Fundamentals of Atomic Structure and Periodic Trends

Foundational Concepts of Atomic Structure

Discovery of Subatomic Particles

• Electron: Discovered by J. J. Thomson (1897). Relative mass: 1/1837. Mass: 9.1 x 10⁻³¹ kg. Charge: -1.
• Proton: Discovered by Ernest Rutherford (1920). Relative mass: 1 atomic mass unit (u). Mass: 1.7 x 10⁻²⁷ kg. Charge: +1.
• Neutron: Discovered by James Chadwick (1932). Relative mass: 1 atomic mass unit (u). Mass: 1.7 x 10⁻²⁷ kg. Charge: 0.

Rutherford's Planetary Model (1911)

Ernest Rutherford proposed the Planetary Model, which is accepted today as the basic structure of the atom.

The atom consists of a very small central region, the nucleus, containing protons and neutrons, surrounded by orbiting electrons. The number of electrons equals the number of protons in the nucleus, resulting in atoms having no net electrical charge.

The nucleus is approximately 100,000 times smaller than the atom but contains virtually all of the atomic mass.

Key Atomic Definitions

• Atomic Number (Z): The number of protons an atom contains. Atoms of a particular element always have the same atomic number. Atoms of different elements have different numbers of protons.
• Mass Number (A): The total number of protons and neutrons contained in the nucleus. (Note: Mass number is often written at the top, and atomic number at the bottom, in isotopic notation.)

Ions and Isotopes

• Ion: Formed when an atom loses or gains electrons.
  • A positive ion is called a Cation (loss of electrons).
  • A negative ion is called an Anion (gain of electrons).
• Isotopes: In 1913, J. J. Thomson discovered two types of neon atoms differing in mass. Most elements consist of a mixture of isotopes in constant proportions. Isotopes of a given element have the same number of protons but different numbers of neutrons.

Electron Configuration and Energy Levels

In 1913, Niels Bohr proposed the idea that electrons are arranged in layers (shells) around the nucleus according to their energy. The outer layers possess higher energy.

The maximum number of electrons (e⁻) in any principal energy level (n) is calculated by the formula: 2n².

The electron configuration of an atom shows precisely how electrons are arranged in these energy levels around the nucleus.

The Periodic Table of Elements

Structure and Classification

The Periodic Table organizes over 100 elements based on their differing physical and chemical properties. Elements are arranged in increasing order of atomic number.

• Horizontal rows are called Periods.
• Vertical columns are called Groups.

Three Main Classes of Elements

1. Metals: Typically shiny, solid at room temperature, and good conductors of electricity.
2. Non-metals: Lack shine, often gases or brittle solids, and are poor conductors of electricity.
3. Metalloids: Possess properties intermediate between metals and non-metals.

Relationship to Atomic Structure

The position of an element in the table relates directly to its atomic structure:

• All elements in the same group have the same number of outer shell electrons (valence electrons). Valence electrons dictate how an element reacts chemically.
• The Group number corresponds to the number of outer electrons (except for Group 0/18).
• The Period number indicates how many electron shells the atom possesses.

Key Groups in the Periodic Table

Group 1A: Alkali Metals

Located on the far left of the table. They are called alkali metals because they form alkaline compounds. They have one electron in their highest energy level, making them extremely reactive. They are often stored in oil to prevent reaction with air or water.

Group 2A: Alkaline Earth Metals

Reactivity generally increases as you move down this group.

Groups 3B–2B: Transition Metals

These are metals used in everyday applications. They are typically hard, have high density, are good conductors, and are malleable and ductile. They are known for forming a great number of colored compounds.

Group 7A: Halogens (Non-metals)

These non-metals are highly reactive at room temperature. Reactivity decreases as you move down the group. Some elements in this group are dangerous and have historically been used as chemical weapons.

Group 0 (or 18): Noble Gases

Their electron shells are full, meaning they are highly stable and generally do not react to form stable compounds. They exist as individual atoms. They have limited uses (e.g., in airships and light bulbs).