A covalent bond is a chemical bond where two atoms share one or more pairs of electrons. Instead of one atom stealing an electron from another (as happens in ionic bonds), both atoms contribute electrons to a shared pair, and that shared pair holds the atoms together. This is the most common type of bond in organic chemistry and the reason molecules like water, oxygen, and DNA hold their shapes.
How Covalent Bonds Form
Every atom “wants” a full outer shell of electrons, which for most elements means eight (a principle chemists call the octet rule). Hydrogen is the exception, needing only two. When atoms can’t easily gain or lose electrons outright, they share instead.
The simplest example is hydrogen gas. Each hydrogen atom has one electron but needs two to fill its outer shell. When two hydrogen atoms approach each other, they each contribute their single electron to a shared pair. That shared pair now counts toward both atoms’ outer shells, making the system far more stable than two isolated hydrogen atoms. The result is a single molecule of H₂, held together by one covalent bond.
The same logic scales up. Carbon has four electrons in its outer shell and needs four more, which is why it forms four covalent bonds in almost every molecule it appears in. Oxygen needs two more, nitrogen needs three. These numbers explain the shapes of familiar molecules: water (H₂O) has two hydrogen atoms each sharing one electron with oxygen, methane (CH₄) has four hydrogens sharing with carbon, and ammonia (NH₃) has three hydrogens sharing with nitrogen.
Single, Double, and Triple Bonds
Atoms can share more than one pair of electrons. A single bond shares one pair, a double bond shares two pairs, and a triple bond shares three. Each additional shared pair pulls the atoms closer together and makes the bond stronger.
Carbon-carbon bonds illustrate this clearly. A single C–C bond is 154 picometers long and takes about 347 kilojoules of energy per mole to break. A double C=C bond shrinks to 134 picometers and requires 614 kJ/mol. A triple C≡C bond is shorter still at 120 picometers and needs 839 kJ/mol to break. More shared electrons means a tighter, stronger connection.
You encounter all three types in everyday chemistry. Ethane (in natural gas) has a single carbon-carbon bond. Ethylene, used to ripen fruit, has a double bond. Acetylene, the gas in welding torches, has a triple bond.
Polar vs. Nonpolar Covalent Bonds
Not all covalent bonds share electrons equally. When two identical atoms bond, like the two oxygens in O₂, they pull on the shared electrons with exactly the same strength. The electrons sit evenly between them. This is a nonpolar covalent bond.
When two different atoms bond, one usually pulls harder. Electronegativity is the term for how strongly an atom attracts shared electrons. A small difference in electronegativity between two atoms creates a polar covalent bond, where the electrons spend more time near the stronger-pulling atom. That atom picks up a slight negative charge, while the other end becomes slightly positive.
Water is the classic polar molecule. Oxygen pulls the shared electrons away from hydrogen, giving the oxygen end a partial negative charge and each hydrogen end a partial positive charge. This polarity is what makes water such an excellent solvent and gives it a relatively high boiling point for such a small molecule. When the electronegativity difference becomes very large, the bond stops being covalent altogether and becomes ionic, with one atom essentially taking the electron rather than sharing it.
Coordinate Covalent Bonds
In a standard covalent bond, each atom donates one electron to the shared pair. In a coordinate (or dative) covalent bond, one atom provides both electrons. The bond itself, once formed, is identical to any other covalent bond. The only difference is where the electrons came from.
This happens when one atom has a lone pair of electrons it can donate to another atom that has an empty spot in its outer shell. It’s common in metal complexes and in acids and bases reacting with each other.
How Covalent Compounds Behave
Substances held together by covalent bonds behave very differently from ionic compounds like table salt. Because covalent molecules are electrically neutral, the attraction between individual molecules is much weaker than the attraction between charged ions in a salt crystal. This has several practical consequences.
Covalent compounds generally have much lower melting and boiling points. Many are liquids or gases at room temperature: think of water, carbon dioxide, oxygen, and methane. In their solid form, they tend to be softer than ionic solids. They also conduct electricity poorly in any state, because there are no free charged particles to carry a current. Ionic compounds, by contrast, conduct electricity well when dissolved in water or melted, because their ions are free to move.
Visualizing Bonds With Lewis Structures
Chemists use Lewis structures (also called Lewis dot structures) to show how electrons are arranged in a molecule. Each dot represents a valence electron, and a line between two atoms represents a shared pair. To draw one, you add up all the valence electrons from every atom in the molecule, arrange the atoms with the least electronegative one in the center, and distribute electrons so that each atom reaches a full outer shell.
If the central atom still doesn’t have eight electrons after placing all the single bonds, lone pairs from neighboring atoms shift in to form double or triple bonds. Hydrogen always sits on the outside of the structure and only needs two electrons. The final check is simple: every valence electron should be accounted for, and every atom (except hydrogen) should have eight electrons around it.
Common Examples
Covalent bonds show up in an enormous range of substances. Simple inorganic molecules like hydrogen (H₂), nitrogen (N₂), chlorine (Cl₂), water (H₂O), and ammonia (NH₃) are all covalently bonded. Every organic compound, from the sugars you eat to the DNA in your cells, is built on a framework of covalent bonds, primarily between carbon, hydrogen, oxygen, and nitrogen.
Even the distinction between “covalent” and “ionic” isn’t always sharp. Many real-world bonds fall on a spectrum, with most being at least slightly polar. But the core idea stays the same: when atoms share electrons rather than transfer them, the bond holding them together is covalent.