Covalent Bonding

Basic Idea of Covalent Bonding

In a covalent bond, two (or more) atoms share electron pairs. Each shared pair is located between the bonding atoms and belongs to both of them at the same time.

Covalent bonding is especially important for:

• Nonmetals bonding with other nonmetals (e.g. $H_2$, $O_2$, $H_2O$, $CH_4$)
• Many compounds in living organisms (sugars, proteins, DNA, etc.)

You can think of a covalent bond as two atoms “holding on” to the same pair of electrons to reach a more stable electron configuration.

Common features:

• The shared electrons are usually found in the region between the nuclei.
• The bond has a definite direction: it connects specific atoms in specific orientations.
• The number of shared pairs determines whether the bond is single, double, or triple.

Electron Pairs and the Octet Rule (Simple View)

For many main-group elements, a simple picture works well:

• Atoms “want” to achieve 8 valence electrons (an octet) in their outer shell (hydrogen is an important exception; it aims for 2).
• By sharing pairs of electrons, each atom in the bond “counts” the bonding pair as part of its own valence shell.

Example: molecule $H_2O$

• Oxygen has 6 valence electrons.
• It forms 2 covalent bonds with 2 hydrogen atoms.
• Each O–H bond is a shared pair: oxygen “counts” 2 extra electrons, reaching 8.
• Each hydrogen “counts” the bonding pair as its 2 electrons.

This simple counting rule works well for many common covalent molecules (e.g. $CH_4$, $NH_3$, $H_2O$, $CO_2$), but there are exceptions and more advanced descriptions that are dealt with elsewhere.

Single, Double, and Triple Covalent Bonds

Single Bonds

A single covalent bond is formed by one shared electron pair (2 electrons in total).

Example: hydrogen molecule $H_2$

• Each hydrogen atom has 1 valence electron.
• Together they share 2 electrons:
• Each hydrogen “sees” 2 electrons in its valence shell.
• The two atoms are held together by one single bond, often drawn as:
• Structural formula: $H-H$

Properties of single bonds (in simple terms):

• Typically longer than double or triple bonds between the same pair of elements.
• Typically weaker than double or triple bonds between the same atoms.

Double Bonds

A double bond consists of two shared electron pairs (4 electrons).

Example: oxygen molecule $O_2$

• Each oxygen atom has 6 valence electrons.
• They share 2 pairs (4 electrons) in a double bond:
• Structural formula: $O=O$
• Each oxygen “counts” 4 bonding electrons plus its own nonbonding electrons to reach an octet.

Compared to a single bond between the same atoms:

• A double bond is shorter.
• A double bond is stronger.

Triple Bonds

A triple bond consists of three shared electron pairs (6 electrons).

Example: nitrogen molecule $N_2$

• Each nitrogen has 5 valence electrons.
• They share 3 pairs of electrons:
• Structural formula: $N\equiv N$
• Each nitrogen “counts” 3 shared pairs plus its remaining nonbonding electrons, reaching an octet.

Triple bonds:

• Are shorter and stronger than single and double bonds between the same atoms.
• Often make molecules relatively rigid and chemically less flexible at that bond.

Bond Length and Bond Strength in Covalent Bonds

Two simple but important relationships:

• Bond length: the average distance between the nuclei of two bonded atoms.
• Bond strength: often measured as bond dissociation energy – the energy required to break the bond between two atoms in a molecule.

Typical trends for bonds between the same two elements:

• As the number of shared pairs increases (single → double → triple):
• Bond length decreases.
• Bond strength increases.

Example (qualitative):

• C–C single bond: relatively long, relatively weak (compared with C=C, C≡C).
• C=C double bond: shorter and stronger.
• C≡C triple bond: shortest and strongest among the three.

These trends strongly influence how molecules behave in chemical reactions, because breaking strong, short bonds requires more energy.

Polar and Nonpolar Covalent Bonds (Qualitative View)

In many covalent bonds, the shared electron pair is not shared equally.

• If two atoms have identical or very similar abilities to attract electrons, the bond is approximately:
• Nonpolar covalent: electrons are shared fairly evenly.
• Example: $H_2$, $Cl_2$, $O_2$, or C–H (approximately).
• If one atom attracts electrons more strongly than the other, the bond is:
• Polar covalent: the shared electrons are nearer to one atom on average.
• This gives:
• A partial negative charge $(\delta^-)$ on the more electron-attracting atom.
• A partial positive charge $(\delta^+)$ on the other atom.

