Periodicity in Properties-II

Classification of elements and periodicity in properties: A Brief Overview

The 19th century saw the discovery of so many elements that it was impossible to study them individually; thus, they had to be classified. Consequently, the elements were meticulously organised, and periodicity in properties became the focus of all chemical discovery and innovation. The Periodic Law states that if the chemical elements are listed in increasing order of atomic number, their properties change cyclically, with similar elements recurring periodically. 

Periodic trends result from the variations in the atomic structure of the chemical elements within their respective groups (vertical columns) or periods (horizontal rows) of the periodic table. 

According to these laws, chemical elements can be classified into periodic tables depending on their properties and atomic structures. Periodic trends can help determine the unknown characteristics of elements; however, there are exceptions, such as the trend in the group 3 ionisation energy trend, the group 17 electron affinity trend, the group 1 element density trend (alkali metals), etc.

A German chemist named Johann Wolfgang Dobereiner organised elements with similar properties into three groups, called triads. Dobereiner proposed that the atomic mass of the triad’s central part should be approximately equal to the mean of the atomic masses of the triad’s other two elements.

Lithium, sodium, and potassium make up this type of triad. Lithium has an atomic mass of 6.94, and potassium has an atomic mass of 39.10. The atomic mass of sodium, the middle element in this triad, is 22.99, roughly equal to the mean of the atomic masses of lithium and potassium (which is 23.02).

In 1866, an English scientist named John Newlands arranged the 56 known elements based on their increasing atomic mass. He noticed a pattern in which every eighth element looked similar to the first. 

According to Newland’s Law of Octaves, the periodicity in properties of two elements separated by a seven-element interval is similar when the elements are arranged in increasing order of atomic mass.

In 1869, Dmitri Ivanovich Mendeleev, a Russian chemist, proposed the periodic table. He discovered that the physical and chemical properties of elements were regularly related to the atomic mass of the elements.

According to the Periodic Law (also known as Mendeleev’s Law), the chemical properties of elements are a periodic function of their atomic weights. This principle underpins the modern periodic table.

Dmitri Mendeleev, however, is considered to be the most important contributor to the modern periodic table. The Periodic Law is named in honor of Mendeleev, the “Father of the Modern Periodic Table.”

Periodic properties

According to the Periodic Law, when chemical elements appear in increasing atomic number order, many of their properties change cyclically, with identical elements returning at regular intervals.

After grouping elements in increasing atomic number order, many of lithium’s physical and chemical features, such as its strong reactivity with water, may be found in sodium, potassium, and caesium.

In 1913, Henry Moseley found that periodicity depends on atomic number, not atomic weight. After several months, Julius Lothar Meyer, who disagreed with Mendeleev’s Periodic rule, arranged the elements based on their atomic weight, with the same valency, arranged in vertical columns, showing a striking similarity to Mendeleev’s table.

At first, the Periodic Law had no theoretical foundation and existed purely as an empirical idea. However, as quantum mechanics progressed, the academic backing for the Periodic Law became obvious.

When elements appear in the order of increasing atomic number, the periodic occurrence of elements with similar physical and chemical properties results directly from the periodic event of identical electronic configurations in respective atoms’ outer shells.

The discovery of the Periodic Law was one of the most meaningful events in the history of chemical science. The Periodic Law has been and continues to be used by almost every chemist.

It also led to the development of the periodic table, which is used across a wide range of professions. Electronegativity, electron affinity, ionisation energy, atomic radii, metallic character, ionic radius, and chemical reactivity can all be classified as major periodic patterns.

The major periodic trends:

• Atomic Radius

The distance between an atom’s nucleus and its outermost stable electron orbital is its atomic radius when it is in equilibrium. Atomic radius tends to decrease from left to right because the atom shrinks as the effective nuclear force on the electrons rises. Atomic radius typically grows as one progresses through a group because adding a new energy level (shell) lowers the size of the atoms with time.

• Ionisation energy

Ionisation potential is the energy required to remove a single electron from an isolated, gaseous, and neutral atom in a molecule. The energy required to remove the first electron is the first ionisation energy. Additionally, the energy that eliminates the atom’s nth electron after the first (n-1) electrons have been removed is known as the nth ionisation energy.

• Electronegativity

Electronegativity refers to an atom’s or molecule’s ability to attract pairs of electrons in the context of a chemical bond. Electronegativity increases because of the stronger attraction that the atoms obtain as the nuclear charge increases as one moves from left to right across a period in the periodic table.

• Valency

The periodic table’s valency increases and then decreases over time. Thus, when you move down a group, nothing changes. However, this periodic trend is less common for heavier elements (those with an atomic number greater than 20), particularly the lanthanide and actinide series.

• Electron Affinity 

Electron affinity is the amount of energy released when an electrons is added to an isolated, gaseous, and neutral atom .

Conclusion:

Only a few elements were known before the 18th century. However, there are 118 known elements at present; thus, studying and remembering their properties can be difficult. Due to the discovery of many elements in the 19th century and the examination of their properties, classification of elements and periodicity in properties became necessary. The repetition of the same valency cell electronic configuration after a definite gap of atomic numbers (magic numbers) such as 2, 8, 8, 18, 18, 32 causes periodicity in the properties of elements placed in any group.