Periodic classification of elements

Periodic Classification of Elements

The periodic classification of elements refers to the organization of chemical elements into a tabular form known as the Periodic Table. This table arranges elements based on their atomic number, electron configurations, and recurring chemical properties.

History of the Periodic Table

1. Dmitri Mendeleev (1869):

   • Mendeleev is credited with creating the first widely recognized periodic table. He arranged elements in order of increasing atomic weight and grouped elements with similar properties together.
   • Mendeleev left gaps for elements that had not yet been discovered and predicted their properties.

2. Henry Moseley (1913):

   • Moseley discovered the concept of atomic number, which led to the reordering of the periodic table by atomic number instead of atomic weight.

Structure of the Periodic Table

1. Periods:

   • Horizontal rows in the periodic table. There are 7 periods.
   • Elements in the same period have the same number of electron shells.

2. Groups:

   • Vertical columns in the periodic table. There are 18 groups.
   • Elements in the same group have similar chemical properties and the same number of valence electrons.

Classification of Elements

1. Metals:

   • Located on the left and middle parts of the periodic table.
   • Properties: Good conductors of heat and electricity, malleable, ductile, and have a shiny appearance.
   • Examples: Sodium (Na), Iron (Fe), Gold (Au).

2. Non-metals:

   • Located on the right side of the periodic table.
   • Properties: Poor conductors of heat and electricity, brittle, and have a dull appearance.
   • Examples: Oxygen (O), Carbon (C), Sulfur (S).

3. Metalloids:

   • Elements that have properties intermediate between metals and non-metals.
   • Located along the staircase line starting from Boron (B) to Polonium (Po).
   • Examples: Silicon (Si), Germanium (Ge).

Blocks of the Periodic Table

1. s-block:

   • Groups 1 and 2, including Hydrogen and Helium.
   • Elements have their outermost electron in an s-orbital.

2. p-block:

   • Groups 13 to 18.
   • Elements have their outermost electron in a p-orbital.

3. d-block:

   • Groups 3 to 12 (transition metals).
   • Elements have their outermost electron in a d-orbital.

4. f-block:

   • Lanthanides and Actinides.
   • Elements have their outermost electron in an f-orbital.

Periodic Trends

1. Atomic Radius:

   • Decreases across a period from left to right due to increasing nuclear charge.
   • Increases down a group due to the addition of electron shells.

2. Ionization Energy:

   • Increases across a period from left to right due to increasing nuclear charge.
   • Decreases down a group due to increased distance of valence electrons from the nucleus.

3. Electronegativity:

   • Increases across a period from left to right.
   • Decreases down a group.

4. Electron Affinity:

   • Generally becomes more negative across a period.
   • Varies down a group.