Evolution of Atomic Models: From Dalton to Quantum Mechanics

Dalton’s Atomic Model

In 1808, Dalton formulated his atomic theory, which broke with traditional ideas (Democritus, Leucippus). It introduces the concept of the discontinuity of matter, being the first scientific theory to consider that matter is divided into atoms. The basic postulates of this theory are:

• Matter is divided into particles called atoms, which are indivisible and unchangeable.
• Atoms are very small particles and cannot be seen with the naked eye.
• All atoms of the same element are equal in mass and properties.
• Atoms of different elements have different masses and different properties.
• Compounds are formed when atoms join together in a constant and simple ratio.
• In chemical reactions, atoms separate or combine, but no atoms are created or destroyed, and no atom of one element becomes an atom of another element.

This view was maintained for nearly a century.

Plum Pudding Model

After the discovery of the electron (discovered by Thomson in 1897, see subatomic particles), in 1898, Thomson proposed an atomic model that considered the existence of this subatomic particle.

His model was static, meaning electrons were at rest within the atom, which was electrically neutral. Thomson’s model was similar to a fruitcake: electrons were embedded in a spherical mass of positive charge. The total negative charge of the electrons was equal to the total positive charge of the sphere, resulting in a neutral atom.

Thomson also explained the formation of ions, both positive and negative. When an atom loses an electron, it becomes positively charged, forming a positive ion. If an atom gains an electron, it becomes negatively charged, forming a negative ion.

Rutherford Model

After the discovery of the proton (a discovery to which Rutherford contributed, see subatomic particles), Rutherford developed his atomic model.

In 1911, Rutherford used alpha particles to determine the internal structure of matter (gold foil experiment, see left column, 1). From this experiment, he concluded that:

• Most particles pass through the foil undeflected (99.9%).
• Some particles are deflected (0.1%).

As Thomson’s model was unsatisfactory, Rutherford developed the nuclear model of the atom. In this model, the atom consists of a nucleus and a surrounding cloud:

• Nucleus: Almost the entire mass of the atom is concentrated here, and it has a positive charge.
• Cloud: Consists of electrons orbiting the nucleus in circular orbits (like a miniature solar system).

He also stated that matter is neutral because the positively charged nucleus and the negatively charged cloud neutralize each other.

Rutherford’s Conclusions:

The atom is mostly empty space; the nucleus is 100,000 times smaller than the atom’s radius. Most alpha particles do not deviate because they pass through the cloud, not the nucleus. Those passing near the nucleus are deflected due to repulsion. When an atom loses electrons, it becomes negatively charged, forming a negative ion. Conversely, if an atom gains electrons, it becomes positively charged. The atom is stable.

Bohr Model

After the discovery of the neutron (see subatomic particles) in 1913, Bohr tried to improve Rutherford’s model by applying Planck’s quantum ideas (see Quantum Theory). To make his atomic model work for the hydrogen atom, he described it as having one proton in the nucleus and an electron revolving around it. The new ideas on energy quantization are:

• The atom is quantized, meaning it can only have specific energy levels.
• The electron moves in circular orbits around the nucleus, and each orbit is a stable state associated with a natural number, “n” (principal quantum number), with values from 1 to 7.
• Each “n” layer is composed of various sub-levels, “l”. These are further divided into others (Zeeman effect), “m”. Finally, there is a fourth quantum number, “s”, related to spin.
• The allowed energy levels are multiples of Planck’s constant.
• When an electron moves from one energy level to another, it absorbs or emits energy. When the electron is in n=1, it is at the ground state (minimum energy level). If the electron absorbs energy, it changes levels and becomes an excited electron.

Bohr placed electrons in specific locations in space.

Quantum-Mechanical Model

This is the current model, proposed in 1925 by Heisenberg and Schrödinger. Its characteristic features are:

• Wave-particle duality: De Broglie proposed that material particles have wave properties and that any moving particle has an associated wave.
• Uncertainty Principle: Heisenberg stated that it is impossible to determine both the exact position and momentum of an electron simultaneously.

The equations of the quantum-mechanical model describe the behavior of electrons inside the atom, incorporating their wave character and the inability to predict exact trajectories.

This led to the concept of the orbital (see left column, 2), the region of space around the atom where the probability of finding an electron is very high.

Characteristics of orbitals:

• Energy is quantized.
• Unlike the Bohr model, this model does not determine the exact position of the electron, but rather the probability distribution.
• Within the atom, the electron is interpreted as a cloud of negative charge. The probability of finding an electron is higher where the cloud density is greater.

The behavior of electrons inside the atom is described by quantum numbers (see left column, 3).

Quantum numbers determine the behavior of electrons and the electron configuration (see left column, 4) and their distribution.

Given the number of elements, a classification system was needed. Today, we use the periodic table, although many other proposals were made before it. In the periodic table, elements are classified according to their atomic number.