The philosophical notion that matter is composed of indivisible particles, or "atomos," dates back to ancient Greek philosophers like Democritus. However, these were largely conceptual ideas lacking empirical support. The true scientific pursuit of the atom began much later, laying the groundwork for modern chemistry.

In the early 19th century, John Dalton, an English chemist, revolutionized chemical thought with his atomic theory, built upon careful observations of chemical reactions. He proposed that all matter is composed of extremely small, solid, and indestructible particles called atoms, which he famously envisioned as miniature, uniform billiard balls. According to Dalton, atoms of a given element were absolutely identical in every respect, including their mass and chemical properties, making them indistinguishable from one another. Conversely, he asserted that atoms of different elements possessed unique masses and distinct properties, accounting for the diversity of substances in the world. Crucially, Dalton's theory posited that chemical reactions were not about the creation or destruction of atoms, but rather about their simple rearrangement and combination in fixed, whole-number ratios to form new compounds. This provided a foundational explanation for the law of conservation of mass and the law of definite proportions in chemical reactions.

However, Dalton's "indivisible" atom was soon challenged by new experimental evidence. Towards the end of the 19th century, J.J. Thomson, a British physicist, conducted groundbreaking experiments using cathode ray tubes. He observed that these rays were deflected by electric and magnetic fields, indicating they were composed of negatively charged particles. Through meticulous measurements, Thomson determined that these particles were far smaller and lighter than any known atom, leading to the astonishing discovery of the electron. This discovery fundamentally undermined Dalton's concept of an indivisible atom, demonstrating that atoms themselves had internal structure.

The most significant shift in our understanding of atomic structure came with Ernest Rutherford, a New Zealand-born physicist, and his associates in the early 20th century. Their iconic gold foil experiment, conducted in 1911, involved firing a beam of positively charged alpha particles at an extremely thin sheet of gold foil. Based on Thomson's plum pudding model, they expected the alpha particles to pass straight through with only minor deflections, as the positive charge was thought to be spread out. However, the experimental results were truly astonishing and defied expectations. While the vast majority of alpha particles did indeed pass straight through the foil, a small but significant fraction were deflected at very large angles, and, remarkably, a tiny percentage even bounced directly back towards the source. Rutherford famously likened this to firing a cannonball at a piece of tissue paper and having it bounce back. This unexpected scattering led Rutherford to a revolutionary conclusion: the atom's positive charge and nearly all of its mass were concentrated in an extraordinarily tiny, dense central region, which he named the nucleus.

While Rutherford's model elegantly accounted for the nucleus and the vast emptiness of the atom, it still had limitations. Classical physics predicted that orbiting electrons, being charged particles, should continuously lose energy and spiral inwards, eventually collapsing into the nucleus. This clearly contradicted the observed stability of atoms. In 1913, Niels Bohr, a Danish physicist, provided a crucial refinement to Rutherford's model by incorporating ideas from quantum theory. Bohr proposed that electrons do not orbit randomly but exist in specific, fixed energy levels or "shells" around the nucleus, similar to the rungs of a ladder. Each energy level corresponds to a distinct amount of energy, and electrons can only occupy these allowed energy levels without radiating energy. An electron can move from a lower energy level to a higher one by absorbing a precise amount of energy (a quantum), and conversely, it can fall from a higher energy level to a lower one by emitting a specific amount of energy, often in the form of light.