Chemical Change and Chemical Bond

Physical and Chemical Changes

Chemical Bonding

Elements are rarely capable of free existence. In a compound, atoms of different elements are held together by bonds. The types of bonds present in a compound are largely responsible for its physical and chemical properties. The different bonds can be classified as strong and weak.

Why do elements undergo bond formation?

Elements are made of atoms, which comprise of protons, electrons, and neutrons. The protons and the neutrons reside in the nucleus and the electrons revolve around in definite paths called orbits. The electrons present in the last shell are called valence electrons. These electrons are responsible for all the chemical reactions of that element.

Every element has a tendency to attain a stable outer octet. To do so, it either gains or loses or shares its electrons; and in this process, it forms the bonds.

Types of strong bonds:

• Ionic or electrovalent bond
• Covalent bond
• Metallic bond

Types of weak bonds:

• Bonds formed due to van der Waal’s interaction
• Hydrogen bond

This representation of elements with valence electrons as dots around elements is referred to as Electron Dot structures for elements. The electron dot structure of some of the elements are: 

Chemical Bonding

Elements are rarely capable of free existence. In a compound, atoms of different elements are held together by bonds. The types of bonds present in a compound are largely responsible for its physical and chemical properties. The different bonds can be classified as strong and weak.

Why do elements undergo bond formation?

Elements are made of atoms, which comprise of protons, electrons, and neutrons. The protons and the neutrons reside in the nucleus and the electrons revolve around in definite paths called orbits. The electrons present in the last shell are called valence electrons. These electrons are responsible for all the chemical reactions of that element.

Every element has a tendency to attain a stable outer octet. To do so, it either gains or loses or shares its electrons; and in this process, it forms the bonds.

Types of strong bonds:

• Ionic or electrovalent bond
• Covalent bond
• Metallic bond

Types of weak bonds:

• Bonds formed due to van der Waal’s interaction
• Hydrogen bond

This representation of elements with valence electrons as dots around elements is referred to as Electron Dot structures for elements. The electron dot structure of some of the elements are: 

A chemical bond is an attractive force which holds various constituents (such as atoms, ions) together in different chemical species.

Kossel-Lewis Approach to Chemical Bonding

• Lewis postulated that atoms attain the stable octet when they are chemically bonded.

• Lewis symbols

• Notations to represent valence electrons in an atom

• Example:

• Significance of Lewis symbols − The number of dots represents the number of valence electrons.

• Octet rule- Atoms tend to gain, lose or share electrons so as to have eight electrons in their valence shells.

            

 

• Lewis dot structure

Representation of molecules and ions in terms of the shared pairs of electrons and the octet rule

Steps to writing Lewis dot structure:

• Add the valence electrons of the combining atoms. This will give the total number of electrons required to write the structure.

• One negative charge means the addition of an electron. Similarly, one positive charge implies the removal of an electron from the total number of electrons.

• The chemical symbol of the atoms and the skeletal structure of the compound should be known. Then, distribute the total number of electrons as bonding shared pairs between the atoms in proportion to the total bonds.

• The least electronegative atom occupies the central position of the molecule/ion. For example in NF₃, nitrogen occupies the central position whereas the three fluorine atoms occupy the terminal positions.

• When the shared pairs of electrons have been accounted for single bonds, utilise the remaining electron pairs for either multiple bonding or count them as lone pairs. Here, the basic requirement is that each bonded atom gets an octet of electrons.

• Lewis representation of some molecules

 

(*- Each hydrogen atom attains the electronic configuration of helium i.e. a duplet of electrons)

• Covalent bond

• Single covalent bond − Sharing of one electron pair

• Double bond − Sharing of two electron pairs

• Triple bond − Sharing of three electron pairs

• Formal Charge

• Example:

Lewis structure of O₃ is

F.C. on the O-1 atom

F.C. on the O-2 atom

F.C. on the O-3 atom

• Smaller the formal charge on the atoms, lower is the energy of the structure.
• The concept of formal charge is based on covalent bonding in which electron pairs are equally shared by neighbouring atoms.

• Limitations of the octet rule:

• Incomplete octet of the central atom

Examples: LiCl, BeH₂, BCl₃

• Odd electron molecules

Examples: NO, NO₂

• Expanded octet

Examples: PF₅, SF₆, H₂SO₄

Some other drawbacks of octet rule:

• It is based upon chemical inertness of noble gases. However, some noble gases can combine to form compounds such as XeF₂, KrF₂, XeOF₂, etc.

• It does not account for shape of molecules

• It does not explain the relative stability of molecules

Conditions for Formation of Covalent Bond 

• Presence of four or more elec…