Structure of an Atom
The atom is the fundamental unit of matter, consisting of a dense nucleus surrounded
by an electron cloud. The nucleus contains protons (positively charged) and
neutrons (neutral), which account for nearly all the atom's mass. Electrons
(negatively charged) orbit the nucleus in energy levels or shells, contributing
negligible mass but determining chemical behaviour.
1. Quantum Mechanical Model (Electron Cloud Model):
o Electrons do not follow fixed orbits but exist in orbitals (regions with high
probability of finding an electron).
o The Schrödinger equation provides wave functions (ψ) that describe
electron behaviour.
o The square of the wave function (|ψ|²) gives the probability density of
finding an electron at a particular location.
Energy Levels & Subshells:
Electrons occupy discrete energy levels (n = 1, 2, 3, ...), each with sublevels (s,
p, d, f).
Each sublevel has orbitals with unique shapes:
o s-orbital (spherical)
o p-orbital (dumbbell-shaped)
o d & f orbitals (more complex shapes)
How the Electron Cloud Spreads Around the Nucleus:
Lower energy levels (e.g., n=1) are closer to the nucleus, with a more
concentrated cloud.
Higher energy levels (e.g., n=2, 3,…) have orbitals that spread out farther,
increasing the atom’s size.
Isotopes
Isotopes are atoms of the same element (same atomic number) with different numbers
of neutrons (different mass numbers). They have similar chemical properties (due to
identical electron configurations) but differ in physical properties like stability or
nuclear behaviour.
1. Hydrogen Isotopes (Atomic Number = 1)
Protium (¹H)
o Protons: 1, Neutrons: 0, Mass Number: 1
Deuterium (²H or D)
o Protons: 1, Neutrons: 1, Mass Number: 2
o Used in nuclear reactors and heavy water (D₂O).
Tritium (³H or T)
o Protons: 1, Neutrons: 2, Mass Number: 3.
2. Carbon Isotopes (Atomic Number = 6)
Carbon-12 (¹²C)
o Protons: 6, Neutrons: 6, Mass Number: 12
Carbon-13 (¹³C)
o Protons: 6, Neutrons: 7, Mass Number: 13
Carbon-14 (¹⁴C)
o Protons: 6, Neutrons: 8, Mass Number: 14.
3. Oxygen Isotopes (Atomic Number = 8)
Oxygen-16 (¹⁶O)
o Protons: 8, Neutrons: 8, Mass Number: 16
Oxygen-17 (¹⁷O)
o Protons: 8, Neutrons: 9, Mass Number: 17.
Oxygen-18 (¹⁸O)
o Protons: 8, Neutrons: 10, Mass Number: 18
Isotopes have identical chemical properties (same protons & electrons)
but different physical properties (mass, stability).
Some isotopes are stable, while others are radioactive (unstable) and decay
over time.
Applications include nuclear energy, medicine (tracers), archaeology
(dating), and climate science.
Mass Number and Atomic Number
Atomic Number (Z): The number of protons in an atom's nucleus, defining the element's identity. It
also equals the number of electrons in a neutral atom.
Mass Number (A): The total number of protons and neutrons in the nucleus, reflecting the atom's
mass (electrons contribute negligible mass).
Calculation: Number of neutrons = A - Z.
Nuance: The atomic number determines the element’s position in the periodic table, while the mass
number varies among isotopes of the same element.
Example: Carbon-12 has Z = 6 (6 protons) and A = 12 (6 protons + 6 neutrons). Carbon-14 has Z =
6 but A = 14 (6 protons + 8 neutrons)
Octet Rule
The octet rule states that atoms tend to achieve a stable electron configuration by having eight electrons in
their valence (outermost) shell, resembling the electron configuration of noble gases. Atoms achieve this
through ionic bonding (electron transfer) or covalent bonding (electron sharing).
⁺Ionic: Sodium (Na) loses one electron to form Na⁺, achieving a neon-like configuration. Chlorine (Cl) gains
one electron to form Cl⁻, also achieving a neon-like configuration, forming NaCl.
Covalent: In methane (CH₄), carbon shares four electrons with four hydrogens, achieving an octet, while
each hydrogen achieves a duet.
Covalent bonding involves the sharing of electron pairs between atoms to achieve a stable electron
configuration, typically between non-metals with similar electronegativities. The shared electrons occupy
overlapping orbitals, forming a molecular bond.
Mechanism:
o Atoms share electrons in their valence orbitals (usually s or p) to form a bond.
o Single, double, or triple bonds involve sharing one, two, or three electron pairs, respectively.
o The electron cloud of the shared pair is concentrated between the nuclei, stabilizing the
molecule.
Orbital Structure:
o Bonding occurs via the overlap of atomic orbitals (e.g., s-s, s-p, p-p).
o The overlap forms molecular orbitals: bonding (σ or π, lower energy) and antibonding
(higher energy).
o Sigma (σ) bonds form from head-on overlap (e.g., s-s or p-p along the bond axis), while pi
(π) bonds form from sideways p-orbital overlap (e.g., in double/triple bonds).
Examples of Ionic (Electrovalent) Compounds
1. Sodium Chloride (NaCl) – Table salt (Na⁺ + Cl⁻)
2. Magnesium Oxide (MgO) – Ionic solid (Mg²⁺ + O²⁻)
3. Calcium Chloride (CaCl₂) – Used in de-icing roads
4. Potassium Bromide (KBr) – Used in photography
Examples of Covalent Compounds
1. Water (H₂O) – Polar covalent bond
2. Methane (CH₄) – non-polar covalent bond
3. Carbon Dioxide (CO₂) – Linear covalent molecule
4. Glucose (C₆H₁₂O₆) – Organic covalent compound
Electrovalent (Ionic) Bonding
Electrovalent bonding occurs when atoms achieve a stable electron configuration (often an octet) by
transferring electrons, forming oppositely charged ions that are held together by electrostatic forces. This
typically happens between a metal (low electronegativity) and a non-metal (high electronegativity).
Thomson’s Atomic Model
Evidence:
Based on cathode ray experiments, which discovered electrons.
Limitations:
Could not explain atomic stability (why electrons don’t collapse into the
positive sphere).
Failed to explain spectral lines or Rutherford’s scattering results.
Rutherford’s Atomic Model
Evidence:
Gold Foil Experiment (α-particle scattering):
o Most α-particles passed straight through (empty space).
o Some deflected at large angles (due to the nucleus).
Limitations:
Could not explain why electrons don’t spiral into the nucleus (classical
physics predicts energy loss via radiation).
Did not explain electron energy levels or atomic spectra.
Bohr’s Atomic Model
Evidence:
Explained hydrogen’s line spectrum (Balmer & Lyman series).
Mathematically predicted electron energy levels.
Limitations:
Only worked for hydrogen (failed for multi-electron atoms).
Did not explain electron wave nature (later addressed by quantum
mechanics).
