Covalent Bond: Definition, Polar vs Nonpolar, Sigma and Pi Bonds

What is a covalent bond and how does it form?

A covalent bond holds two atoms together through a pair of electrons they share. Both nuclei attract that pair, and the attraction keeps the atoms close. Forming each covalent bond releases energy, which is why the bonded atoms are more stable than the separate ones. Covalent bonds are typical of very electronegative atoms, the nonmetals.

Hydrogen is the simplest case. Each atom has one electron; by sharing them, both end up with two, the configuration of helium, and the H₂ molecule stops at a single bond.

Valence bond theory describes this as the overlap of two atomic orbitals, one from each atom, each holding one electron. In the overlap region the two electrons pair up, and the more the orbitals overlap, the stronger the bond. As two atoms approach, their energy falls to a minimum, and the distance between the nuclei at that point is the bond length.

Octet rule, bonding pairs and Lewis structures

In most simple molecules, shared electrons give each atom the configuration of a noble gas: eight electrons in the outer shell, two for hydrogen. This is Lewis’s octet rule. Chlorine has seven valence electrons, one short of an octet, and in Cl₂ it fills the shell by sharing a pair.

In a Lewis drawing the shared pair is the bonding pair, and the pairs that belong to a single atom are lone pairs. In water the oxygen has two bonding pairs and two lone pairs, and each hydrogen, with its shared pair, has a full shell of two electrons. When drawing a more complex structure, the central atom is usually the least electronegative element in the compound.

Practical example: to draw Cl₂, start from the seven valence electrons of each atom. One pair becomes the line between the two symbols, the bonding pair, and the remaining electrons stay on the atoms as lone pairs.

The rule has exceptions. In boron trifluoride the boron has only 3 outer electrons and cannot reach eight by sharing. In phosphorus pentachloride the phosphorus has 5 pairs in its outer shell, so more than eight electrons.

Polar vs nonpolar covalent bond

To see how the pair is divided you need electronegativity, the relative ability of an atom to attract bonding electrons. On the Pauling scale, the most widely used, fluorine is set at 4.0 by convention; carbon, nitrogen and oxygen are about 2.5, 3.0 and 3.5. The value grows from left to right across a period of the periodic table and falls going down a group. Noble gases are left off the scale because they do not usually share electrons.

Nonpolar covalent bond

When the pair is shared in perfectly equal parts, the bond is nonpolar. This normally happens between atoms of the same element, as in H₂ and Cl₂.

Polar covalent bond

When one atom pulls the pair harder, the sharing is unequal. The more electronegative atom takes a partial negative charge and the less electronegative one a partial positive charge, both smaller than a whole charge. The H-Cl bond of hydrogen chloride and the O-H bonds of water are the standard examples. The bigger the electronegativity difference, the larger the partial charges. The C-H bond has a small difference and is usually treated as nonpolar.

From covalent to ionic: a continuum

There is no sharp border between covalent and ionic bonds. A difference of zero gives a nonpolar covalent bond, a small one gives a polar covalent bond, and a large one gives an ionic bond. One teaching source points out that there is no exact answer to where polar ends and ionic begins: lithium iodide, for example, is described as ionic with some covalent character, and it dissolves in organic solvents such as ethanol.

The numbers in textbooks are conventions. A course based on CK-12 material treats a bond with a difference below 0.4 as nonpolar, between 0.4 and 1.7 as polar covalent, and above 1.7 no longer as mostly covalent. An Italian teaching site puts the upper limit of polar at 1.9, so the two cutoffs differ.

The ionic bond and the other bond types are compared in the Recap on chemical bonds.

Single, double and triple covalent bonds

A double bond is two electron pairs shared between the same two atoms, instead of one. O₂ has a double bond: sharing two pairs gives both oxygen atoms a stable structure. N₂ has a triple bond, since each nitrogen has 5 outer electrons and shares 3. In the Lewis model the number of pairs holding two atoms together is the bond order: 1, 2 or 3.

