Periodic Classification of Elements

Periodic Classification of Elements Dobereiner’s Triads Johann Wolfgang Dobereiner - 1817 ● The German scientist tried to arrange the elements with similar chemical properties. ● He identified some groups of 3 elements having the same chemical properties and thus called them triads. ● He showed that when the three elements in the triads are arranged in ascending order of atomic masses, the average of the atomic masses (A/no. of protons and neutrons) of the 1st and the 3rd element is roughly equal to the mass of the second element. Limitation He could only accommodate nine elements in the triads at that time and other elements couldn’t get a place. This law was discarded as a mere coincidence.

Newlands’ Law of Octaves John Newlands - 1866 ● The English Scientist John Newland tried to arrange the then-known elements in order of their increasing atomic masses. ● After he observed the arrangement he noticed that every 8th element had similar properties to its respective 1st element. ● He compared this to the octaves found in music, therefore, naming it as the “Law of Octaves”, now known as the “Newlands’ Law of Octaves Limitations 1. It was only applicable up to Ca after that the 8th element wasn’t similar to that of the first. 2. John assumed that only 56 elements existed in nature and no more wud be discovered in the future. But later on, several new elements were discovered and they didn’t fit into the “Law of Octaves”. 3. To fit the elements in the table he adjusted two elements with dissimilar properties under the same note. Co and Ni were placed in the same column as F, Cl, and Br whereas Fe which is similar to Ni and Co was placed far away from it.

Mendeleev’s Periodic Table Dmitri Ivanovich Mendeleev - 1869 ● The Russian Chemist examined the relationship between the physical and chemical properties of an element and its atomic mass. ● For chemical properties, he chose the formulas of hydrides and oxides as one of the basic properties as O2 and H2 form compounds with nearly every element. ● He then took 63 cards for the 63 elements which were discovered until then and wrote down properties of each element and arranged or pinned them in the ascending order of their atomic masses on a wall. ● He observed periodic recurrence of elements with similar chem and physical properties and thus formulated his Periodic Law. ● Mendeleev’s Periodic Law states “The properties of the elements are the periodic function of their atomic masses.”

Achievements 1. Mendeleev boldly predicted the existence of some of the elements that had not been discovered at that time by leaving gaps for the undiscovered elements to fill in. The properties listed by Mendeleev for the elements in the gap were highly similar to the later discovered elements. 2. The noble gases which were discovered much later could be placed in a new group without disturbing the existing order. Limitations 1. No fixed position could be provided to hydrogen. 2. Isotopes of the elements which have the same chemical properties but different atomic masses posed a challenge to the PT made by Mendeleev. 3. Atomic masses didn’t increase in a regular manner and thus it was not possible to predict no. of undiscovered elements between two discovered elements.

Modern Periodic Table Henry Moseley - 1913 ● Henry Moseley showed that the atomic number (Z/ no. of protons) of an element is a more fundamental property than atomic mass. ● Mendeleev’s PT was changed accordingly and atomic no. was adopted as the fundamental base of the Modern Periodic Table (MPT). ● Modern Periodic Law states “The properties of the elements are the periodic function of their atomic number.”.

TRENDS IN THE MODERN PERIODIC TABLE Valency ● Across a period the valency starts from 1 increases up to 4 and then decreases to 0. ● Down a group, the valency remains the same Atomic Radii ● Atomic radii decrease as we go along a period as there is an increase of nuclear charge which tends to pull the valence electrons closer to the nucleus of the cell.

● Atomic radii increases as we go down a group because the no. of shells increase as we go down a group. Metallic Properties or Effective Nuclear Charge ● Metallic properties decrease as we go across a period as the effective nuclear charge increases which reduces the tendency to lose electrons. [ELECTRONEGATIVITY INCREASES ACROSS A GROUP] ● Metallic properties increase as we go down a group as the atomic radii increases due to an increase in the no. of shells which further leads to low effective nuclear charge, thus, as we go down a group elements tend to lose electrons. [ELECTROPOSITIVITY INCREASES DOWN A GROUP]