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Single, double and triple bonds Knowledge-Based Study

Chemistry · Chemical Bonding & Molecular Structure · Chemistry · Covalent Bonding

Chemistry topic explanations on Single, double and triple bonds.

Study notes

  1. A covalent bond forms when two nonmetal atoms share valence electrons to attain a stable, noble gas electron configuration. Depending on the number of electron pairs shared between the two bonded nuclei, covalent bonds are classified as single, double, or triple bonds.

    The concept of bond order quantifies this relationship:

    • Single bond: Bond order of 1 (one shared pair / 2 electrons).
    • Double bond: Bond order of 2 (two shared pairs / 4 electrons).
    • Triple bond: Bond order of 3 (three shared pairs / 6 electrons).

    A single covalent bond consists of one shared pair of electrons between two atoms. It is represented in Lewis structures by a single line segment or two dots between the atomic symbols.

    Orbital Overlap and Nature

    Single bonds are typically σ (sigma) bonds, formed by the direct head-on (end-to-end) overlap of atomic orbitals along the internuclear axis. Because the electron density is concentrated along the line connecting the nuclei, σ bonds permit free rotation around the bond axis.

    Representative Examples

    • Hydrogen gas (H2): Each hydrogen atom shares its single 1s electron, resulting in a single H–H bond.
    • Chlorine gas (Cl2): Two chlorine atoms share one pair of p-electrons to complete an octet, forming a Cl–Cl bond.
    • Methane (CH4): The central carbon atom forms four separate single C–H bonds with four hydrogen atoms.

    A double covalent bond occurs when two atoms share two pairs of electrons (a total of 4 electrons). In structural formulas, it is denoted by two parallel lines (=).

    Orbital Overlap: σ and π Bonds

    A double bond consists of:

    • One σ (sigma) bond formed by end-to-end orbital overlap along the internuclear axis.
    • One π (pi) bond formed by the sideways (lateral) overlap of unhybridized parallel p orbitals above and below the internuclear axis.

    Because the π bond fixes the relative orientation of the overlapping p orbitals, free rotation around a double bond is restricted at standard temperatures.

    Representative Examples

    • Oxygen gas (O2): Each oxygen atom contributes two valence electrons, sharing two pairs (O=O) to achieve an octet.
    • Ethene (C2H4): Contains a double bond between the two carbon atoms (H2C=CH2).
    • Carbon dioxide (CO2): Carbon shares two electron pairs with each of the two oxygen atoms, forming two double bonds (O=C=O).

    A triple covalent bond involves the sharing of three pairs of electrons (a total of 6 electrons) between two atoms. It is symbolized by three parallel lines (≡).

    Orbital Overlap: σ and Two π Bonds

    A triple bond consists of:

    • One σ (sigma) bond directed along the internuclear axis.
    • Two mutually perpendicular π (pi) bonds formed by the side-by-side overlap of two sets of parallel p orbitals.

    The high electron density between the nuclei pulls them closer together, making triple bonds exceptionally short and strong.

    Representative Examples

    • Nitrogen gas (N2): Two nitrogen atoms share three pairs of electrons (N≡N), producing an exceptionally stable molecule with high dissociation energy.
    • Ethyne / Acetylene (C2H2): Features a carbon-carbon triple bond (H–C≡C–H).
    • Carbon monoxide (CO): Features a triple bond between carbon and oxygen (:C≡O:).

    Comparing covalent bonds between the same pair of elements (such as C–C, C=C, and C≡C) reveals fundamental physical and thermodynamic relationships:

    • Bond Order: Increases from single (1) to double (2) to triple (3).
    • Bond Length: Decreases as bond order increases. Greater electrostatic attraction between more shared electrons and both nuclei draws the nuclei closer together:
      Single > Double > Triple (C–C ≈ 154 pm, C=C ≈ 134 pm, C≡C ≈ 120 pm).
    • Bond Dissociation Energy (Strength): Increases as bond order increases. More energy is required to break multiple shared pairs:
      Triple > Double > Single (C–C ≈ 348 kJ/mol, C=C ≈ 614 kJ/mol, C≡C ≈ 839 kJ/mol).

    Note: A double bond is stronger than a single bond, but not twice as strong, because a π bond provides less effective orbital overlap than a σ bond.