bond

Examples of bond 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.

• Double and Triple Covalent Bonds

  • The double bond between the two carbon atoms consists of a sigma bond and a π bond.
  • A triple bond involves the sharing of six electrons, with a sigma bond and two $\pi$ bonds.
  • Triple bonds are stronger than double bonds due to the the presence of two $\pi$ bonds rather than one.
  • Experiments have shown that double bonds are stronger than single bonds, and triple bonds are stronger than double bonds.
  • Double bonds have shorter distances than single bonds, and triple bonds are shorter than double bonds.

• Hydrogen Bonding

  • A hydrogen bond is a type of dipole-dipole interaction; it is not a true chemical bond.
  • This hydrogen atom is a hydrogen bond donor.
  • Greater electronegativity of the hydrogen bond acceptor will create a stronger hydrogen bond.
  • Hydrogen bonds are shown with dotted lines.
  • Where do hydrogen bonds form?

• 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.

• Ionic vs Covalent Bond Character

  • The bond formed between any two atoms is not a purely ionic bond.
  • In the conventional presentation, bonds are designated as ionic when the ionic aspect is greater than the covalent aspect of the bond.
  • Bonds that fall in between the two extremes, having both ionic and covalent character, are classified as polar covalent bonds.
  • This bond is considered to have characteristics of both covalent and ionic bonds.
  • Discuss the idea that, in nature, bonds exhibit characteristics of both ionic and covalent bonds

• Bond Energy

  • Since reactions of organic compounds involve the making and breaking of bonds, the strength of bonds, or their resistance to breaking, becomes an important consideration.
  • Bond energy is the energy required to break a covalent bond homolytically (into neutral fragments).
  • Bond energies are commonly given in units of kcal/mol or kJ/mol, and are generally called bond dissociation energies when given for specific bonds, or average bond energies when summarized for a given type of bond over many kinds of compounds.
  • The following table is a collection of average bond energies for a variety of common bonds.
  • First, a single bond between two given atoms is weaker than a double bond, which in turn is weaker than a triple bond.

• Bond Lengths

  • The distance between two atoms participating in a bond, known as the bond length, can be determined experimentally.
  • The bond length is the average distance between the nuclei of two bonded atoms in a molecule.
  • Bonds involving hydrogen can be quite short; the shortest bond of all, H–H, is only 74 pm.
  • Atoms with multiple bonds between them have shorter bond lengths than singly bonded ones; this is a major criterion for experimentally determining the multiplicity of a bond.
  • For example, the bond length of $C - C$ is 154 pm; the bond length of $C = C$ is 133 pm; and finally, the bond length of $C \equiv C$ is 120 pm.

• Types of Bonds

  • Ionic bonds can form between nonmetals and metals, while covalent bonds form when electrons are shared between two nonmetals.
  • Pure ionic bonding cannot exist: all ionic compounds have some degree of covalent bonding.
  • Thus, an ionic bond is considered a bond where the ionic character is greater than the covalent character.
  • Bonds with partially ionic and partially covalent character are called polar covalent bonds.
  • This difference in charge is called a dipole, and when the covalent bond results in this difference in charge, the bond is called a polar covalent bond.

• Covalent Bonds

  • Covalent bonding interactions include sigma-bonding (σ) and pi-bonding (π).
  • Single bonds occur when two electrons are shared and are composed of one sigma bond between the two atoms.
  • 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.

• Bond Enthalpy

  • Bond enthalpy is defined as the enthalpy change when a covalent bond is cleaved by homolysis.
  • In other words, breaking a bond is an endothermic process, while the formation of bonds is exothermic.
  • For instance, the bond enthalpy, or bond-dissociation energy, for one of the C-H bonds in ethane (C2H6) is defined by the process:
  • Each bond in a molecule has its own bond dissociation energy, so a molecule with four bonds will require more energy to break the bonds than a molecule with one bond.
  • As each successive bond is broken, the bond dissociation energy required for the other bonds changes slightly.