Periodic Table 

The periodic table arranges elements by increasing atomic number and repeating properties. Rows represent periods; columns represent groups. Elements in the same group have similar electron configurations and chemical behaviors. Within a period, valence electrons increase sequentially, correlating with rising energy levels and more energy sublevels per level.

1.0What Is a Periodic Table?

• The periodic table systematically organizes all 118 known elements by atomic number, indicating the number of protons in the nucleus. Each subsequent element increases by one atomic number as you move left to right.
• The table is named for the recurring appearance of elements with similar traits at consistent intervals, known as periods, organizing rows based on shared properties.
• The first 94 elements occur naturally, while those from 95 to 118 have been synthesized in labs or reactors. The modern periodic table evolved from earlier 19th and 20th-century models, with Dimitri Mendeleev credited for its development, building upon the work of predecessors like John Newlands and Antoine-Laurent de Lavoisier.

2.0Origin of Periodic Classification

• The modern periodic table, born from nearly a century of collaboration among scientists, organizes elements based on their atomic numbers and recurring properties. Lavoisier's categorization of elements as metals and nonmetals in 1789 and Döbereiner's triads in the early 19th century laid initial groundwork. 
• Mendeleev's predictive table, refined from Meyer's work, became iconic, while Newlands' "law of octaves" introduced an early attempt at classification by atomic mass.

3.0Dobereiner’s Triad Rule

In 1789, French chemist Antoine Lavoisier categorized elements as metals and nonmetals. Forty years later, German physicist Johann Wolfgang Döbereiner observed similarities among certain elements' physical and chemical properties. He grouped them in triads based on increasing atomic weight, noting that the properties of the middle element often approximated the average of the other two.

$X=\frac{35.5+127}{2}=81.2$ (Where, X = average atomic weight)  

$X=\frac{40+137}{2}=88.5$   (Where X = average atomic weight)

$X=\frac{7+39}{2}=23$    (Where X = average atomic weight)

4.0Newlands Octave Rule

In 1864, chemist John Newlands organized 62 known elements by ascending atomic weights, noticing a pattern: every eighth element shared properties. This led to the Newland Law of Octaves, demonstrating the orderly nature of chemistry. For example, sodium, the eighth element after lithium, shares properties with lithium and potassium. Similarly, chlorine, the eighth element following fluorine, shares properties with fluorine.

Newlands Octave Rule (1865):

(i) Elements were organized by atomic mass, with every eighth element resembling the first, akin to musical octaves.

(ii) Inert gases were unknown at that time.

(iii) Lithium properties resemble sodium, beryllium properties resemble magnesium, and so forth.

5.0Mendeleev's Periodic Table

• Mendeleev's Periodic Table was a monumental scientific achievement, representing the first systematic organization of all known elements. Dmitri Mendeleev's visionary brilliance and unwavering dedication to understanding the chemical world led to this groundbreaking milestone.
• Mendeleev also considered chemical properties, prioritizing the formulae of hydrides and oxides for categorization.
• His ingenious arrangement positioned elements with akin characteristics in vertical columns, which he dubbed "groups" within the periodic table. He termed horizontal rows "periods." This structuring ensured that elements with similar properties assembled, forming a cohesive framework for understanding the underlying principles of chemistry.

6.0What is Modern Periodic Law?

A Russian chemist, Dmitri Mendeleev, observed a pattern among the elements: as their atomic masses increased, their chemical and physical properties repeated those of preceding elements.

This principle became known as Mendeleev's Periodic Law. However, with only 63 known elements, this law was only universally applicable to some elements.

To refine Mendeleev's Periodic Law, Henry Moseley proposed the Modern Periodic Law, which states:

"The properties of elements are a periodic function of their atomic numbers, rather than their atomic masses." 

The modern periodic law, attributed to Henry Moseley, is also known as Moseley's Law.

7.0Modern Periodic Table

The modern periodic table, refined by Bohr, organizes elements by atomic number, not mass. It has 18 groups (1-18) and 7 periods (1-7). Groups have similar outer shell configurations, revealing chemical similarities. In Periods, elements are grouped by electron shell count, aiding comprehension of atomic structure. This system enables the prediction of element properties and behaviors.