For comparable bonds, length falls as bond order rises, and bond energy rises. The relation is not linear: a double bond is not half as long as a single one.

Molecule	Bond	Order	Length (pm)	Dissociation energy (kJ/mol)
Ethane	C-C	1	153.5	376
Ethene	C=C	2	133.9	728
Ethyne	C≡C	3	120.3	965
Oxygen O₂	O=O	2	120.7	498.4
Nitrogen N₂	N≡N	3	109.8	945.3

The values come from a Chemistry LibreTexts table for gas-phase molecules at 298 K, drawing on the CRC Handbook (2004) and NIST data. Other tables, built on different conditions, give slightly different figures for the same bonds.

Practical example: breaking the C-C bond of ethane takes about 376 kJ/mol, and the C=C bond of ethene takes about 728. The C=C bond measures 133.9 pm, against 153.5 for C-C.

Among single bonds, N-N, O-O and F-F are unusually weak compared with those of the heavier elements in the same group. The explanation offered, with caution, is strong repulsion between the lone pairs on neighbouring atoms.

Sigma and pi bonds

A sigma bond comes from head-on overlap of two atomic orbitals. In H₂ these are the two 1s orbitals, and the bond has cylindrical symmetry around the line joining the nuclei. A pi bond comes from sideways overlap of two p orbitals. Its electrons are off that line, above and below the plane of the molecule. In ethene the pi bond between the two carbons is weaker than the sigma bond, because it lies a little farther from the control of the nuclei.

Every double bond, between any pair of atoms, is built from one sigma bond and one pi bond. The triple bond of ethyne has two pi bonds at right angles to each other: one above and below the molecule, the other in front and behind.

Dative (coordinate) covalent bond

Usually each atom contributes one electron to the pair, though it can go otherwise. When both electrons come from the same atom, the bond is called dative, or coordinate. It links a donor atom, which has a lone pair, to an acceptor atom, which needs a pair to complete a shell.

In diagrams an arrow points from donor to acceptor. In the ammonium ion, once formed, it is indistinguishable from the three ordinary covalent bonds.

Practical example: in boron trifluoride the boron is short of electrons. The lone pair on the nitrogen of ammonia fills the gap, and the compound that forms has a dative bond.

Covalent bond examples

Substance	Bond	Type
H₂, Cl₂	single	nonpolar, atoms of the same element
O₂	double	nonpolar
N₂	triple	nonpolar
H-Cl	single	polar, chlorine pulls the pair
H₂O	two O-H bonds	polar, polar molecule
CO₂	two C-O bonds	polar, nonpolar molecule
Diamond	single, through the whole structure	giant covalent network

Polar bonds and polar molecules

A molecule’s dipole moment is the vector sum of the dipoles of its bonds. In bigger molecules, three-dimensional geometry and symmetry decide the outcome.

Practical example: in CO₂ the two C-O bonds are polar, but the molecule is linear. The dipoles are equal and point 180 degrees apart, so they cancel, and experiments find no dipole moment. Water is bent, with an angle of 104.5° between its two O-H dipoles, so they do not cancel.

In highly symmetric geometries (tetrahedral, square planar, trigonal bipyramidal, octahedral) the bond dipoles cancel completely, as in CCl₄ and BCl₃, while NH₃ and CHCl₃ have a net dipole. The polarity of water lets it form hydrogen bonds with other polar or charged species, including other water molecules.

Giant covalent structures

Diamond is a covalent substance with no separate molecules. Each carbon shares electrons with four other carbons through four single bonds, in a structure that extends through all three dimensions. It is not a molecule, since the number of atoms joined depends on the size of the crystal.

Subliming diamond means breaking the very strong carbon-carbon bonds of the whole structure. The electrons are held tight between the atoms and cannot move, which is why diamond does not conduct electricity. In graphite each carbon forms single bonds with three neighbours, and the fourth electron is delocalised in the plane.