Bond Energy

Bond dissociation energy is commonly referred to as bond energy. It is usually interpreted positively since it refers to the amount of energy required rather than the amount of energy released. It is the energy of one mole of a specific bond type, not one mole of the material. Bond energy refers to the bond dissociation energy of a diatomic molecule. On the other hand, bond dissociation energy is determined by the type of bond and the molecule in which it exists. The mean estimate of the dissociation energy of a specific bond is taken when a substance’s molecule contains more than one bond of the same type. Bond energy is the aggregate bond dissociation energy necessary to break every bond in a molecule.

An example

It is imperative to remember that a chemical bond’s energy in a compound is the mean of all of the chemical bonds’ unique bond dissociation enthalpies. Imagine two hydrogen atoms A and B approaching each other and having nuclei NA and NB, respectively, with electrons eA and eB. There is no contact between two atoms when a great distance separates them. New attraction and repulsion energies emerge when these two atoms reach one other.

Representation

∆bH or ∆bond H Chemical bonding is linked to enthalpy changes. Enthalpy variations linked with chemical bonding are expressed using two separate terminologies. They are:

• Bond dissociation enthalpy
• Mean or Average bond enthalpy

The bond dissociation enthalpy is used for diatomic compounds, but for polyatomic molecules, a calculation based on the mean of bond energy is used.

Bond order

The total number of covalently bound electron pairs connecting two atoms in a molecule is the bond order of a bond. It may be discovered by sketching the molecule’s Lewis structure and calculating the absolute number of electron pairs among the atoms.

• Single bonds possess a bond order of 1
• Double bonds possess a bond order of 2
• Triple bonds possess a bond order of 3

According to molecular orbital theory, a covalent bond’s order is equivalent to half of the discrepancy among the number of bonding and antibonding electrons, as expressed by the following formula: Bond Order Formula = (½)*(total bonding electrons – total antibonding electrons)

Factors Affecting Bond Energy

Many factors affect the bond energy:

• The bond length is proportional to the atom’s size and the bond dissociation energy. This means less bond energy applies to smaller atom sizes

• The bond dissociation energy of a bond between two identical atoms increases as the bond multiplicity increases

• The repulsion between the atoms increases as the number of isolated pairs of electrons on the connected atoms increases, and therefore the bond dissociation energy decreases

• As the s orbital contribution to the hybrid orbitals rises, the bond energy rises. As a result, bond energy reduces in the sequence listed as: sp > sp² > sp³

• The larger the electronegativity discrepancy, the stronger the bond polarity, and thus the bond strength (bond energy). H-F > H-Cl > HBr > H-I

• Resonance raises bond energy among halogens Cl – Cl > F – F > Br – Br > I – I, descending sequence of bond energy

Different rules of bonding

• Octet rule

According to the electrical theory of valence, every atom seeks to achieve the octet arrangement, the existence of eight electrons in its valence shell by shedding, gaining or splitting electrons. The octet rule is the name for this concept. The Octet rule has a few exceptions:

• Species having an odd number of electrons include the following: NO, NO
• LiCl, BeH, and BCl form an incomplete octet for the centre atom.
• PF, SF, and H SO are the expanded octets for the centre atom.

• Lewis structures

Lewis invented the notion of portraying valence electrons with dots, which are known as Lewis symbols, to help people understand the basic concept of valence electrons.

• A ‘dash’ (-) commonly represents a ‘bond’ which shows a pair of bonded electrons, whereas ‘dots’ indicate lone pairs or ‘non-bonded’ electrons.
• Valence electrons are electrons found in the final shell of an atom.

• Ionic bonding

The electrostatic attraction among positive and negative ions in a chemical molecule forms an ionic bond, also known as an electrovalent bond. If one atom’s valence (outermost) electrons are irreversibly transferred to some other atom, a bond is formed. The transport of electrons between a nonmetal and metal results in the creation of an ionic bond.

• Covalent bonding

A covalent bond occurs when two atoms, whether the same or different, combine their valence electrons to achieve the noble gas state. The sharing of electrons forms a covalent link between two nonmetals. Bond pairs are electron pairs that engage in bonding. The interatomic bond forms when two atoms share an electron pair. A covalent bond forms when the total energy of the bound atoms is lower than that of widely spaced atoms.

Ionic Bond	Covalent Bond
An ionic bond is essentially a bond between metals and nonmetals. Because the nonmetal pulls the electron, it is as if the metal gives it one of its own electrons.	Two nonmetals with identical electronegativities form a covalent bond. In the outer orbitals, atoms share electrons.
The electronegativity values of the atoms in an ionic connection differ as an electron is transferred from one atom to another.	The electron gets shared evenly among the atoms forming the covalent connection in a real covalent bond, and the bond is considered to be nonpolar.
Ionic bonds do not have a particular shape.	The shape of covalent bonds is defined.
Example: Sulphuric Acid (H2SO4) Sodium chloride (NaCl)	Example: Hydrochloric acid (HCl) Methane (CH4)