How Periodic Trends Relate To Periodic Law

Periodic trends are unique patterns in chemical element properties revealed by the periodic table of elements. Electronegativity, ionisation energy, electron affinity, atomic radii, ionic radius, metallic character, and chemical reactivity are all major periodic patterns.

Periodic trends are caused by changes in chemical elements’ atomic structure within their respective periods and groups (horizontal rows and vertical columns) in the periodic table. Based on their atomic structures and properties, these principles allow chemical elements to be grouped in the periodic table. The periodic trends allow for a partial inference of any element’s unknown attributes.

There are a few outliers, such as the ionisation energy trend of group 3, the electron affinity trend of group 17, the density trend of group 1 elements (alkali metals), and so on.

How periodic trends relate to periodic law

The periodic trends are based on the Periodic Law, which asserts that if the chemical elements are enumerated in increasing atomic number order, many of their properties undergo cyclical changes, with elements with comparable attributes recurring at regular intervals. Many of the physical and chemical features of lithium, such as its intense reactivity with water, recur in sodium, potassium, and cesium after arranging elements in increasing atomic numbers.

After a number of nineteenth-century scientific experiments, Russian chemist Dmitri Mendeleev developed this idea in 1871. An elemental periodic system was also devised by Mendeleev, who proposed an elemental periodic system that included not only an atomic weight, but also chemical and physical qualities. Henry Moseley found in 1913 that periodicity is dictated by the atomic number, not the atomic weight. Lothar Meyer’s table was presented a few months after Mendeleev’s, although he disagreed with Mendeleev’s Periodic law. Initially, there was no scientific explanation for the Periodic Law, and it was simply utilised as an empirical principle; but, with the discovery of quantum physics, it became possible to grasp the theoretical basis for the Periodic Law.

When elements are listed in increasing atomic number order, the periodic recurrence of elements with comparable physical and chemical properties emerges directly from the periodic recurrence of identical electronic configurations in respective atoms’ outer shells.

The Periodic Law’s discovery was a watershed moment in chemistry’s evolutionary trajectory. For example, the Periodic Law had an impact on the development of the periodic table.

Radius of the atom

The atomic radius is the distance between the atomic nucleus and the atom’s outermost stable electron orbital. Because rising effective nuclear force on electrons causes the atom to shrink, it tends to decrease during a period from left to right. Because of the inclusion of a new energy level, the atomic radius normally increases as one moves down a group (shell). However, because the quantity of electrons has a greater effect than the sizable nucleus, atomic radii tend to increase diagonally. Lithium, for example, has a smaller atomic radius (145 picometer) than magnesium (150 picometer).

 The atomic radius can be classified in four ways:

• Covalent radius is half the distance between two singly bound atoms in a diatomic molecule.
• Van der Waals radius: half the distance between the nuclei of various molecules in a covalent lattice.
• Metallic radius is defined as half the distance between two adjacent atom nuclei in a metallic lattice.
• Ionic radius is defined as half the distance between two nuclei of an ionic compound.

The energy of ionisation

The ionisation potential is the amount of energy necessary to remove one electron from each isolated, neutral, and gaseous atom in a mole. The first ionisation energy is the energy required to remove the first electron, and the nth ionisation energy is the energy required to remove the atom’s nth electron after the initial (n1) electrons have been removed. Ionisation energy tends to grow over time because a greater quantity of protons (higher nuclear charge) pulls the circling electrons more strongly, increasing the energy required to remove one of the electrons. Ionisation energy and ionisation potentials are not the same thing. Potential is an intensive property that is measured in “volts,” whereas energy is an extensive attribute that is represented in “eV” or “kJ/mol.”

Electron attraction

The energy released by an atom when an electron is added to it, or the energy required to detach an electron from a singly charged anion, can be used to explain an atom’s electron affinity. The sign of the electron affinity can be somewhat perplexing, because atoms that become more stable with the addition of an electron (and hence have a higher electron affinity) show a reduction in potential energy, implying that the energy gained by the atom appears to be negative. The atom’s electron affinity is positive in this circumstance. Potential energy increases for atoms that become less stable after adding an electron, implying that the atom obtains energy. The atom’s electron affinity is negative in this sequence  . However, if electron affinity is defined as the energy required to remove an electron from an anion, the energy value obtained will be of the same magnitude but with the opposite sign. This is due to the fact that atoms with a high electron affinity are less likely to give up an electron, requiring more energy to remove the electron from the atom. In this, the atom with the higher positive energy value has a greater attraction for electrons. The electron affinity increases as one moves from left to right over a period.

Electronegativity

Electronegativity is a measure of an atom’s or molecule’s ability to attract pairs of electrons in the setting of a chemical bond.

Using the Pauling scale, the difference in electronegativity between the atoms involved determines the sort of connection formed. The electronegativity increases as one proceeds from left to right through a period in the periodic table due to the stronger attraction that the atoms receive as the nuclear charge increases. As one moves down the group, the electronegativity drops due to an increase in the distance between the nucleus and the valence electron shell, reducing the atom’s attraction to electrons.

However, electronegativity increases from aluminium to thallium in group (iii) elements.

Fluorine is the element with the highest electronegativity.

Electrons with valence

The electrons in the outermost electron shell of an isolated atom of an element are known as valence electrons. As we move from left to right in a period, the number of valence electrons increases. In a group, however, this periodic tendency is continuous, i.e. the number of valence electrons remains constant.

Valency

Across a period, the periodic table’s valency increases and subsequently decreases. There is no difference when you go down a group.This periodic tendency, however, is less common for heavier elements (elements with atomic numbers greater than 20), particularly the lanthanide and actinide family.

The more core electrons there are, the more electrons are shielded from the nucleus’s core charge. As a result, ionisation energy is lower for elements lower in a group, while species polarizability is higher for elements lower in a group. The valency does not alter as one moves down a group because the core electrons have no effect on the bonding behaviour. Non-bonding interactions, such as the ones mentioned above, are, however, influenced by core electrons.

Properties of metallic and non-metallic materials

Metallic characteristics often grow down groups as diminishing affinity between nuclei and outermost electrons causes these electrons to be less tightly bonded and therefore capable of conducting heat and electricity. The decreasing metallic nature is caused by the rising attraction between the nuclei and the outermost electrons over each cycle, from left to right.

Non-metallic character, on the other hand, reduces down groups and grows over time.

The majority of metals are shiny (when freshly broken, polished, or prepared), ductile, malleable, and sonorous, whereas the majority of nonmetals are not.

CONCLUSION

From the following article we can conclude that The periodic trends are based on the Periodic Law, which asserts that if the chemical elements are enumerated in increasing atomic number order, many of their properties undergo cyclical changes, with elements with comparable attributes reoccurring at regular intervals. Periodic trends are unique patterns in chemical element properties revealed by the periodic table of elements. Electronegativity, ionisation energy, electron affinity, atomic radii, ionic radius, metallic character, and chemical reactivity are all major periodic patterns.