Chemical Bonding

Atoms of elements enter into chemical combination to form stable substances (molecules or compounds). The reason for their entering into the combination is for them to attain the inert octet or duplet structure possessed by the noble gases. Thus, chemical bonding is the coming together of atoms of the same or different elements to form a stable structure. When atoms combine, it is the electrons on the outermost shell of the atom that react by being exchanged, or shared. The attractive force between atoms when they combine chemically is called a chemical bond. There are two main types of chemical bonds, namely:

They are broadly categorized into:

• Intramolecular Forces – forces within a molecule.
• Intermolecular Forces – forces between molecules.

These forces determine the structure, stability, physical properties (like boiling/melting points), and reactivity of substances.

 

Intramolecular Forces

a. Electrovalent or ionic bonding.

b. Covalent bonding

c. Coordinate covalent or dative bonding

d. Metallic bonding

 

Intermolecular Forces

e. Hydrogen bonding

f. Van der Waal forces

g. Ion-dipole

h. Dipole-dipole

 

Electrovalent Bonding

This is also known as ionic bonding. It is a bonding that occurs between metal and non-metal. It involves the transfer of valence electron or electrons from metal to non-metal. The metal loses its valence electron while the non-metal accepts it and the resultant ions are attracted together by a strong force called electrostatic force. The metal acquires positive charge or charges because it contains excess proton or protons in its nucleus while the non-metal acquires negative charge or charges because it has gained extra electron or electrons.

Thus, sodium combines with chlorine to form sodium chloride

 

Examples of Ionic Bond

1. Sodium Chloride (NaCl)

2. Magnesium Oxide (MgO)

3. Calcium Chloride (CaCl₂)

4. Potassium Oxide (K₂O)

 

Characteristics of electrovalent compounds

1. Electrovalent compounds consist of aggregates of ions (charged particles).

2. They have high melting and boiling points.

3. They are soluble in water and other polar solvents.

4. They are strong electrolytes

5. They are insoluble in organic solvents because they are inorganic compounds.

6. They are good conductors of heat and electricity.

  

Covalent Bonding

In this type of bonding, there is sharing of a pair of electrons. Each reacting atom contributes an electron to be shared by the two atoms, so that each has the stable duplet or octet structure of rare gases. This is the type of bonding that occurs between the non-metals. Diatomic molecules of elements are formed by covalent combination. Eg HCl, H₂ and Cl₂ molecules.

Hydrogen combines with chlorine to form hydrogen chloride gas.

 

 

Types of Covalent

Non-polar covalent: Non-polar covalent bonds are covalent bonds formed between two non-metals of the same or about the same electronegativity. In non-polar molecules, the electron pair is shared equally, e.g N₂, H₂, Cl₂, and Br₂

Polar covalent bond: Polar covalent bonds are covalent bonds formed between non–metal atoms having different electronegativity. The electron pair is not shared equally, e.g H₂O, HCl, H₂S, HF.

 

Characteristics of covalent compounds

1. Covalent compounds consist of molecules. They are referred to as molecular compounds.

2. They have low melting and boiling points.

3. They are insoluble in water and other polar solvents.

4. They are non-electrolyte

5. They are usually soluble in organic solvents such as benzene and ethanol

6. They are non-conductors of heat and electricity.

 

Co-ordinate Covalent Bonding

Is also known as dative covalent bonding. In this type of bonding, the shared pair of electrons

(lone pair) are donated by only one of the participants. For dative covalent bonding to occur, one of the reacting substances must possess a lone pair. Example of compounds that exhibit co-ordinate covalent bond: NH₄Cl, NH₄⁺, H₃O⁺.

Note: Ammonia and water molecules possess lone pairs, which enable them to form dative covalent bonds

with a hydrogen proton (H⁺).

 

The ammonia molecule (NH₃) combines with a hydrogen proton (H⁺) to form the ammonium radical (NH₄⁺).

 

 

Metallic Bonding

It is an attractive force that held metal atoms together in crystal lattice. In a crystal lattice, the valence electrons of metal tend to separate from its orbital and move at random forming an electron cloud. The metallic ions formed tend to repel one another but are held together by the attractive force which exists between the electron cloud and the positively charged metallic ions.

This force of attraction is referred to as metallic bond. This metallic bond increases with an increase in the number electron cloud and the positively charged metallic ions. This force of attraction is referred to as metallic bond. This metallic bond increases with an increase in the number of valence electrons for metals in a particular group. For instance, the decreasing order of metallic bond in the crystal lattice of sodium, magnesium and aluminum is as follows:

 

Hydrogen Bonding

It is an intermolecular attraction between two dipoles which occurs when the positive pole of one molecule attracts the negative pole of another. The intermolecular force arises when hydrogen atom is covalently bonded to highly electronegative atoms like oxygen, nitrogen or fluorine. Because of the electronegativity difference between hydrogen and highly electronegative elements, the latter tend to pull the shared pair of electrons in the covalent bonds towards themselves. This brings about the formation of dipole. Hydrogen becomes partially positive and nitrogen, oxygen or fluorine becomes partially negative. As a result of this, electrostatic attraction arises between the partially positive pole of one molecule and the partially negative

pole of another molecule.

 

 

Hydrogen bonding causes the trans association of molecules eg Hydrogen fluoride (HF) can exist

as H₂F₂ (a dimer) or H₃F₃ (a trimer).

 

Vander Waals Forces

It is first studied and described by a Dutchman named J. D. Van der Waals. Hence the force is given the above name. It is a weak force that exists between discrete molecules. It is very important in the formation of molecular lattices in iodine and naphthalene crystals. It is also important in the liquefaction of gases.

This force of attraction is also known as the London Dispersion Force (LDF).

LDFs exist for all substances, whether composed of polar or nonpolar molecules. LDF arises from the formation of temporary instantaneous polarities across a molecule from the circulation of electrons.

An instantaneous polarity in one molecule may induce an opposing polarity in the adjacent molecule, resulting in a series of attractive forces among neighboring molecules. E.g N₂, O₂.

 

Ion-Dipole Forces

It occurs between ions and polar molecules. Very important in solutions, especially in dissolving ionic compounds in water.

Example: Na⁺ interacting with water molecules in NaCl solution.

 

Dipole-Dipole Interactions

It occurs between polar molecules with permanent dipoles. The positive end of one molecule is attracted to the negative end of another.

Examples: HCl, SO₂
Stronger than dispersion forces but weaker than hydrogen bonds. 

 

Comparison of Intramolecular and Intermolecular Forces

Feature	Intramolecular Forces	Intermolecular Forces
Acts Between	Atoms within a molecule	Molecules
Strength	Strong	Weaker
Examples	Covalent, Ionic, Metallic	Dipole-dipole, London, H-bond
Affects	Chemical properties	Physical properties
Bond Formation	Involves sharing/transfer	Involves attractions