Covalent Bonding
Imagine two roommates splitting the rent on an apartment. Neither can afford the place alone, so they pool their resources and share the cost. Both benefit, and the arrangement only works because each person contributes. Covalent bonding follows the same logic - two atoms share electrons because neither one can simply take them from the other. This is the dominant bonding strategy between nonmetal atoms, which have similar (and relatively high) electronegativities.
Why Sharing Instead of Transferring?
In ionic bonding, one atom clearly dominates: the metal has such a low ionization energy that it readily gives up electrons. In covalent bonding, neither atom is willing to surrender electrons because both have relatively high electronegativities. The compromise is sharing.
When two nonmetals approach each other, their atomic orbitals overlap and the shared electrons spend time in the region between both nuclei. This concentration of negative charge between the two positive nuclei acts as “electrostatic glue,” holding the atoms together.
Single, Double, and Triple Bonds
Atoms can share more than one pair of electrons. The number of shared pairs determines the bond type:
Ionic to covalent is one axis, not two boxes
Scroll sideways to see the whole map.
Why the cut-offs are dashed1.7 is a convention, not a physical boundary, and different textbooks put it at 1.7 or 1.8. HF sits at 1.78 and is unmistakably a covalent molecular gas, which is exactly the point: nothing changes state at the line. Use the number to rank bonds, not to classify them absolutely.
What a dipole actually isA polar bond has a δ+ and a δ− end. A polar molecule needs those bond dipoles not to cancel: CO₂ has two strongly polar bonds and no net dipole because they point opposite ways, while water's two bonds are bent apart and add up.
Where the lattice comes fromIonic compounds are not molecules. Once electrons are transferred, every cation attracts every nearby anion, so the ions stack into a repeating lattice held by electrostatic attraction in all directions. That is why they are hard, brittle, high-melting, and conduct only once melted or dissolved.
- Single bond: one shared pair (2 electrons). Example: H-H in H₂, C-H in methane.
- Double bond: two shared pairs (4 electrons). Example: O=O in O₂, C=O in CO₂.
- Triple bond: three shared pairs (6 electrons). Example: N≡N in N₂, C≡C in acetylene.
More shared pairs mean a stronger, shorter bond. A triple bond is stronger and shorter than a double bond, which is stronger and shorter than a single bond between the same two atoms. This relationship between bond order, bond strength, and bond length is tested repeatedly on the MCAT.
Bond Order
Bond order is simply the number of bonding electron pairs shared between two atoms. For a single bond, bond order = 1. For a double bond, bond order = 2. For a triple bond, bond order = 3.
In molecules with resonance structures, bond order can be a non-integer. For example, each C-O bond in the carbonate ion (CO₃²-) has a bond order of (approximately 1.33), because four bonding pairs are distributed across three equivalent C-O bonds.
Nonpolar Covalent Bonds
When two identical atoms form a covalent bond, they share electrons perfectly equally. There is zero electronegativity difference, so neither atom pulls the shared electrons toward itself. These are nonpolar covalent bonds.
Examples include H₂, O₂, N₂, Cl₂, and F₂ - any diatomic molecule made from the same element. The electron density is distributed symmetrically between the two nuclei.
This is one end of the bonding spectrum. As the electronegativity difference between the two atoms increases, the bond becomes polar covalent, and eventually ionic. We will explore the polar covalent region in the next section.
Coordinate Covalent Bonds
In a standard covalent bond, each atom contributes one electron to the shared pair. In a coordinate covalent bond (also called a dative bond), both electrons in the shared pair come from the same atom.
The atom that donates both electrons is called the Lewis base (it has a lone pair to give). The atom that accepts the electron pair is the Lewis acid (it has an empty orbital to receive).
Classic examples:
- NH₄+ (ammonium): Ammonia (NH₃) has a lone pair on nitrogen. When H+ approaches, nitrogen donates both electrons to form the fourth N-H bond. All four N-H bonds in NH₄+ are identical - you cannot tell which one was the coordinate covalent bond.
- H₃O+ (hydronium): Water has two lone pairs on oxygen. One lone pair is donated to H+ to form the third O-H bond.
The critical MCAT takeaway: once a coordinate covalent bond forms, it is indistinguishable from a regular covalent bond. The distinction only matters when describing how the bond was formed, not how it behaves.
Properties of Covalent Compounds
Covalent compounds (also called molecular compounds) have properties that differ sharply from ionic compounds because they exist as discrete molecules held together by relatively weak intermolecular forces, rather than as a lattice of ions held by strong electrostatic forces.
Lower melting and boiling points. To melt or boil a covalent compound, you only need to overcome the intermolecular forces between molecules - not break covalent bonds. These forces (London dispersion, dipole-dipole, hydrogen bonds) are much weaker than the ion-ion attractions in a crystal lattice.
Poor electrical conductors. Covalent compounds have no free ions or delocalized electrons to carry charge. Even when dissolved, they typically do not produce ions (with notable exceptions like acids).
Variable solubility. Polar covalent compounds tend to dissolve in polar solvents (like water). Nonpolar covalent compounds tend to dissolve in nonpolar solvents (like hexane). “Like dissolves like” is the guiding principle.
Ionic vs. Covalent: A Side-by-Side Comparison
| Property | Ionic Compounds | Covalent Compounds |
|----------|----------------|-------------------|
| Formation | Electron transfer (metal + nonmetal) | Electron sharing (nonmetal + nonmetal) |
| Basic unit | Formula unit (crystal lattice) | Discrete molecule |
| Melting/boiling point | High (strong lattice forces) | Low to moderate (weak intermolecular forces) |
| Electrical conductivity | Conducts when dissolved or molten | Generally does not conduct |
| Solubility | Soluble in polar solvents | Polar in polar solvents, nonpolar in nonpolar solvents |
| State at room temperature | Usually solid | May be solid, liquid, or gas |
| Hardness | Hard but brittle | Soft or waxy (many exceptions) |
Network Covalent Solids: The Exception
Not all covalent compounds have low melting points. Network covalent solids like diamond (C), silicon dioxide (SiO₂), and silicon carbide (SiC) consist of atoms connected by covalent bonds in a continuous three-dimensional network - no discrete molecules at all. These materials have extremely high melting points, often exceeding those of ionic compounds.
Diamond melts at roughly 3,550 degrees C, far above NaCl’s 801 degrees C. The difference is that breaking a network covalent solid requires breaking actual covalent bonds, not just overcoming intermolecular forces.
If a passage describes a covalent compound with an unusually high melting point, think network covalent solid.
A coordinate covalent bond is formed when one atom donates both electrons in the shared pair. However, once formed, it is completely indistinguishable from a regular covalent bond. The distinction only describes the bond’s origin, not its properties. Examples include the fourth N-H bond in NH₄+ and the third O-H bond in H₃O+.
Increasing bond strength: C-C < C=C < C≡C. Decreasing bond length: C-C > C=C > C≡C. As bond order increases, more shared electron pairs pull the nuclei closer together (shorter bond) and require more energy to break (stronger bond). A triple bond is the shortest and strongest; a single bond is the longest and weakest.
A network covalent solid. Unlike typical molecular covalent compounds (which have low melting points because only weak intermolecular forces must be overcome), network covalent solids like diamond (C) and quartz (SiO₂) have continuous three-dimensional covalent bonding. Melting requires breaking strong covalent bonds throughout the entire structure, producing extremely high melting points.