Example: $HCl$

• Chlorine attracts electrons more strongly than hydrogen.
• The bond is polar covalent:
• $H^{\delta+}-Cl^{\delta-}$

These partial charges create bond dipoles, which are important for:

• Determining whether a whole molecule is polar or nonpolar.
• Influencing boiling points, solubilities, and intermolecular interactions.

The quantitative description of how strongly an atom attracts electrons (electronegativity) and how to compare bond types belongs to other chapters; here the focus is on the idea that covalent bonds can be either nonpolar or polar.

Localized and Delocalized Covalent Bonding (Conceptual)

For many simple molecules (e.g. $H_2O$, $CH_4$), we can picture covalent bonds as localized:

• Each pair of atoms has its own distinct shared electron pair directly between them.
• We draw these as a single line, double line, etc.

However, in some molecules and ions, electrons are spread out over more than two atoms. This is called delocalized covalent bonding.

Resonance Structures (Qualitative Idea)

Sometimes, no single Lewis structure (with localized bonds) can accurately describe the bonding. For example:

• In certain molecules and ions, two or more valid drawings (same arrangement of atoms, different placement of electrons) can be written.
• These are called resonance structures (or resonance forms).
• The real molecule is best thought of as a hybrid of these structures.
• In such cases, some bonds have an intermediate character between single and double, etc., reflecting delocalization of electrons.

Example (no full details here):

• In some polyatomic ions or aromatic compounds, all bonds between certain atoms are equivalent in reality, even though simple drawings suggest single and double bonds.

Delocalized covalent bonding has important consequences for:

• Bond lengths (often intermediate between typical single and double bonds).
• Chemical stability.
• Reactivity and color in some organic and inorganic compounds.

A more detailed treatment of resonance, aromaticity, and advanced bonding models is given elsewhere.

Coordinate (Dative) Covalent Bonds

In most covalent bonds, each atom contributes one electron to the shared pair.
In a coordinate (dative) covalent bond, both electrons in the shared pair come from the same atom.

Despite this difference in origin:

• Once formed, a coordinate bond is still a covalent bond.
• It is usually drawn and treated just like an ordinary covalent bond.

Conceptually:

• One atom acts as an electron pair donor.
• The other acts as an electron pair acceptor.
• The donor provides a lone pair (nonbonding pair) to form a bond with an empty orbital on the acceptor.

In structural formulas, coordinate bonds are sometimes drawn with an arrow from donor to acceptor, e.g.:

• Donor → Acceptor

This type of bonding is particularly important in:

• Many cation–molecule interactions (e.g. $NH_3$ donating a lone pair to a proton).
• Complex ions and coordination compounds (treated in detail in coordination chemistry).

Covalent Bonding and Molecular Shape (Qualitative Connection)

Covalent bonds are directional: the shared electron pair is located in a specific region relative to the nuclei. As a result:

• Atoms in a covalent molecule are not randomly arranged.
• They adopt geometries that:
• Minimize repulsions between electron pairs.
• Maximize stability of the molecule.

Important points:

• Lone pairs (nonbonding electron pairs) and bonding pairs around an atom influence the overall molecular shape.
• Bond angles (e.g. about $109.5^\circ$, $120^\circ$, $180^\circ$) are characteristic of certain arrangements.

Examples (qualitative):

• $CH_4$: four single covalent C–H bonds arranged approximately in a tetrahedral geometry around carbon.
• $H_2O$: two O–H bonds plus two lone pairs on oxygen; this leads to a bent (V-shaped) molecular structure.

The systematic prediction of 3D shapes from electron pair arrangements and the detailed models used to do this are covered separately; here, the key idea is that covalent bonds fix specific spatial relationships between atoms.

Comparison to Other Main Bond Types (Brief, Qualitative)

Within the overall context of main bond types, covalent bonds are characterized by:

• Sharing of electron pairs between specific atoms (directional bonds).
• Typical occurrence between nonmetal atoms.
• Formation of discrete molecules or extended covalent networks (e.g. diamond, quartz).
• Strong dependence of properties on molecular structure (e.g. shape, polarity, functional groups).

In contrast (details in other chapters):

• Ionic bonds involve overall transfer of electrons and electrostatic attraction between ions.
• Metallic bonds involve a lattice of positive centers with delocalized electrons spread over the whole solid.

Covalent bonding is therefore the key type of bonding for understanding the structures and properties of most molecular substances, especially in organic and biological chemistry.