Bond Energy

Examples of Bond Energy in the following topics:

• Bond Energy

  • 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.
  • Tables of bond energies may be found in most text books and handbooks.
  • The following table is a collection of average bond energies for a variety of common bonds.
  • Such average values are often referred to as standard bond energies, and are given here in units of kcal/mole.

• Bond Energy

  • Bond energy is the measure of bond strength.
  • Bond energy is a measure of a chemical bond's strength, meaning that it tells us how likely a pair of atoms is to remain bonded in the presence of energy perturbations.
  • These energy values (493 and 424 kJ/mol) required to break successive O-H bonds in the water molecule are called 'bond dissociation energies,' and they are different from the bond energy.
  • The bond energy is the average of the bond dissociation energies in a molecule.
  • Identify the relationship between bond energy and strength of chemical bonds

• Bond Enthalpy

  • Bond enthalpy, also known as bond dissociation energy, is defined as the standard enthalpy change when a bond is cleaved by homolysis, with reactants and products of the homolysis reaction at 0 K (absolute zero).
  • 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.
  • Bond dissociation energies for different element pairings are listed.

• Energy Changes in Chemical Reactions

  • Due to the absorption of energy when chemical bonds are broken, and the release of energy when chemical bonds are formed, chemical reactions almost always involve a change in energy between products and reactants.
  • By the Law of Conservation of Energy, however, we know that the total energy of a system must remain unchanged, and that oftentimes a chemical reaction will absorb or release energy in the form of heat, light, or both.
  • The energy change in a chemical reaction is due to the difference in the amounts of stored chemical energy between the products and the reactants.
  • This means that the energy required to break the bonds in the reactants is less than the energy released when new bonds form in the products.
  • This means that the energy required to break the bonds in the reactants is more than the energy released when new bonds form in the products; in other words, the reaction requires energy to proceed.

• Bonding and Antibonding Molecular Orbitals

  • MO modeling is only valid when the atomic orbitals have comparable energy; when the energies differ greatly, the bonding mode becomes ionic.
  • This MO is called the bonding orbital, and its energy is lower than that of the original atomic orbitals.
  • The reduction these electrons' energy is the driving force for chemical bond formation.
  • Whenever symmetry or energy make mixing an atomic orbital impossible, a non-bonding MO is created; often quite similar to and with energy levels equal or close to its constituent AO, the non-bonding MO creates an unfavorable energy event.
  • Bonding and antibonding levels in the hydrogen molecule; the two electrons in the hydrogen atoms occupy a bonding orbital that is lower in energy than the two separate electrons, making this an energy-favorable event.

• Bond Order

  • Bond order is the number of chemical bonds between a pair of atoms.
  • 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.
  • In the second diagram, one of the bonding electrons in H2 is "promoted" by adding energy and placing it in the antibonding level.
  • By adding energy to an electon and pushing it to the antibonding orbital, this H2 molecule's bond order is zero, effectively showing a broken bond.

• Double and Triple Covalent Bonds

  • The newly formed hybrid orbitals all have the same energy and have a specific geometrical arrangement in space that agrees with the observed bonding geometry in molecules.
  • Covalent bonds can be classified in terms of the amount of energy that is required to break them.
  • Based on the experimental observation that more energy is needed to break a bond between two oxygen atoms in O2 than two hydrogen atoms in H2, we infer that the oxygen atoms are more tightly bound together.
  • Therefore, it would take more energy to break the triple bond in N2 compared to the double bond in O2.
  • Double bonds have shorter distances than single bonds, and triple bonds are shorter than double bonds.

• Single Covalent Bonds

  • Single covalent bonds are sigma bonds, which occur when one pair of electrons is shared between atoms.
  • There are four hierarchical levels that describe the position and energy of the electrons an atom has.
  • Principal energy levels are made out of sublevels, which are in turn made out of orbitals, in which electrons are found.
  • 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.

• Bond Lengths

  • The bond length is the average distance between the nuclei of two bonded atoms in a molecule.
  • This is because a chemical bond is not a static structure, but the two atoms actually vibrate due to thermal energy available in the surroundings at any non-zero Kelvin temperature.
  • 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.
  • The potential energy function for this system is also indicated.
  • The minimum energy occurs at the equilibrium distance r0, which is where the bond length is measured.

• 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.
  • The electrons donated by the ligand move into hybridized orbitals of higher energy, which are then filled by electron pairs donated by the ligand.