Atomic Models and Structure

The Development of Atomic Models

Thomson's Model (1897), the "Plum Pudding"

J.J. Thomson discovered the electron. He proposed that an atom was a sphere of uniformly distributed positive charge with electrons embedded throughout it, like raisins in a pudding. This model predicted that alpha particles fired at atoms would pass through or be deflected by only small amounts.

The Geiger-Marsden Experiment (1909)

Rutherford directed Geiger and Marsden to fire alpha particles at a thin gold foil and detect where they landed:

• Expected (if Thomson correct): almost all particles pass straight through; minor small-angle scattering.
• Observed: most particles did pass straight through; but some were deflected through large angles, and a few bounced almost directly back.

Rutherford famously said it was "as if you fired a 15-inch shell at tissue paper and it came back and hit you."

Conclusion: most of the atom is empty space; positive charge and most of the mass are concentrated in a tiny, dense nucleus. Electrons orbit the nucleus in the surrounding empty space.

Rutherford's Nuclear Model (1911)

Rutherford proposed the atom as a miniature solar system: a tiny, dense, positively charged nucleus at the centre, with electrons orbiting it like planets around the Sun. The nucleus diameter is about $10^{-4}$ times the atom diameter.

Bohr's Model (1913)

Niels Bohr modified Rutherford's model by proposing that electrons occupy specific energy levels called shells or orbits. Electrons in a shell have a fixed energy and do not radiate energy as long as they remain in that shell. When an electron jumps to a lower shell, it emits energy as light of a specific frequency.

Chadwick's Discovery of the Neutron (1932)

James Chadwick confirmed the existence of the neutron, an uncharged particle with nearly the same mass as a proton, inside the nucleus. This explained why nuclear masses were always greater than the number of protons alone would predict.

Structure of the Atom

The atom is electrically neutral: number of protons = number of electrons.

Nuclide Notation

${}^{A}_{Z}\text{X}$

where:

• $\text{X}$ is the chemical symbol of the element.
• $Z$ is the atomic number (proton number), number of protons.
• $A$ is the mass number (nucleon number), total number of protons + neutrons.

Number of neutrons: $N = A - Z$.

Example: ${}^{40}_{20}\text{Ca}$ has 20 protons, 20 neutrons, and (as a neutral atom) 20 electrons.

Isotopes

Isotopes are atoms of the same element (same $Z$) with different numbers of neutrons (different $A$). Isotopes have identical chemical properties but different nuclear properties.

Examples:

• Hydrogen: ${}^{1}_{1}\text{H}$, ${}^{2}_{1}\text{H}$ (deuterium), ${}^{3}_{1}\text{H}$ (tritium).
• Carbon: ${}^{12}_{6}\text{C}$ (stable), ${}^{14}_{6}\text{C}$ (radioactive, used in dating).
• Uranium: ${}^{235}_{92}\text{U}$ (fissile) and ${}^{238}_{92}\text{U}$.

The Electron Shell Model

Electrons are arranged in shells around the nucleus. Each shell can hold a maximum number of electrons:

The first two shells give electronic configurations for elements up to calcium. For example, sodium (Na, $Z = 11$): 2, 8, 1, the outermost electron is responsible for sodium's chemistry.

FLASHCARD SET • 6 CARDSAtomic Models and Structure: Quick Recall