History of the Periodic Table and Atomic Properties
How to Identify Elements
In 1830, John Jacob Berzelius, a Swedish chemist (1779-1848), proposed a method to represent the elements: using the first letter in Latin or, if two or more elements had the same initial, the initial letter followed by another present in the Latin name. For example, N for nitrogen, Na for sodium, and Ni for nickel.
Early Grouping of Elements
Many studies early this century established that the elements could be grouped into families with similar chemical properties, such as sodium and potassium, or chlorine, bromine, and iodine. The two properties most researched by scientists at the time to characterize a new element were: the atomic weight (a physical property now known as relative atomic mass or Ar) and the valence (a chemical property, numerically expressing the combining capacity of atoms, now known as oxidation number).
Dobereiner’s Proposal
In 1817, Johann Dobereiner (1780-1849) observed that the atomic weight of strontium was very close to the arithmetic mean of the atomic weights of calcium and barium, and all three are chemically similar elements were grouped into one family. In 1829, he established the same regularity of the atomic weights for several sets of three items that he called triads, where the atomic weight of the central element of the triad was almost equal to the average of the other two. For example: Chlorine (35.47) – Bromine (79.916) – Iodine (126.91) — Average: 81.18
Newlands’ Octaves
In 1864, John R. Newlands (1837-1898) ordered the then-known elements in increasing order of atomic weights and observed that the properties of elements repeated in periods of seven, in a similar way as do the musical notes in the octave of a piano keyboard. The properties of the eighth item in a series were similar to those of the first, so these periods of seven elements were called Newlands’ octaves.
Meyer and Mendeleev’s Work
Between 1868 and 1870, the work of Lothar Meyer (1830-1895) in Germany and Dmitri Mendeleev (1834-1907) in Russia led to the discovery of the periodic law of the chemical elements. Meyer ordered previously known elements in increasing order of atomic weights and related this scale to another: the atomic volume. Representing the atomic volumes of the elements according to atomic weights, Meyer found that the chart formed a series of peaks, corresponding to groups of elements with similar properties: lithium, sodium, potassium, rubidium, and cesium. He also determined that each peak, with its ups and downs, was a period of elements.
Mendeleev’s Periodic Table
In 1869, Mendeleev published the first edition of the periodic table, ordering the 63 elements then known. Once ordered by atomic weight, Mendeleev studied their chemical properties, especially in terms of their valences. He noted that the first elements of the list showed a progressive change in their valence, with increasing and decreasing values. He established periods: the first single for hydrogen, the next two with seven items each, and the others with more than seven items. To coincide properties, Mendeleev did not hesitate to relocate some items. He also left empty spaces to form groups of items with the same properties and predicted, with uncanny accuracy, the properties of the elements that would occupy those vacancies once discovered. From Mendeleev’s work, the periodic law of elements was established.
Electronic configuration of atoms of these elements ends the same way, in general:
- The atoms of elements belonging to a group have the same outer electron configuration (EEC).
- By contrast, when analyzing the electronic configuration of atoms of the elements located in the same period, it shows that they have the same number of energy levels.
According to this structure in groups and periods, the table is divided into 4 main blocks: S, P, D, and F, whichever is the last occupied orbital of the EEC.
- The building blocks S and P correspond to the elements, comprising representative metals and nonmetals. Some are metalloids, such as silicon or arsenic.
- D block elements are called transition elements and are all metals.
- The F block is composed of inner transition elements, which are also metals, most obtained by artificial synthesis.
Effective Nuclear Charge: The electrons that are closest to the core (core electrons) have a screening effect on the positive charge of the nucleus (Z). For this reason, the outermost electrons are attracted to the core with a smaller force; the net charge is called the effective nuclear charge (Zeff).
Atomic Radius: According to the quantum mechanical model, the electron density distribution in an atom does not have a clearly defined limit. However, if the atom is considered a sphere, the distance between the outer electron and the nucleus can be determined experimentally. This distance is called the atomic radius.
Ionic Radii: When neutral atoms lose or gain electrons, they become ions. Cations are formed by losing electrons and have a net positive charge, while anions are formed by gaining electrons and become negatively charged. The size of a cation is smaller than the corresponding neutral atom; however, the size of an anion is greater than the neutral atom source.
Ionization Energy: Ionization energy is the energy required to remove an electron from a neutral atom, in the gaseous state, and in its ground state. The atom becomes a positive single ion (i.e., with a single positive charge).
Electron Affinity (EA): Electron affinity is the energy exchanged when a neutral atom, in the gaseous state and in its ground state, captures an electron and becomes a single negative ion, i.e., with a single negative charge.
Electronegativity (EN): Electronegativity is the relative ability of an atom to attract electrons to itself when in a chemical bond with another atom.