Atomic Structure: Foundations of the Atom

Atomic Structure

Introduction

The study of atomic structure is fundamental to understanding the composition and behavior of matter. Atoms are the smallest units of elements, consisting of subatomic particles that determine the chemical and physical properties of substances.

• Atom: The smallest unit of an element, retaining its chemical identity.

• Element: A pure substance consisting of only one type of atom.

• Subatomic particles: Protons, neutrons, and electrons.

Historical Aspects

• Early atomic theory dates back to ancient Greek philosophers, but scientific models began with John Dalton in the early 19th century.

• Dalton's atomic theory provided the first modern scientific description of the atom.

Dalton’s Atomic Theory

Dalton's theory (1808) laid the groundwork for modern chemistry:

• All matter is composed of indivisible atoms.

• Atoms of the same element are identical in mass and properties.

• Atoms of different elements differ in mass and properties.

• Compounds are formed by the combination of atoms in simple whole-number ratios.

• Chemical reactions involve rearrangement of atoms; atoms are neither created nor destroyed.

Limitations: Discovery of subatomic particles and isotopes showed that atoms are divisible and can have different masses.

Discovery of Fundamental Particles

Experiments in the late 19th and early 20th centuries revealed the existence of subatomic particles:

• Electrons: Discovered by J.J. Thomson using cathode ray tube experiments.

• Protons: Discovered by Goldstein and further characterized by Rutherford.

• Neutrons: Discovered by James Chadwick in 1932.

Properties of Cathode Rays (Electrons)

• Travel in straight lines from cathode to anode.

• Negatively charged (deflected by electric and magnetic fields).

• Consist of particles much smaller than atoms (electrons).

Properties of Anode Rays (Protons)

• Travel from anode to cathode in a discharge tube.

• Positively charged particles (protons).

• Mass much greater than electrons.

Atomic Models

Thomson’s Atomic Model (Plum Pudding Model)

• Atoms are spheres of positive charge with electrons embedded within.

• Could not explain the arrangement of positive and negative charges.

Rutherford’s α-Particle Scattering Experiment

• Gold foil experiment showed that most α-particles passed through, but some were deflected at large angles.

• Concluded that atoms have a small, dense, positively charged nucleus.

Rutherford’s Atomic Model

• Atom consists mostly of empty space.

• Nucleus at the center contains protons (and later, neutrons).

• Electrons revolve around the nucleus in orbits.

Drawbacks: Could not explain the stability of atoms or the arrangement of electrons.

Bohr’s Atomic Model

• Electrons revolve in fixed orbits (energy levels) without radiating energy.

• Energy is absorbed or emitted only when an electron jumps between orbits.

Fundamental Particles: Properties and Comparison

Particle	Symbol	Charge (C)	Relative Mass	Location
Electron	e-	-1.602 × 10-19	1/1836	Outside nucleus
Proton	p+	+1.602 × 10-19	1	Inside nucleus
Neutron	n	0	1	Inside nucleus

Atomic Number, Mass Number, and Isotopes

• Atomic number (Z): Number of protons in the nucleus; defines the element.

• Mass number (A): Total number of protons and neutrons in the nucleus.

• Isotopes: Atoms of the same element with different numbers of neutrons.

Example: Carbon-12 and Carbon-14 are isotopes of carbon.

Electronic Configuration of Atoms

• Electrons are arranged in shells or energy levels around the nucleus.

• Maximum number of electrons in a shell is given by $2n^2$, where n is the shell number.

Shell	Symbol	Maximum number of electrons
1	K	2
2	L	8
3	M	18
4	N	32

Valence Shell and Valence Electrons

• Valence shell: Outermost shell of an atom.

• Valence electrons: Electrons in the valence shell; determine chemical properties and reactivity.

Example: Sodium (Na, Z = 11) has electronic configuration 2, 8, 1. The single electron in the outermost shell is the valence electron.

Sample Calculations and Examples

• Number of protons = Atomic number (Z)

• Number of neutrons = Mass number (A) - Atomic number (Z)

• Number of electrons = Number of protons (for a neutral atom)

Example: For an atom with Z = 17 and A = 35 (Chlorine): Number of protons = 17 Number of neutrons = 35 - 17 = 18 Number of electrons = 17

Key Equations

• Maximum electrons in a shell: $2n^2$

• Number of neutrons: $N = A - Z$

Summary Table: Atomic Structure Concepts

Concept	Description
Atomic Number (Z)	Number of protons in the nucleus
Mass Number (A)	Total number of protons and neutrons
Isotopes	Atoms with same Z but different A
Electronic Configuration	Arrangement of electrons in shells
Valence Electrons	Electrons in the outermost shell

Additional info:

• Practice questions and concept application exercises are included in the original material to reinforce understanding of atomic structure concepts.

• Tables and diagrams in the original notes illustrate the structure of atoms, subatomic particles, and electronic configurations.