ATOMIC STRUCTURE AND BONDING FREE NOTES

1. What is Atomic Structure ?

Atomic structure refers to the constitution of an atom, which is the smallest unit of ordinary matter that forms a chemical element. Every solid, liquid, gas, and plasma is composed of neutral or ionized atoms.

Basic Concept

An atom consists of three primary subatomic particles, each defined by its mass and electrical charge:

• Protons: Found in the central nucleus, protons have a positive charge (+1). The number of protons determines the atomic number (Z) and identifies the element.
• Neutrons: Also located in the nucleus, these have no charge (neutral). They add mass and act as a “buffer” to stabilize the protons.
• Electrons: These are much smaller particles with a negative charge (-1). They reside in the space surrounding the nucleus.

2. Atomic Models

2.1 Bohr’s Model

The Bohr Atomic Model was proposed by Niels Bohr in 1913 to overcome the limitations of Ernest Rutherford’s nuclear model. It successfully explained the structure of the hydrogen atom and the origin of its line spectrum.

According to Bohr, electrons revolve around the nucleus only in certain fixed circular paths called orbits or energy levels without losing energy.

2.2 Quantum Mechanical Model

The Quantum Mechanical Model (also called the Electron Cloud Model or Wave Mechanical Model) is the modern and most widely accepted model of the atom. It was developed in 1926 by Erwin Schrödinger based on the ideas of Louis de Broglie, Werner Heisenberg, Max Born, and Paul Dirac.

Unlike Niels Bohr’s model, the quantum mechanical model does not assume that electrons move in fixed circular orbits. Instead, electrons are described by wave functions, and their exact positions cannot be known with certainty. The model predicts the probability of finding an electron in a particular region around the nucleus.

Electronic Configuration

• Arrangement of electrons in shells and subshells
• Governed by:
  • Aufbau Principle (lowest energy first)
  • Pauli Exclusion Principle (max 2 electrons/orbital)
  • Hund’s Rule (maximize unpaired electrons)

Example:

• Sodium (Na): 1s² 2s² 2p⁶ 3s¹

Periodic Properties (Important for Engineering Materials)

• Atomic Radius → size of atom
• Ionization Energy → energy to remove electron
• Electron Affinity → tendency to gain electron
• Electronegativity → attraction for electrons

These properties influence bonding and material behavior.

2. Interatomic Bonding

Interatomic bonding is the attractive force that holds atoms together to form molecules, compounds, and solids. Atoms form bonds to achieve a stable electronic configuration, usually by completing their outermost (valence) electron shell. Most atoms become stable by attaining the electron configuration of a noble gas (the octet rule for many main-group elements).

Interatomic bonds determine the physical, chemical, electrical, and mechanical properties of materials.

2.1 Primary Bonds (Strong Bonds)

2.1.1. Ionic Bond

An ionic bond (also called an electrovalent bond) is a type of chemical bond formed by the complete transfer of one or more electrons from one atom to another. This transfer creates positively charged ions (cations) and negatively charged ions (anions), which are held together by a strong electrostatic force of attraction.

Ionic bonding usually occurs between a metal and a non-metal. Metals tend to lose electrons, while non-metals tend to gain electrons to achieve a stable electron configuration (usually the octet rule).

Formation of an Ionic Bond

The formation of an ionic bond involves three main steps:

1. A metal atom loses one or more valence electrons to become a cation.
2. A non-metal atom gains those electrons to become an anion.
3. The oppositely charged ions attract each other through electrostatic forces, forming an ionic bond.

Properties:

• High melting point
• Brittle
• Good electrical conductivity in molten state

2.1.2. Covalent Bond

• Formed by sharing of electrons
• Between non-metals

Properties:

• Strong and directional
• Low electrical conductivity
• Example: Diamond, Silicon

2.1.3. Metallic Bond

• “Sea of free electrons” around positive ions

Properties:

• High electrical & thermal conductivity
• Ductile and malleable
• Example: Iron, Copper

2.2 Secondary Bonds (Weak Bonds)

2.2.1. Van der Waals Forces

• Weak attraction between molecules
• Important in polymers

2.2.2. Hydrogen Bonding

• Stronger than Van der Waals
• Occurs in water, polymers

3. Energy vs Interatomic Distance Curve

Key Concepts:

• Attractive Forces → pull atoms together
• Repulsive Forces → push atoms apart
• Equilibrium (r₀) → stable spacing
• Bond Energy → energy required to separate atoms

This curve explains:

• Elasticity
• Thermal expansion
• Strength of materials

4. Relation Between Bonding and Material Properties

Bond Type	Strength	Electrical Conductivity	Example Materials
Ionic	High	Low (solid)	Ceramics
Covalent	Very High	Very Low	Diamond, Si
Metallic	Moderate	High	Metals
Secondary	Low	Very Low	Polymers

5. Engineering Significance

• Mechanical Properties:
  • Strong bonds → high strength & hardness
• Electrical Properties:
  • Free electrons → conductivity (metals)
• Thermal Properties:
  • Bond strength affects melting point
• Material Selection:
  • Engineers choose materials based on bonding type

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