Chemical Bonds: Ionic, Covalent and Metallic Bond Types

Two atoms bond when the system they form together has lower potential energy than the separate atoms. The underlying mechanism is electrostatic: opposite charges attract, so nuclei and electrons of neighboring atoms pull the system into a steadier state. For sodium and chlorine, for example, a solid of ions has less energy than the same atoms scattered as a gas.

Chemical bonding definition: why atoms bond

The electrons that take part are the ones in the outer shell, the valence electrons. A widely used explanation puts the stability down to the electron configuration of the noble gases, and it is the basis of the octet rule: atoms tend to reach eight valence electrons. LibreTexts presents it as an interpretation.

The rule has three kinds of exception. In some molecules the central atom has fewer than eight electrons, as in BeH2 and BCl3. Others have an odd number of valence electrons: NO2 has 17. And from the third period onward atoms can hold more than eight electrons, as in PCl5 and SF6.

Types of chemical bonds: ionic, covalent and metallic

Textbooks distinguish three types: in an ionic bond electrons pass from one atom to another, in a covalent bond they are shared in pairs, and in a metallic bond the valence electrons are mobile throughout the solid.

Ionic bond and examples of ionic compounds

An ionic bond is the attraction between positive and negative ions. When sodium reacts with chlorine, each chlorine atom takes an electron from a sodium atom and Na+ and Cl- form. After the transfer each ion has eight electrons in its valence shell. Bringing oppositely charged ions together is strongly favorable in energy terms, and the energy released as the lattice forms stabilizes the solid.

Practical example: table salt, sodium chloride, is a lattice of Na+ and Cl-. Salt water conducts electric current, while a dry salt crystal does not.

In ionic crystals each ion is attracted to many oppositely charged ions at once, which makes the lattice very strong: NaCl melts at about 800 °C. Molten or dissolved, ionic compounds conduct current, because free ions can move and carry it. In the solid they are locked in the lattice. The crystals are generally hard but brittle. If one is struck, layers slide, ions of the same charge find themselves side by side, and the repulsion shatters the crystal.

Covalent bond: shared pairs of electrons

In a covalent bond the outer electrons are pulled by the nuclei of both atoms at the same time, and a shared pair creates a mutual attraction. As a rule each atom contributes one of its valence electrons. The hydrogen molecule, H2, is the simplest model: with the shared pair, each atom reaches two valence electrons. One shared pair is a single bond, two pairs a double bond, three pairs a triple bond. Many molecules complete their octet by sharing more than one pair. Nitrogen, N2, is rather inert because of the strong triple bond between its two atoms. Sometimes the whole pair comes from a single atom. That is called a dative, or coordinate, bond, and the ammonium ion, NH4+, is an example: the fourth hydrogen attaches using the lone pair on the nitrogen of ammonia.

Single, double and triple bonds: sigma and pi bonds

Sigma and pi bonds describe how these bonds are built. A sigma bond comes from head-on overlap of orbitals, with the electron density concentrated between the two nuclei. A pi bond comes from sideways overlap, with the electron density on either side of the plane that holds the nuclei.

Every single bond is a sigma bond. A double bond is one sigma and one pi, and a triple bond is one sigma and two pi.

Bond order also changes length and energy. The Howard University textbook gives these values for the bond between two carbon atoms (the energies are dissociation energies, measured in gas-phase molecules at 298 K):

C-C bond	Example	Distance	Dissociation energy
Single	ethane	153.5 pm	376 kJ/mol
Double	ethene	133.9 pm	728 kJ/mol
Triple	ethyne	120.3 pm	965 kJ/mol

The higher the bond order, the closer the nuclei and the more energy it takes to break the bond. Energy does not rise in proportion, though.

Polar and nonpolar bonds

Electronegativity is the relative ability of an atom in a bond to attract electrons. There are several scales, all of them conventions; the best known is the Pauling scale, which runs from 0 to 4.

Between identical atoms, as in H2, O2 or N2, the electronegativity difference is zero and the pair is shared equally. The bond is nonpolar. Between different atoms the pair shifts toward the more electronegative one. The bond is then polar and has partial ionic character, which can be expressed as a percentage.

A polar bond and a polar molecule are different things. A molecule needs at least one polar bond to be polar, but that is not sufficient. CO2 has two polar C=O bonds, yet they are dipoles of equal strength pointing in opposite directions: they cancel, and the molecule is nonpolar. Water’s O-H dipoles do not cancel, which makes that molecule polar. A polar molecule behaves as an electric dipole, and its polarity is measured by the permanent electric dipole moment. “Polar” is also often a relative term: one molecule can be more or less polar than another.

Metallic bond: definition and example

The metallic bond model says the valence electrons are free, delocalized and mobile, with no particular atom as their owner. A metal is sometimes described as a lattice of positive ions in a sea of electrons, a simplified picture. Because the electrons move freely, metals conduct electric current.

The electron sea also acts as a cushion: when a metal is hammered, the atoms slide past one another and the bond holds. That accounts for malleability, the ability to be shaped into thin sheets such as aluminum foil, and for ductility, which allows the metal to be drawn into wires such as copper wire.

Telling the bond type from electronegativity

To judge which type a bond between two different atoms resembles, look at the electronegativity difference. An introductory Fresno City College textbook, on the Pauling scale, uses these cut-offs:

Electronegativity difference	Most likely bond type
0 to 0.4	pure covalent (nonpolar)
0.5 to 2.0	polar covalent
above 2.0	probably ionic

The cut-offs vary from textbook to textbook and are only indicative. Real substances often mix the characters of the three types in varying degrees, and even within covalent bonding the polar kind has a share of ionic character.

Practical example: in hydrogen, H2, the electronegativity difference is zero and the bond is pure covalent. In salt, NaCl, the bond is ionic. In water the O-H bonds are polar covalent.

Covalent vs non-covalent bonds: the weak forces between molecules

The attractions between molecules are usually far weaker than a covalent bond. All molecules, polar or nonpolar, attract one another through London dispersion forces, which for similar substances grow with the size of the molecule. Dipole-dipole interactions are much weaker than those between two ions.

A hydrogen bond is a particularly strong dipole-dipole interaction between molecules, formed when hydrogen is attached to a highly electronegative atom such as O, N or F. According to Fresno City College, its typical dissociation energies are only 15 to 25 kJ/mol, far less than a covalent bond. Hydrogen bonding between its molecules gives water an unusually high boiling point.