Composition and Structure of the Atom and the Periodic System of Elements

Composition and Structure of the Atom

An atom is the fundamental building particle of matter, which is further indivisible by chemical means. It is composed of two primary parts: the nucleus and the electron envelope. The history of the atom begins with the Greek philosophers Democritus and Leucippus in the 5th century BC, who first proposed that matter consists of indivisible particles called atoms, derived from the Greek word "atomos," meaning indivisible. In 1803, British chemist John Dalton formulated the first atomic theory, which established that elements are composed of very small indivisible atoms. He stated that atoms of a single element are identical, while atoms of different elements differ in properties such as mass. During chemical reactions, atoms are combined, rearranged, or separated without changing into atoms of a different element, and they are neither created nor destroyed. Furthermore, Dalton's theory posited that molecules of chemical compounds are formed by the combination of atoms from two or more elements, with a specific compound always containing the same relative number of atoms for each element, such as in water, where there are always two hydrogen atoms for every one oxygen atom.

Historical Atomic Models

Following Dalton, the understanding of atomic structure evolved through several key models. In 1897, British physicist Joseph John Thomson proposed the first model, known as the Pudding Model. He suggested the atom was composed of a positively charged mass (the pudding) with negatively charged electrons distributed within it like raisins. In 1911, New Zealand physicist Ernest Rutherford introduced the Planetary Model after discovering the atomic nucleus and protons. He divided the atom into an extremely small nucleus with a diameter of approximately $10^{-14}\,m$, which contains 99.9% of the atom's mass, and a significantly larger electron envelope measuring around $10^{-10}\,m$ with negligible mass. In this model, negatively charged electrons move in circular orbits around the positively charged nucleus. However, this theory faced criticism because, according to classical physics, moving electrons should radiate electromagnetic waves, causing them to lose energy, decrease their orbital radius, and eventually be swallowed by the nucleus. Since atoms are stable, this model was incomplete.

In 1913, Danish physicist Niels Bohr proposed the Quantum-mechanical Model. He theorized that electrons orbit the nucleus only on predetermined stationary (stable) paths, which are circles with specific radii. In these orbits, electrons have constant energy and do not radiate electromagnetic waves. An electron can move to a higher energy path only if it receives a specific dose of energy called a quantum; conversely, it must radiate a quantum of energy when returning to a lower path. This model introduced the concept of quantization, showing that electrons exist only in states with certain energy levels and change energy in steps (quanta) rather than continuously. While Bohr viewed the electron as a small ball (corpuscle), his theory was limited to hydrogen and cations with a single electron, such as $He^{+}$ or $Li^{2+}$, and failed to explain other chemical phenomena like chemical bonding.

Finally, in 1923, Austrian physicist Erwin Schrödinger proposed the Wave-mechanical Model. This model was based on the experimentally confirmed hypothesis that electrons have a dualistic character, behaving simultaneously as particles (corpuscles) and as waves. Schrödinger’s model is mathematical and describes the atom through the Schrödinger equation, the solutions of which determine the shapes of orbitals.

The Atomic Nucleus and Elementary Particles

The study of the atomic nucleus is the focus of atomic physics. The nucleus consists of two types of elementary particles: protons and neutrons, which are collectively referred to as nucleons. The nucleus is approximately $10^{-14}\,m$ in size and accounts for 99.9% of the atom's total mass. The number of protons in the nucleus is defined by the proton number ($Z$), which also determines the number of electrons in a neutral atom and the element's position in the Periodic System of Elements (PSP). A proton is a positively charged particle with a charge of $Q_p = e = 1.6021 \times 10^{-19}\,C$ and a rest mass of $m_p = 1.6725 \times 10^{-27}\,kg$. The electron, while not a part of the nucleus, is a negatively charged particle in the envelope with a charge of $Q_e = -e = -1.6021 \times 10^{-19}\,C$ and a rest mass of $m_e = 9.11 \times 10^{-31}\,kg$, which is negligible compared to the nucleus. The neutron number ($N$) indicates the number of neutrons in the nucleus; neutrons are neutral particles with no charge ($Q_n = 0$) and a rest mass of $m_n = 1.6748 \times 10^{-27}\,kg$. The nucleon number ($A$) is the sum of protons and neutrons in the nucleus, expressed as $A = Z + N$.

Related to these numbers are the definitions of substance types. An element is a substance composed of atoms with the same proton number ($Z$), though they may have different neutron numbers. A nuclide is a substance composed of atoms that are absolutely identical, having the same number of protons and the same number of neutrons. Isotopes are atoms of the same element that differ only in their number of neutrons. For example, hydrogen exists as three nuclides: protium ($H_1^1$), deuterium (heavy hydrogen, $H_1^2$), and tritium ($H_1^3$). Most elements occur in nature as a mixture of isotopes. While isotopes have identical chemical properties due to having the same electron configuration, they possess slightly different physical properties. The elementary charge ($e$) is the smallest existing charge in nature, though partial charges ($\delta+$ and $\delta-$) can exist during bond shifts, albeit never independently of each other.

