bond order

(noun)

the number of overlapping electron pairs between a pair of atoms

Related Terms

Examples of bond order in the following topics:

  • Bond Order

    • Bond order is the number of chemical bonds between a pair of atoms.
    • Bond order is the number of chemical bonds between a pair of atoms; in diatomic nitrogen (N≡N) for example, the bond order is 3, while in acetylene (H−C≡C−H), the bond order between the two carbon atoms is 3 and the C−H bond order is 1.
    • Bond order indicates the stability of a bond.
    • Bond order is also an index of bond strength, and it is used extensively in valence bond theory.
    • For a bond to be stable, the bond order must be a positive value.

  • Bonding in Coordination Compounds: Valence Bond Theory

    • Valence bond theory is used to explain covalent bond formation in many molecules.
    • Valence bond theory is a synthesis of early understandings of how chemical bonds form.
    • Lewis proposed that the basis of chemical bonding is in the ability of atoms to share two bonding electrons.
    • Where bond order is concerned, single bonds are considered to be one sigma bond, double bonds are considered to contain one sigma and one pi bond, and triple bonds consist of one sigma bond and two pi bonds.
    • Valence bond theory is used to explain covalent bond formation in many molecules, as it operates under the condition of maximum overlap, which leads to the formation of the strongest possible bonds.

  • Explanation of Valence Bond Theory

    • Valence bond theory states that overlap between two atomic orbitals forms a covalent bond between two atoms.
    • In chemistry, valence bond (VB) theory is one of two basic theories—along with molecular orbital (MO) theory—that use quantum mechanics to explain chemical bonding.
    • Both types of overlapping orbitals can be related to bond order.
    • Single bonds have one sigma bond.
    • Double bonds consist of one $\sigma$ and one $\pi$ bond, while triple bonds contain one $\sigma$ and two $\pi$ bonds.

  • Bond Energy

    • Bond energy is the measure of bond strength.
    • In order to turn one mole of a molecule into its constituent atoms, an amount of heat equal to the bond energy needs to be put into the system.
    • The bond energy is the average of the bond dissociation energies in a molecule.
    • At internuclear distances in the order of an atomic diameter, attractive forces dominate.
    • Identify the relationship between bond energy and strength of chemical bonds

  • Atomic and Molecular Orbitals

    • In order to explain this covalent bonding, Linus Pauling proposed an orbital hybridization model in which all the valence shell electrons of carbon are reorganized.
    • In order to explain the structure of methane (CH4), the 2s and three 2p orbitals are converted to four equivalent hybrid atomic orbitals, each having 25% s and 75% p character, and designated sp3.
    • When these bonding orbitals are occupied by a pair of electrons, a covalent bond, the sigma bond results.
    • Since bonds consisting of occupied π-orbitals (pi-bonds) are weaker than sigma bonds, pi-bonding between two atoms occurs only when a sigma bond has already been established.
    • The overall bonding order depends on the number of antibonding orbitals that are occupied.

  • Covalent Bonds

    • Covalent bonding requires a specific orientation between atoms in order to achieve the overlap between bonding orbitals.
    • Covalent bonding interactions include sigma-bonding (σ) and pi-bonding (π).
    • Double bonds occur when four electrons are shared between the two atoms and consist of one sigma bond and one pi bond.
    • Triple bonds occur when six electrons are shared between the two atoms and consist of one sigma bond and two pi bonds (see later concept for more info about pi and sigma bonds).
    • Unlike an ionic bond, a covalent bond is stronger between two atoms with similar electronegativity.

  • Single Covalent Bonds

    • Single covalent bonds are sigma bonds, which occur when one pair of electrons is shared between atoms.
    • The strongest type of covalent bonds are sigma bonds, which are formed by the direct overlap of orbitals from each of the two bonded atoms.
    • A single covalent bond can be represented by a single line between the two atoms.
    • The shapes of the first five atomic orbitals are shown in order: 1s, 2s, and the three 2p orbitals.
    • Notice that the area of overlap always occurs between the nuclei of the two bonded atoms.

  • Introduction to Lewis Structures for Covalent Molecules

    • In covalent molecules, atoms share pairs of electrons in order to achieve a full valence level.
    • Other elements in the periodic table react to form bonds in which valence electrons are exchanged or shared in order to achieve a valence level which is filled, just like in the noble gases.
    • It therefore has 7 valence electrons and only needs 1 more in order to have an octet.
    • In order to achieve an octet for all three atoms in CO2, two pairs of electrons must be shared between the carbon and each oxygen.
    • Since four electrons are involved in each bond, a double covalent bond is formed.

  • Halogenation

    • The reactivity of the halogens decreases in the following order: F2 > Cl2 > Br2 > I2.
    • alkyl radical stability increases in the order: phenyl < primary (1º) < secondary (2º) < tertiary (3º) < allyl ≈ benzyl.
    • Indeed, the stability order of alkyl radicals (primary (1º) < secondary (2º) < tertiary (3º) < allyl ≈ benzyl) varies inversely with the corresponding C–H bond dissociation energies, as expected.
    • The ionic character of a single covalent bond increases with the electronegativity difference between the bonded atoms.
    • It is noteworthy that the C–O BDEs of alcohols are found to increase in the order: secondary (2º) > tertiary (3º) > primary (1º) > methyl, nearly opposite the order of C–H BDEs.

  • The Shape of Molecules

    • This shape is dependent on the preferred spatial orientation of covalent bonds to atoms having two or more bonding partners.
    • In order to represent such configurations on a two-dimensional surface (paper, blackboard or screen), we often use perspective drawings in which the direction of a bond is specified by the line connecting the bonded atoms.
    • A wedge shaped bond is directed in front of this plane (thick end toward the viewer), as shown by the bond to substituent B; and a hatched bond is directed in back of the plane (away from the viewer), as shown by the bond to substituent D.
    • Some texts and other sources may use a dashed bond in the same manner as we have defined the hatched bond, but this can be confusing because the dashed bond is often used to represent a partial bond (i.e. a covalent bond that is partially formed or partially broken).
    • In the linear configuration (bond angle 180º) the bond dipoles cancel, and the molecular dipole is zero.