• We classify elements in the periodic table in rows, columns, and blocks that are discussed below,

In Periods and Groups

Elements are categorized into 7 periods and 18 groups, with elements in each group arranged based on the electrons in their outermost shell, and elements in each period organized according to the number of electron shells they possess.

The number of elements in each period:

• The first period has 2 elements, Hydrogen and Helium.
• The second period has 8 elements, from Lithium to Neon.
• The third period has 8 elements, from Sodium to Argon.
• The fourth period has 18 elements, from Potassium to Krypton.
• The fifth period has 18 elements, from Rubidium to Xenon.
• The sixth period has 32 elements.
• The seventh period is incomplete.

In Blocks

The Modern Periodic Table is divided into four blocks that are as follows:

• s-Block
• p-Block
• d-Block
• f-Block

8.0Electronic Configurations Of Elements

An element's position in the periodic table corresponds to the quantum numbers of its last filled orbit. This section explores the direct correlation between elements' electronic configurations and their placement within the extended format of the Periodic Table.

In Periods

In periods, the value of n in the electronic configuration corresponds to the outermost or valence shell. Each successive period in the Periodic Table signifies the filling of the next higher principal energy level (n = 1, n = 2, etc.). Notably, the number of elements in each period is twice the number of atomic orbitals available in the energy level being filled.

In Groups 

Elements within a vertical column or group share akin valence shell electronic configurations, possessing identical numbers of electrons in their outer orbitals, leading to analogous properties. For instance, the Group 1 elements (alkali metals) uniformly display an ns¹ valence shell electronic configuration.

9.0Periodic Trends In Properties Of Elements

Atomic Radius

Atomic radius measures the size of an atom from its nucleus to the outermost electron. Across a period, it decreases left to right; down a group, it increases. This concept helps determine the covalent radius for nonmetals and the metallic radius for metals based on proton count and electron-proton attraction. There are three types of atomic radii: Metallic, Covalent, and Van der Waals'.

Atomic radii are influenced by :

• Number of shells: An increase in the number of shells corresponds to an increase in atomic radii.
• Effective nuclear charge (Z_eff): Atomic radius decreases as the effective nuclear charge within an atom increases.
• Screening effect.: Greater screening effect from electrons in filled orbitals results in a larger atomic radius.

Ionic Radius

Ionic radius measures the distance between the nucleus's center and the outermost electron in an ion. It diminishes from left to right across a period and enlarges from top to bottom within a group in the periodic table. Typically, anions exhibit larger ionic radii compared to cations.

Ionization Energy

Ionization energy represents the minimum energy needed to extract an electron from an isolated gaseous atom in its ground state.

Ionization energy escalates as one progresses from left to right across a period and diminishes when descending a group. The magnitude of ionization energy significantly influences the nature of chemical bonds and molecular geometry. Essentially, ionization energy quantifies an element's propensity to relinquish an electron.

The lower the ionization energy, the more readily a neutral atom can transform into a positive ion.

Factors Influencing Ionization Energy:

• Atomic Size
• Effective Nuclear Charge (Z_eff)
• Screening Effect
• Penetration Effect
• Electronic Configuration

Electron Gain Enthalpy

• Electron gain enthalpy is the change in enthalpy when an electron is added to a neutral gaseous atom, forming a gaseous anion. Positive values indicate energy absorption, while negative values indicate energy release. Positive enthalpy means decreased stability, while negative enthalpy means increased stability. It increases from left to right and decreases down a group in the periodic table.

Factors that influence electron gain enthalpy include:

• Atomic size
• Effective nuclear charge (how strongly the nucleus attracts electrons)
• Screening effect (how much inner electrons shield outer electrons)
• Electronic configuration (arrangement of electrons in the atom)

Electronegativity 

Electronegativity refers to an atom's ability within a molecule to pull the shared pair of electrons closer to itself. It's not a fixed property but varies depending on the element to which it's bonded. Electronegativity aids in predicting the nature of the bond established between two atoms.

Factors Influencing Electronegativity:

• Atomic Size
• Effective Nuclear Charge (Z_eff)
• Magnitude of Positive Charge on the Atom