Stability of the Atomic Nucleus

The stability of an atomic nucleus is related to the mass defect ($\Delta m$), which is the difference between the theoretical rest mass of the nucleus (the sum of the masses of individual nucleons, $m_j' = Z \cdot m_p + (A - Z) \cdot m_n$) and the actual measured rest mass ($m_j$). The actual mass is always smaller than the theoretical mass: $\Delta m = m_j' - m_j$. This phenomenon is explained by Einstein’s Special Theory of Relativity (STR), which states that $\Delta E = \Delta m \cdot c^2$. When a stable nucleus forms from individual nucleons, mass is lost and energy is released. The binding energy of the nucleus ($E_v$) is the energy released during the formation of a nucleus from nucleons or the energy required to split it back into nucleons. The higher the binding energy, the more stable the nucleus. The most stable nuclei have nucleon numbers ($A$) between 30 and 130. The unit for this energy is the electronvolt ($eV$), where $1\,eV = 1.6 \times 10^{-19}\,J$. The binding energy per nucleon ($\epsilon = \frac{E_v}{A}$) serves as a measure of stability; a larger value indicates a more stable nucleus and more energy released during its formation.

History of the Periodic System of Elements

The organization of elements began in the 18th century when approximately 40 elements were known. Early attempts to use "atomic weights" (relative atomic masses) were often inaccurate. In 1829, German chemist Johann Wolfgang Döbereiner grouped elements with similar properties into triads, where the atomic weight of the middle element was the arithmetic average of the other two. In 1860, the adoption of Avogadro's constant led to more accurate atomic weights. In 1862, Alexandre Émile Béquyer de Chancourtois arranged 62 elements by increasing atomic weight in a three-dimensional "telluric screw" and introduced the concept of the period. In 1864, John Alexander Reina Newlands introduced ordinal numbers and the "law of octaves," noting that every eighth element had similar properties. William Odling also arranged elements by weight in 1864, including exceptions for tellurium and iodine and leaving gaps for undiscovered elements. Simultaneously, Julius Lothar Meyer arranged 28 elements based on valency (oxidation state).

In 1869, Russian chemist Dmitri Ivanovich Mendeleev discovered the Periodic Law, stating that the properties of elements and their compounds are a periodic function of their atomic weights. He predicted the existence and properties of undiscovered elements and corrected the atomic weights of others. By 1870, his table was titled the "Natural System of Elements." In the 20th century, W. Ramsay added noble gases, B. Brauner helped place lanthanides, and T. Seaborg placed actinides. Eventually, with the description of the electron envelope, the proton number replaced atomic weight as the primary organizational criterion. The modern Periodic Law states: "The properties of elements are a periodic function of their proton number."

Structure and Organization of the Periodic Table

The Periodic Table is a graphical representation of the periodic law. It consists of seven horizontal rows called periods ($n = 1$ to $n = 7$). Each period corresponds to the principal quantum number ($HK\check{C}$) of the highest occupied electron shell. Period 1 ($n=1$) contains 2 elements; periods 2 and 3 each contain 8; periods 4 and 5 contain 18 each; period 6 contains 32; and period 7 is incomplete, containing 26 elements. All periods except the last end with a noble gas.

There are 16 vertical columns known as groups. Groups are labeled with Roman numerals I-VIII and divided into A (main/non-transitional) and B (subsidiary/transitional) types, or more modernly numbered 1-18. Elements in the same group have similar properties because they share the same valence electron configuration in their outer shell. Specific group names include alkali koval (I.A), alkaline earth metals (II.A), triels (III.A), tetrels (IV.A), pentels (V.A), chalcogens (VI.A), halogens (VII.A), and noble gases (VIII.A). Group VIII.B includes the iron triad (Fe, Co, Ni), light platinum metals (Ru, Rh, Pd), and heavy platinum metals (Os, Ir, Pt). Lanthanides and actinides (internal transitional elements) consist of 14 elements each from the 6th and 7th periods, respectively.

Classification and Trends in the Periodic System

Elements are classified by their valence electrons into blocks: s-elements ($ns^{1-2}$, groups 1-2), p-elements ($ns^2\,np^{1-6}$, groups 13-18), d-elements (transitional koval, groups 3-12), and f-elements (internal transitional koval, lanthanides and actinides). Metals constitute 80% of elements and are characterized by metallic luster, conductivity, malleability, low electronegativity, and the formation of cations and bases. They are arranged in the Becketov series based on redox potential; those to the left of hydrogen react with acids to produce hydrogen, while noble metals to the right do not. Non-metals, found in the upper right, have high electronegativity, form anions and acids, and have oxidative effects. Metalloids like B, Si, Te, Sb, and As exhibit properties of both.

Periodic trends include changes in vertical, horizontal, and diagonal directions. Vertically down a group ($\downarrow$), electronegativity and non-metallic character decrease while metallic character, the ability to form cations, and base-forming character increase. This is because valence electrons are further from the nucleus and less strongly attracted. Horizontally across a period to the right ($\rightarrow$), electronegativity, non-metallic character, the ability to form anions, and acid-forming character increase, while metallic character and reducing abilities decrease. Diagonally toward the top right ($\nearrow$), electronegativity and non-metallic character increase, while toward the bottom left ($\swarrow$), metallic character and reducing abilities increase.