Atomic Structure – Chemistry Notes 🧪

An atom is the smallest particle of an element that can exist independently and retain the chemical properties of that element.

Atomic theory has evolved over time. Key contributors:

• Dalton (1803): Atom is indivisible and indestructible; all atoms of an element are identical.
• Thomson (1897): Electrons exist in a positively charged "plum pudding" model.
• Rutherford (1911): Atom has a small dense nucleus (positive) with electrons orbiting.
• Bohr (1913): Electrons occupy fixed energy levels (shells) around the nucleus.

• Proton (p⁺): Positive charge (+1), mass ≈ 1 amu, found in nucleus.
• Neutron (n⁰): Neutral charge (0), mass ≈ 1 amu, found in nucleus.
• Electron (e⁻): Negative charge (-1), mass ≈ 1/1836 amu, orbits nucleus.

Atomic Number (Z): Number of protons in the nucleus of an atom. Determines the element.

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

Example: Carbon-12 → Z = 6 (protons), A = 12 (protons + neutrons), neutrons = 12 - 6 = 6.

Isotopes are atoms of the same element with the same atomic number but different mass numbers.

Example: Hydrogen has three isotopes:

• Protium (¹H): 1 proton, 0 neutrons
• Deuterium (²H): 1 proton, 1 neutron
• Tritium (³H): 1 proton, 2 neutrons

Electron configuration is the arrangement of electrons in the energy levels (shells) of an atom.

Rule: Maximum electrons in shell n = 2n²

Example: Oxygen (Z=8): 2, 6 (2 electrons in K shell, 6 in L shell)