Fundamental Concepts of Atomic Structure and Chemical Bonding

Posted on Aug 17, 2025 in Chemistry

Subatomic Particles and Their Discoverers

  • Electron: Discovered by J.J. Thomson in 1897.
    • Mass: Approximately 1/1837 of a proton’s mass.
    • Mass (kg): 9.1 x 10-31 kg.
    • Charge: -1 (relative charge).
  • Proton: Discovered by E. Rutherford in 1920.
    • Mass (atomic mass unit, U): 1 U.
    • Mass (kg): 1.7 x 10-27 kg.
  • Neutron: Discovered by J. Chadwick in 1932.
    • Mass (atomic mass unit, U): 1 U.
    • Mass (kg): 1.7 x 10-27 kg.
    • Charge: 0 (neutral).

Models of the Atom

Rutherford’s Planetary Model

The planetary model of the atom was proposed by E. Rutherford. In this model, the nucleus, made up of protons and neutrons, is surrounded by a cloud of orbiting electrons.

Atomic Number and Mass Number

  • The atomic number (Z) of an atom is the number of protons it contains.
  • The mass number (A) of an atom is the total number of protons and neutrons it contains.

Ions

When an atom loses or gains electrons, it becomes an ion. A positively charged ion is called a cation, while a negatively charged ion is called an anion.

Isotopes

In 1913, J.J. Thomson discovered that there were two types of neon atoms differing in mass. He called these “isotopes,” meaning “the same place.” Isotopes of a given element have the same number of protons but a different number of neutrons.

Electronic Configuration

Bohr’s Model of Electron Arrangement

In 1913, to explain the nature of light emitted by atoms, Niels Bohr put forward the idea that electrons were arranged in energy levels around the nucleus according to their energy.

  • The maximum number of electrons an ‘n’ level can accommodate is given by the formula 2n2 (e.g., 8 electrons in the second energy level).
  • Electronic configuration shows how electrons are arranged in levels around the nucleus.

Representing Electronic Configuration

  • The nucleus is shown as a black spot.
  • Each energy level is shown as a circle around the nucleus.
  • Each electron is shown by a dot or a cross.

The Modern Periodic Table

The modern periodic table displays all the elements in horizontal rows called periods and vertical columns called groups. Elements are arranged in order of increasing atomic number.

Types of Elements

  • Metals: Typically shiny, solid at room temperature, and good conductors of electricity.
  • Non-metals: Not shiny, often gases at standard temperature and pressure, and poor conductors of electricity.
  • Metalloids: Possess properties intermediate between metals and non-metals.

Periodic Table Properties

  1. All elements in a group have the same number of electrons in their outer shell. These are called valence electrons, and they dictate how an element reacts.
  2. The group number is generally the same as the number of outer electrons, except for Group 0 (or 18), where the atoms have full shells.
  3. The period number indicates how many electron shells an atom has.

Common Elements and Their Atomic Numbers

  • H Hydrogen (1)
  • He Helium (2)
  • Li Lithium (3)
  • Be Beryllium (4)
  • B Boron (5)
  • C Carbon (6)
  • N Nitrogen (7)
  • O Oxygen (8)
  • F Fluorine (9)
  • Ne Neon (10)
  • Na Sodium (11)
  • Mg Magnesium (12)
  • Al Aluminium (13)
  • Si Silicon (14)
  • P Phosphorus (15)
  • S Sulphur (16)
  • Cl Chlorine (17)
  • Ar Argon (18)
  • K Potassium (19)
  • Ca Calcium (20)
  • Fe Iron (26)
  • Cu Copper (29)
  • Zn Zinc (30)
  • Br Bromine (35)
  • Ag Silver (47)
  • Sn Tin (50)
  • I Iodine (53)
  • Au Gold (79)
  • Hg Mercury (80)
  • Pb Lead (82)
  • U Uranium (92)

Periodic Table Groups

  • Group 1A (Alkali Metals): Located on the left side of the periodic table, they form alkaline compounds.
  • Group 2A (Alkaline Earth Metals): Less reactive than Group 1A metals.
  • Groups 3B-2B (Transition Metals): Include most of the metals.
  • Group 7A (Halogens): Highly reactive at room temperature.
  • Group 0 (Noble Gases): Have completely full electron shells and are generally unreactive, forming stable compounds.

Nuclear Radioactivity

The nucleus is the very small core of an atom, containing the positive charge. It is made up of nucleons (protons and neutrons).

Protons and neutrons are bound together in the nucleus by the strong nuclear force.

Radioactive Decay and Radioisotopes

Radioactive decay is the process by which unstable atoms (radioisotopes) decompose, emitting energetic radiation and changing the nucleus. Radioactive decay is a random process.

Types of Radioactive Emissions

There are three kinds of radioactive emissions:

  • Alpha (α) particles: Energetic helium nuclei.
  • Beta (β) particles: Fast electrons.
  • Gamma (γ) rays: High-energy electromagnetic radiation.

Chemical Bonding

Group 0 elements (noble gases) are unreactive or stable and exist as single atoms. Their electron configuration shows an outer electron shell completely filled with electrons, which brings about stability.

Atoms can bond in different ways:

  1. Ionic Bonding
  2. Covalent Bonding
  3. Metallic Bonding

Ionic Bonding: Metals and Non-metals Reacting

When metal atoms react with non-metal atoms, the metal atom loses electrons and becomes a cation (positively charged ion). The non-metal atom accepts those electrons and becomes an anion (negatively charged ion).

Many cations and anions formed this way cluster together so that each ion is surrounded by the maximum number of ions of opposite charge that their size ratio allows, forming a giant 3-D structure of ions.

Properties of Ionic Substances

  • Ionic compounds have high melting and boiling points, requiring a lot of heat energy to break up the lattice structure. Ionic compounds are solid at room temperature.
  • Ionic crystals are hard but brittle.
  • Ionic compounds are poor conductors of heat and electricity in their solid state.
  • Ionic compounds are usually soluble in water.

Covalent Bonding

When non-metal atoms react together, they need to gain electrons to reach full shells. They achieve this stability by sharing electrons, forming a covalent bond.