When a nonmetal bonds with a nonmetal, the atoms share electrons rather than transferring them. This type of bond is called a covalent bond, and it’s fundamentally different from what happens when a metal bonds with a nonmetal (which produces ionic bonds through electron transfer). The shared electrons hold both atoms together because both nuclei are attracted to the same pair of electrons, creating a stable arrangement.
Why Nonmetals Share Instead of Transfer
Atoms bond to reach a more stable electron arrangement, typically eight electrons in their outer shell (the “octet rule”). Metals achieve this by giving away electrons, but nonmetals need to gain electrons. When two nonmetals meet, neither is willing to surrender its electrons to the other. The compromise: they share. Each atom contributes one electron to form a shared pair, and both atoms count that pair as part of their outer shell.
The number of bonds a nonmetal forms depends on how many electrons it needs to complete its octet. Carbon has four unpaired outer electrons and forms four bonds. Nitrogen has three unpaired electrons and typically forms three. Oxygen needs two more electrons and forms two bonds. Halogens like chlorine and fluorine need just one more electron, so they usually form a single bond.
Single, Double, and Triple Bonds
Atoms can share more than one pair of electrons. When two atoms share one pair, that’s a single bond. Sharing two pairs creates a double bond, and sharing three pairs creates a triple bond. Oxygen gas (O₂) has a double bond between its two oxygen atoms because each oxygen needs two electrons. Nitrogen gas (N₂) has a triple bond because each nitrogen needs three.
More shared pairs means a stronger, shorter bond. A carbon-carbon single bond has a bond energy of about 347 kJ/mol, meaning that’s how much energy it takes to break it apart. A carbon-hydrogen bond is stronger at 413 kJ/mol, and an oxygen-hydrogen bond is stronger still at 467 kJ/mol. These differences in bond strength help explain why some molecules are more stable or reactive than others.
Polar vs. Nonpolar Covalent Bonds
Not all sharing is equal. When two identical atoms bond (like two oxygen atoms in O₂), they pull on the shared electrons equally, creating a nonpolar covalent bond. But when two different nonmetals bond, one atom usually pulls the electrons closer to itself. This uneven sharing creates a polar covalent bond, where one end of the bond is slightly negative and the other slightly positive.
How uneven the sharing is depends on the difference in electronegativity between the two atoms. Electronegativity is a measure of how strongly an atom attracts shared electrons. Fluorine is the most electronegative element at 4.0 on the Pauling scale. Oxygen comes in at 3.5, while chlorine and nitrogen both sit at 3.0. The bigger the gap between two bonded atoms, the more polar the bond:
- 0.0 to 0.4 difference: nonpolar covalent (like C-H or C-C bonds)
- 0.5 to 0.9: slightly polar (like H-Cl)
- 1.0 to 1.3: moderately polar (like C-O)
- 1.4 to 1.7: highly polar (like H-O in water)
Once the electronegativity difference exceeds about 1.8, the bond starts behaving more like an ionic bond, with electrons essentially transferred rather than shared. The hydrogen-fluorine bond, with a difference of 1.9, sits in this borderline zone.
Properties of Nonmetal-Nonmetal Compounds
Compounds formed from covalent bonds between nonmetals behave very differently from ionic compounds (the kind formed between metals and nonmetals). Ionic compounds like table salt are always solid at room temperature, conduct electricity when dissolved in water, and have high melting points. Covalent compounds are far more varied.
Nonmetal-nonmetal compounds can be gases, liquids, or solids at room temperature. Water (H₂O) is a liquid, carbon dioxide (CO₂) is a gas, and sugar (C₁₂H₂₂O₁₁) is a solid. They generally have lower melting and boiling points than ionic compounds because the forces between individual molecules are weaker than the forces holding an ionic crystal together. They also conduct electricity poorly in any state, since they don’t contain free-moving charged particles.
How These Compounds Are Named
Because nonmetals can combine in multiple ratios (carbon and oxygen can form both CO and CO₂, for instance), naming these compounds uses Greek prefixes to specify exactly how many of each atom are present. The prefixes run from mono (1) through di (2), tri (3), tetra (4), penta (5), hexa (6), and so on up to deca (10).
The first element in the name never gets the “mono” prefix. If no prefix appears before it, you assume there’s one atom. The second element always gets a prefix and ends in “-ide.” So CO is carbon monoxide (one carbon, one oxygen), CO₂ is carbon dioxide, and N₂O₅ is dinitrogen pentoxide. This system is specific to compounds made of two nonmetals. Ionic compounds use a different naming system entirely.
Molecular Shape and Electron Repulsion
Covalent molecules aren’t flat stick figures. They have three-dimensional shapes determined by how electron pairs around the central atom push away from each other. Both bonding pairs (shared between atoms) and lone pairs (electrons not involved in bonding) create regions of negative charge that repel one another, spreading out as far apart as possible.
This principle explains why water is bent rather than straight. Oxygen has two bonding pairs (connecting to hydrogen atoms) and two lone pairs. Those four regions of electron density push into a roughly tetrahedral arrangement, but since two of those regions are invisible lone pairs rather than bonds, the visible shape is a bent “V.” Similarly, ammonia (NH₃) has three bonds and one lone pair, giving it a pyramidal shape rather than a flat triangle. These shapes directly affect a molecule’s polarity, boiling point, and how it interacts with other molecules.
Exceptions to the Octet Rule
The octet rule works well for most nonmetal-nonmetal compounds, but several important molecules break it. These exceptions fall into three categories.
Some molecules have an odd number of electrons, making it impossible for every atom to have a complete octet. Nitric oxide (NO) is a classic example. Other molecules have atoms with fewer than eight electrons. Boron trichloride (BCl₃), used in manufacturing reinforced sports equipment like tennis rackets, has a boron atom surrounded by only six valence electrons.
Elements in the third row of the periodic table and beyond can also exceed the octet, holding more than eight electrons. Sulfur hexafluoride (SF₆), used by the electric power industry to insulate high-voltage equipment, has sulfur bonded to six fluorine atoms, giving it twelve electrons in its outer shell. Phosphorus pentafluoride (PF₅) is another example. These expanded octets are possible because larger atoms have access to additional orbital space that smaller atoms like carbon and nitrogen lack.
Coordinate Covalent Bonds
In a standard covalent bond, each atom contributes one electron to the shared pair. But in a coordinate (or dative) covalent bond, both electrons come from the same atom. Once formed, a coordinate bond is identical in strength and character to any other covalent bond. The distinction only matters when you’re tracking where the electrons originally came from.
A common example is the formation of the hydronium ion (H₃O⁺) when an acid dissolves in water. A hydrogen ion, which is just a bare proton with no electrons, attaches to a lone pair on the oxygen atom of a water molecule. Oxygen supplies both electrons for that new bond. Carbon monoxide also features a coordinate bond: in addition to two standard covalent bonds between carbon and oxygen, a lone pair from oxygen forms a third bond, giving CO an unusually strong triple bond.