Dall'atomismo al modello nucleare dell'atomo

Historical Roots of Atomism and the Philosophical Precursors

The study of the fundamental building blocks of matter began long before the advent of modern scientific methodology. Historically, atomism originated as a philosophical and metaphysical doctrine suggesting that reality is composed of atoms—extremely small, indivisible particles in eternal motion. According to this view, the aggregation and disaggregation of these particles give rise to different bodies, their distinct properties, and the nature of their change. This philosophical tradition can be traced back to ancient Greece with figures such as Leucippus and Democritus ($460$ – $370$ B.C.). It was further developed by Epicurus ($341$ – $270$ B.C.) and preserved for the Latin-speaking world through the poetic work De Rerum Natura by Titus Lucretius Carus ($96$ – $53$ B.C.).

In this classical atomistic framework, the universe consists strictly of matter (the atoms) and the void within which those atoms move. By combining in an nearly infinite number of ways, atoms produce the physical reality we observe. The term "atom" itself is derived from the Greek word for "indivisible." In a modern context, while we now distinguish between pure substances, elements, and compounds, the atom remains defined as the smallest identifiable unit of an element. However, for millennia, a critical question remained unanswered: what exactly is an atom, and what is its physical structure? Over twenty centuries passed before the philosophy of atomism transitioned into a rigorous scientific theory based on experimental observation.

The Scientific Foundation and Dalton's Atomic Theory

By the early 19th century, specifically in $1808$, John Dalton unified several experimental observations into what is known as Dalton's Atomic Theory. This theory was built upon three fundamental laws formulated shortly before his time. The first was the Law of Conservation of Mass, or Lavoisier's Law ($1789$), which states that matter is neither created nor destroyed in a chemical reaction. Experimental evidence showed that the total mass of reactants (typically solids or liquids) equals the total mass of the products. While some processes like combustion initially seem to lose mass (e.g., wood weighing more than ash), this is an illusion caused by the escape of gases like $CO_2$ and water vapor. If these gases are captured and weighed, the total mass remains constant.

The second foundation was the Law of Definite Proportions, or Proust's Law ($1797$), which asserts that all samples of a given compound contain the same proportions of their constituent elements by mass, regardless of their source or preparation. For instance, water always maintains an atomic ratio of $2$ to $1$ for hydrogen and oxygen. In terms of mass, every $18-g$ of water contains $16-g$ of oxygen and $2-g$ of hydrogen, meaning there is $8$ times more oxygen than hydrogen by mass. This principle implies that pure substances are identical whether they are natural or synthetic; for example, vanillin extracted from a plant is chemically indistinguishable from laboratory-synthesized vanillin.

The third foundation, the Law of Multiple Proportions, was announced by Dalton himself in $1804$. It explains that when two elements, $A$ and $B$, form more than one compound, the masses of element $B$ that combine with a fixed mass of $A$ are in ratios of small whole numbers. Consider carbon monoxide ($CO$) and carbon dioxide ($CO_2$). If $1-g$ of carbon reacts with $1.33-g$ of oxygen to form $CO$, it will react with exactly $2.67-g$ of oxygen (double the amount) to form $CO_2$. The ratio of these oxygen masses is $2.67 / 1.33 = 2$. Dalton's Atomic Theory summarized these laws with four main postulates: 1) Each element is composed of extremely small particles called atoms. 2) All atoms of a given element are identical in mass and properties, but different from atoms of other elements. 3) Chemical reactions do not transform atoms of one element into another; they are neither created nor destroyed. 4) Compounds form when atoms of different elements combine in fixed, definite proportions.

The Discovery of the Electron and Thomson's Model

Despite the success of Dalton's theory, it did not define the internal structure of the atom; atoms were still treated as undefined "building blocks," which could hypothetically be any shape as long as they were tiny. This changed in the mid-1800s with the observation of cathode rays. By creating a high vacuum in a glass tube and applying a high potential difference between two electrodes, physicists noticed a radiation originating from the negative electrode (the cathode) and moving toward the positive electrode (the anode). These rays were invisible but could be detected with a fluorescent screen.

