Early Atomic Models

John Dalton's model, created in the early 19th century, introduced the idea that matter is made up of small particles called atoms. He suggested that these atoms are indivisible, meaning they cannot be broken down into smaller parts. According to Dalton, atoms of the same element are identical in mass and properties, which implies that all carbon atoms are alike, and all oxygen atoms are alike. Chemical reactions, in this model, simply involve the rearrangement of these atoms.

However, Dalton's model has its limitations. For example, it did not take into account the existence of subatomic particles, which are the smaller units within an atom, like protons, neutrons, and electrons. Also, it could not explain observations made in experiments involving electricity or spectroscopy, which study the interaction of light and matter.

J. J. Thomson built upon Dalton's model after discovering the electron in 1897 during experiments with cathode rays. He proposed a new model called the 'plum-pudding' model, where he envisioned atoms as a positively charged sphere with negatively charged electrons embedded within it, akin to plums scattered in a pudding. This design allowed for an electrically neutral atom overall, balancing the positive and negative charges.

However, like Dalton's model, Thomson's plum-pudding model had limitations. It was unable to accurately describe results from later experiments, particularly those conducted by Ernest Rutherford, who found that atoms had a much more complex structure, including a dense nucleus that Thomson's model did not account for.

In 1911, Rutherford conducted an experiment where he fired alpha particles at a thin foil made of gold. He observed that while most of the particles passed through the foil with little disturbance, some were deflected at large angles. From this, he concluded that most of the atom is empty space, and the atom's positive charge and most of its mass lie in a small, dense core called the nucleus. This was a significant shift from previous models that could not account for such observations.

Key features of Rutherford's model include the nucleus, which is incredibly small compared to the entire atom, containing protons and neutrons, and the electron cloud surrounding it, where electrons exist but do not follow fixed orbits.

Niels Bohr, in 1913, expanded upon Rutherford’s model by adding the concept of quantized orbits for electrons in a hydrogen atom. He proposed that rather than moving randomly, electrons occupy fixed paths or orbits around the nucleus, and can only exist in certain allowed states without radiating energy. This means that as long as the electron stays in a particular orbit, it doesn't lose energy, which explains the stability of atoms. When electrons move between these orbits, they either emit or absorb energy in the form of photons, which manifests as light or other electromagnetic radiation. Bohr's calculations worked well for hydrogen and predicted its emissions accurately, but the model became limited when applied to more complex atoms with multiple electrons.

In 1926, Erwin Schrödinger revolutionized the understanding of atomic structure by introducing a wave-based model for electrons. His approach treated electrons not as particles moving in fixed orbits, as Bohr suggested, but as wavefunctions that describe probabilities of finding an electron in a given position around the nucleus. Schrödinger's equation provides a way to calculate these wavefunctions, which lead to quantized energy levels similar to Bohr's but also included new quantum numbers to describe the shape and orientation of these orbitals. The concept of orbitals represents regions of space where there is a high probability of finding an electron, challenging the simpler circular paths described in Bohr's model.