Chemical Bonding

There are different types of chemical bonds:

1. Ionic or Electrovalent Bonding
2. Covalent Bonding
3. Coordinate Covalent (Dative) Bonding
4. Hydrogen Bond
5. Metallic Bond
6. Van der Waals Force

Electrovalent (Ionic) Bonding

An electrovalent bond involves the transfer of electrons from metallic atoms to non-metallic atoms during a chemical reaction. It follows the donor-acceptor principle, where the metal atom becomes positively charged after losing electrons, and the non-metal becomes negatively charged after gaining electrons. These oppositely charged ions are held together by a strong electrostatic force, forming an ionic bond.

Examples of ionic compounds include calcium chloride (CaCl₂), magnesium oxide (MgO), and magnesium chloride (MgCl₂).

Factors Influencing Ionic Bond Formation

• Ionization energy
• Electron affinity
• Electronegativity difference

Properties of Electrovalent (Ionic) Compounds

• Solid at room temperature
• High melting and boiling points
• Composed of aggregates of ions, not molecules
• Dissolve easily in water and polar solvents
• Conduct electricity when molten or in solution (good electrolytes)
• Insoluble in non-polar solvents like benzene, toluene, ether, and trichloromethane

Covalent Bonding

A covalent bond forms when two atoms share a pair of electrons. This overlap of electron clouds can occur between atoms of the same element or between atoms with similar electronegativities. Each atom contributes one electron to the shared pair, and the bond is often represented by a horizontal line, such as H–H or H–Cl. Double and triple bonds involve the sharing of two and three pairs of electrons, respectively.

Factors Influencing Covalent Bond Formation

• Ionization energy
• Electron affinity
• Electronegativity difference

Properties of Covalent Compounds

• Consist of molecules with definite shapes
• Gases or volatile liquids at room temperature
• Low melting and boiling points
• Soluble in non-polar solvents like toluene and carbon disulfide
• Non-conductors of electricity (non-electrolytes)

Coordinate Covalent Bond (Dative Bond)

A coordinate covalent bond is a type of covalent bond where the shared pair of electrons is donated by only one atom. It is represented by an arrow pointing from the donor to the acceptor atom.

Examples

• Formation of Hydroxonium ion (H₃O⁺)
• Ammonium ion (NH₄⁺)
• Hydrated copper (II) ion [Cu(H₂O)₄]²⁺
• Tetraamine copper (II) ion [Cu(NH₃)₄]²⁺
• Phosphorus oxochloride (POCl₃)

Compounds with coordinate covalent bonds generally share similar properties with covalent compounds, though they tend to be less volatile.

Hydrogen Bond

A hydrogen bond is a dipole-dipole intermolecular attraction that occurs when hydrogen is covalently bonded to highly electronegative elements with small atomic sizes, such as oxygen (O₂), nitrogen (N₂), and fluorine (F₂).

Due to their strong affinity for electrons, these electronegative elements attract the shared pair of electrons closer, making hydrogen partially positively charged and the other atom partially negatively charged. An electrostatic force of attraction is formed between the two dipoles, known as the hydrogen bond.

The strongest hydrogen bonds are observed in hydrogen fluoride (HF). Although hydrogen bonds are relatively weak, they have significant effects on the physical properties of compounds like hydrogen fluoride and water.

Metallic Bonding

Metallic bonding occurs only in metals. It involves loosely held electrons in the outermost shells and the positively charged protons in the nuclei of metal atoms. This type of bonding is responsible for many of the typical properties of metals.

Metallic bonds are not strongly directional, which allows metals to bend or deform without breaking their crystal structures under pressure, such as during hammering. As a result, metals are malleable and ductile. They are also good conductors of heat and electricity because the outer electrons are loosely held and free to move throughout the solid structure.

When a metal is heated, these free electrons carry heat energy; when a potential difference is applied, they carry electric charge. Due to the loose hold on their outermost electrons, metals easily lose electrons during chemical reactions.

In a metal lattice, the free electrons are not attached to a specific atom and are described as delocalized electrons. These delocalized electrons are attracted to surrounding atomic nuclei, helping to bind the atoms together. The more delocalized electrons present per atom, the stronger the metallic bond, leading to higher melting points.

Van der Waals Forces

Van der Waals forces are weak attractive forces that exist even between separate molecules. They were first described by J.D. Van der Waals. Although much weaker than ionic or covalent bonds, Van der Waals forces play a crucial role in the liquefaction of gases and the formation of molecular lattices, such as those found in iodine and naphthalene crystals.