J.J. Thomson demonstrated that these cathode rays were sensitive to electromagnetic fields. By applying voltage to plates or using a magnet, he could deflect the beam, proving the rays consisted of negatively charged particles, now called electrons. Thomson succeeded in calculating the charge-to-mass ratio ($e/m$) for these particles. In $1909$, Robert Millikan conducted the oil-drop experiment to determine the specific charge of an electron. By suspending charged oil drops between metal plates using an electric field, he found that the charge on any drop was always a multiple of a fundamental unit: $-1.6 \times 10^{-19}\,C$. Combining this with Thomson's ratio, the mass of an electron was calculated to be approximately $2000$ times smaller than that of the lightest atom (hydrogen), identifying the electron as a subatomic particle.

Concurrent with these discoveries was the identification of radioactivity. Henri Becquerel observed spontaneous high-energy emission from uranium, and Marie and Pierre Curie identified the uranium atom as the source. Ernest Rutherford later characterized these emissions into three types based on their behavior between charged plates: beta ($\beta$) rays, which are high-energy electrons with a negative charge; gamma ($\gamma$) rays, which are high-energy electromagnetic radiation with no charge; and alpha ($\alpha$) rays, which consist of positively charged particles with a charge of $+2$ (later identified as helium-4 nuclei). Based on the discovery of the tiny, negative electron, Thomson proposed the "Plum Pudding" model in $1900$. He envisioned the atom as a sphere of positive charge with electrons embedded throughout it, similar to raisins in a panettone cake.

Rutherford's Experiment and the Nuclear Model

Ernest Rutherford overturned Thomson's model through his famous gold foil experiment. He bombarded a very thin sheet of gold with positively charged alpha ($\alpha$) particles. According to the Plum Pudding model, the alpha particles should have passed through the gold foil with only minor deflections (less than one degree). However, Rutherford observed that while most particles passed through, some were deflected at very large angles, and a few even bounced back toward the source.

To explain these results, Rutherford introduced the nuclear model of the atom. He concluded that the atom's mass and positive charge are concentrated in a very small region called the nucleus. The rest of the atom is essentially empty space through which the electrons move. The discovery was revolutionary because it suggested that the "void" identified by ancient philosophers was not just between atoms, but inside them. To provide a sense of scale, if an atom was expanded to a diameter of $100\,m$, its nucleus would only be about $1\,mm$ wide. Despite appearing solid, atoms are mostly empty space and are not dense in the way previously imagined. The typical size of an atom is measured in Angstroms, where $1\,Å = 10^{-10}\,m$.

Modern Atomic Structure and Subatomic Particles

The modern view of the atom identifies three primary subatomic particles: protons, neutrons, and electrons. Protons and neutrons are located in the nucleus, while electrons occupy the vast space surrounding it. Using the fundamental charge unit ($1.6 \times 10^{-19}\,C$), electrons are assigned a charge of $-1$ and protons a charge of $+1$. Neutrons, as the name suggests, carry no net electrical charge. For an atom to be electrically neutral, it must possess an equal number of protons and electrons.

Nearly all of an atom's mass is concentrated in the nucleus, as a proton is approximately $2000$ times heavier than an electron. Consequently, the density of the nucleus is staggering, estimated between $10^{13}$ and $10^{14}\,g/cm^3$. To visualize this, a single cubic centimeter of matter with the density of a nucleus would weigh roughly $10,000,000,000\,kg$. However, atoms themselves are not "heavy" in absolute terms; the mass of a single atom is on the order of $10^{-27}\,kg$. To simplify measurements, scientists use the atomic mass unit (amu), also known as the Dalton ($Da$). One amu is defined as $1.66054 \times 10^{-27}\,kg$. Under this system, the mass of both a proton and a neutron is approximately $1\,amu$, providing a convenient scale for chemical and biological calculations.