Structure of the atom

Subatomic particles

Electron, proton, and neutron are the basic subatomic particles.

Discovery of electron

J.J. Thomson discovered the electron using cathode ray tube experiments.

Charge of electron

Robert Millikan determined the charge of an electron through the oil drop experiment: -1.602 × 10⁻¹⁹ C.

Mass of electron

9.1 × 10⁻³¹ kg.

Cathode rays

Streams of electrons moving from cathode to anode in a discharge tube under low pressure and high voltage.

Anode rays

Positive rays or canal rays discovered by Goldstein, consisting of protons.

Discovery of proton

E. Goldstein discovered protons through canal rays; charge: +1.602 × 10⁻¹⁹ C, mass: 1.672 × 10⁻²⁷ kg.

Neutron

James Chadwick discovered neutron; neutral particle, mass ≈ proton, no charge.

Atomic number (Z)

Number of protons in an atom; defines the element.

Mass number (A)

Sum of protons and neutrons in the nucleus.

Isotopes

Atoms of same element with same Z but different A (e.g., ¹H, ²H, ³H).

Isobars

Atoms with same mass number but different atomic number (e.g., ⁴⁰Ca and ⁴⁰Ar).

Thomson's model

Plum pudding model; atom is a positively charged sphere with electrons embedded like raisins.

Limitations of Thomson's model

Couldn't explain stability and results of alpha scattering experiment.

Rutherford's gold foil experiment

Alpha particles bombarded on thin gold foil; most passed through, few deflected, very few rebounded.

Conclusions from Rutherford's experiment

Atom has a small dense positively charged nucleus, rest is empty space, electrons orbit nucleus.

Rutherford's model limitations

Couldn't explain stability of atom and spectra of hydrogen.

Bohr's model (postulates)

1. Electrons revolve in fixed orbits (energy levels). 2. Energy is emitted/absorbed only when electron jumps between orbits.

Energy level

Fixed circular path around nucleus where electron revolves without radiating energy.

Quantum

Smallest packet of energy absorbed or emitted during transition of electron.

Formula for energy emitted/absorbed

ΔE = E₂ - E₁ = hν, where h = Planck's constant, ν = frequency.

Planck's constant (h)

6.626 × 10⁻³⁴ Js.

Hydrogen spectrum

Series of lines due to transitions of electron between orbits; e.g., Lyman, Balmer, Paschen series.

Balmer series

Visible region; electron transition to n=2 from higher levels.

Limitations of Bohr's model

Works only for hydrogen, doesn't explain Zeeman effect, Stark effect, or fine structure.

de Broglie hypothesis

Matter shows wave-like behavior; λ = h/p = h/(mv).

Heisenberg uncertainty principle

Δx × Δp ≥ h/4π; can't simultaneously know exact position and momentum of a particle.

Quantum mechanical model

Uses wave equations to predict probability distribution of electron (developed by Schrödinger).

Atomic orbital

Region in space with high probability of finding electron; defined by quantum numbers.

Principal quantum number (n)

Indicates main energy level or shell; n = 1, 2, 3…; affects size and energy of orbital.

Azimuthal quantum number (l)

Indicates shape of orbital; l = 0 to n-1; s, p, d, f represent l = 0, 1, 2, 3 respectively.

Magnetic quantum number (m)

Indicates orientation of orbital in space; m = -l to +l.

Spin quantum number (s)

Indicates spin of electron; values = +½ or -½.

s orbital

Shape: spherical; l = 0.

p orbital

Shape: dumbbell; l = 1; 3 orientations (m = -1, 0, +1).

d orbital

Shape: cloverleaf; l = 2; 5 orientations (m = -2 to +2).

Pauli Exclusion Principle

No two electrons in an atom can have the same set of four quantum numbers.

Aufbau Principle

Electrons fill orbitals of lowest energy first.

Hund's Rule

Electrons occupy degenerate orbitals singly with parallel spins before pairing.

Electronic configuration

Arrangement of electrons in orbitals; e.g., carbon: 1s² 2s² 2p².

Orbital vs orbit

Orbit: circular path (Bohr model); Orbital: 3D region with high electron probability.

Nodes

Regions with zero probability of finding an electron; number of radial nodes = n - l - 1.

Shape of orbitals

Depends on azimuthal quantum number: s = sphere, p = dumbbell, d = clover.

Subshell

Defined by l; group of orbitals with same energy level and shape.

Shell

Defined by n; collection of subshells.

Number of orbitals in shell n

Total orbitals = n².

Maximum electrons in shell n

Total electrons = 2n².

Electronic configuration notation

Use nl^x format; e.g., 1s² 2s² 2p⁶.

Stability of half-filled and fully-filled orbitals

Extra stability due to symmetry and exchange energy; e.g., Cr: [Ar] 3d⁵ 4s¹ not 3d⁴ 4s².

Valence electrons

Electrons in outermost shell; determine chemical reactivity.