7.2 Covalent Bonding

Covalent bonding describes how nonmetal atoms share electrons to achieve stable electron configurations. Understanding this type of bond is key to predicting molecular properties like melting point, solubility, and conductivity.

Covalent Bonding

Formation of Covalent Bonds

Covalent bonds form when two atoms share one or more pairs of electrons. Each atom contributes electrons to the shared pair, and both nuclei are attracted to those shared electrons. That mutual attraction is what holds the molecule together.

• Atoms form covalent bonds to achieve a stable octet (8 electrons in their outer shell). Hydrogen is an exception: it only needs 2 electrons to fill its valence shell.
• Covalent bonds typically form between nonmetals like carbon, oxygen, and nitrogen. These elements have high electronegativities, so they tend to attract and share electrons rather than give them up entirely.
• Bond energy is the amount of energy required to break a covalent bond. Higher bond energy means a stronger, more stable bond.

Atoms can share more than one pair of electrons:

• A single bond shares 1 pair of electrons (e.g., $\ce{H2}$)
• A double bond shares 2 pairs (e.g., $\ce{O2}$)
• A triple bond shares 3 pairs (e.g., $\ce{N2}$)

Triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds.

Polarity in Covalent Bonds

Not all covalent bonds share electrons equally. The key factor is electronegativity, which measures how strongly an atom attracts shared electrons. Fluorine is the most electronegative element (4.0 on the Pauling scale), while metals on the left side of the periodic table have the lowest values.

The electronegativity difference between two bonded atoms determines whether the bond is polar or nonpolar:

• Nonpolar covalent bonds form when the electronegativity difference is very small (roughly 0 to 0.4). The electrons are shared equally. This happens when identical atoms bond together, like in $\ce{H2}$, $\ce{Cl2}$, or $\ce{O2}$.
• Polar covalent bonds form when there is a moderate electronegativity difference (roughly 0.5 to 1.7). The electrons are pulled closer to the more electronegative atom, creating an unequal distribution of charge.

In a polar bond, the more electronegative atom develops a partial negative charge ($\delta-$), and the less electronegative atom develops a partial positive charge ($\delta+$). For example, in $\ce{HCl}$, chlorine is more electronegative than hydrogen, so the shared electrons spend more time near chlorine. Water ($\ce{H2O}$) and ammonia ($\ce{NH3}$) also contain polar covalent bonds.

Covalent vs. Ionic Compounds

Because covalent and ionic compounds are held together by different types of forces, they behave very differently. Here's how they compare:

• Melting and boiling points: Covalent compounds generally have much lower melting and boiling points. The forces between covalent molecules (intermolecular forces) are weaker than the strong electrostatic attractions between ions in an ionic compound. For example, table sugar (a covalent compound) melts at about 186°C, while table salt (ionic) melts at 801°C.
• Electrical conductivity: Covalent compounds are typically poor conductors of electricity in any state because they lack free-moving charged particles. Plastics and oils are good examples. Ionic compounds don't conduct as solids either, but they do conduct when melted or dissolved in water because their ions become free to move and carry charge.
• Solubility: A useful rule is "like dissolves like." Most covalent compounds are nonpolar, so they dissolve well in nonpolar organic solvents (like oil dissolving in hexane) but poorly in water. Ionic compounds are the opposite: they tend to dissolve in polar solvents like water (think of salt dissolving easily) but not in organic solvents.

Molecular Structure and Representation

The way atoms are arranged in a molecule determines its shape, polarity, and reactivity.

• Lewis structures are diagrams that show how valence electrons are distributed in a molecule. They display both bonding pairs (shared between atoms) and lone pairs (not shared). Drawing Lewis structures is the first step in predicting molecular shape.
• Molecular geometry describes the three-dimensional arrangement of atoms in a molecule. The shape is determined by both bonding pairs and lone pairs around the central atom, following VSEPR (Valence Shell Electron Pair Repulsion) theory. Electron pairs repel each other and spread out as far as possible, which dictates shapes like linear, bent, trigonal planar, and tetrahedral.
• Hybridization is the concept of atomic orbitals mixing to form new hybrid orbitals suited for bonding. For example, carbon in methane ($\ce{CH4}$) undergoes $sp^3$ hybridization, producing four equivalent orbitals arranged in a tetrahedral shape. The type of hybridization ($sp$, $sp^2$, $sp^3$) determines the geometry and bond angles of the molecule.