting hydrogen molecule (H₂) is far more stable than either atom alone. Carbon has four electrons in its outer shell and needs eight, so it forms four covalent bonds. That’s why carbon shows up in so many different molecules: it has four connection points to work with.
Single, Double, and Triple Bonds
The number of electron pairs two atoms share determines whether the bond is single, double, or triple:
- Single bond: one shared pair (2 electrons). The bond between hydrogen and carbon in methane (CH₄) is a single bond.
- Double bond: two shared pairs (4 electrons). Each carbon-oxygen bond in carbon dioxide (CO₂) is a double bond.
- Triple bond: three shared pairs (6 electrons). The two nitrogen atoms in atmospheric nitrogen (N₂) are held together by a triple bond.
More shared electrons means a stronger, shorter bond. Looking at carbon-carbon bonds as an example: a single C–C bond is about 154 picometers long and takes 376 kJ/mol of energy to break. A double C=C bond shrinks to 134 picometers and requires 728 kJ/mol. A triple C≡C bond is shorter still at 120 picometers and needs 965 kJ/mol to break. In practical terms, triple bonds are roughly two and a half times harder to break than single bonds.
Polar vs. Nonpolar Covalent Bonds
Not all atoms pull on shared electrons with equal strength. Electronegativity is the measure of how strongly an atom attracts electrons toward itself. When two identical atoms share electrons, they pull equally, and the bond is nonpolar. The H–H bond in hydrogen gas is a perfect example.
When the atoms are different, the more electronegative atom hogs the shared electrons slightly, creating an uneven charge distribution. Chemists classify these bonds by the electronegativity difference between the two atoms. A difference below 0.5 is considered nonpolar covalent. Between 0.5 and 2.1, the bond is polar covalent. Above 2.1, the electron transfer is so lopsided that the bond is essentially ionic.
Water is the classic polar molecule. Oxygen is significantly more electronegative than hydrogen, so the shared electrons in each O–H bond spend more time near the oxygen. This gives the oxygen end a slight negative charge and the hydrogen ends a slight positive charge. That polarity is the reason water is such an excellent solvent and why it has a relatively high boiling point for such a small molecule.
How Orbitals Overlap
At a slightly deeper level, covalent bonds form when the electron clouds (orbitals) of two atoms overlap. There are two types of overlap that produce two types of bonds. Sigma bonds form when orbitals overlap end to end, directly along the line connecting the two nuclei. This puts the highest concentration of shared electrons right between the atoms. Pi bonds form when orbitals overlap side to side, placing the electron density above and below that connecting line rather than along it.
Every single bond is a sigma bond. A double bond consists of one sigma bond plus one pi bond. A triple bond is one sigma bond plus two pi bonds. Pi bonds are generally weaker than sigma bonds because the side-to-side overlap is less direct, which is one reason double bonds aren’t simply twice as strong as single bonds.
How Bonds Shape Molecules
Covalent bonds don’t just hold atoms together. They determine a molecule’s three-dimensional shape. Electron pairs around a central atom repel each other and spread out as far apart as possible, like balloons tied together at their bases. This principle lets chemists predict molecular geometry based on the number of bonds and lone pairs (pairs of electrons not shared with another atom) around each atom.
Four groups of electrons around a central atom arrange themselves into a tetrahedron, with angles of about 109.5° between them. That’s the shape of methane, where carbon sits at the center of four evenly spaced hydrogen atoms. Water also has four electron groups around its oxygen (two bonds to hydrogen, plus two lone pairs), but because two of those groups are invisible lone pairs rather than bonds, the visible shape is bent rather than tetrahedral. Two electron groups produce a linear shape (180°), and three produce a flat triangle (120°). These shapes matter because they control how molecules interact with each other, how they fit into biological receptors, and whether they dissolve in water.
Coordinate Covalent Bonds
In a standard covalent bond, each atom contributes one electron to the shared pair. In a coordinate (or dative) covalent bond, both electrons come from the same atom. The result looks identical to any other covalent bond once it’s formed. The distinction is only in how it forms.
A common example is the ammonium ion (NH₄⁺). Ammonia (NH₃) has a lone pair of electrons on its nitrogen. When a bare hydrogen ion (just a proton, no electrons) approaches, nitrogen donates its lone pair to form a fourth N–H bond. All four bonds in the ammonium ion are identical in strength and length. You’d never be able to tell which one was the coordinate bond just by looking at the finished product. The hydronium ion (H₃O⁺), formed when water accepts a hydrogen ion, works the same way. Carbon monoxide also contains a coordinate bond: oxygen donates a lone pair to carbon on top of two regular shared bonds.
Properties of Covalent Compounds
Molecules held together by covalent bonds behave differently from ionic compounds in several noticeable ways. Most covalent compounds have low melting and boiling points. That’s because melting or boiling them doesn’t require breaking the covalent bonds themselves, just overcoming the much weaker attractions between neighboring molecules. Table sugar, rubbing alcohol, and vegetable oil are all covalent compounds that exist as liquids or soft solids at room temperature.
The exception is network covalent solids, where atoms are linked by continuous covalent bonds extending in all directions rather than forming discrete molecules. Diamond is pure carbon with every atom covalently bonded to four neighbors in a rigid lattice. Breaking that lattice means breaking actual covalent bonds, which is why diamond is one of the hardest materials known and melts only above 3,500°C. Silicon dioxide (quartz) is another network covalent solid.
Covalent Bonds in Biology
Your body runs on covalent chemistry. Proteins are long chains of amino acids, each linked to the next by a covalent peptide bond. These bonds are stable enough to maintain the chain’s structure under normal body conditions but can be broken by digestive enzymes when you need to recycle the amino acids. DNA’s backbone relies on covalent bonds between sugar molecules and phosphate groups, giving the double helix the structural integrity needed to store genetic information reliably through billions of cell divisions.
The covalent bonds within biological molecules are complemented by weaker forces (hydrogen bonds, for example) that handle the more flexible, reversible interactions, like the pairing of DNA strands or the folding of a protein into its working shape. But without the strong covalent scaffold underneath, none of those higher-level structures would hold together.
Drawing Lewis Structures
Lewis dot structures are the standard shorthand for representing covalent bonds on paper. Each line between two atoms represents one shared pair of electrons (one covalent bond), and dots represent lone pairs. To draw one, you follow a straightforward process: count up all the valence electrons from every atom in the molecule, arrange the atoms with the least electronegative one in the center, place single bonds between connected atoms, then distribute remaining electrons as lone pairs starting with the outer atoms. If the central atom still doesn’t have a full octet after that, convert lone pairs on neighboring atoms into double or triple bonds until it does.
For a molecule like CO₂, carbon is the central atom with 4 valence electrons, and each oxygen brings 6, totaling 16. After placing double bonds between carbon and each oxygen (using 8 electrons) and filling in lone pairs on the oxygens (using the remaining 8), every atom has a full octet. The structure shows two double bonds, which matches CO₂’s known properties: a strong, stable molecule with a linear shape.