JEE Chemistry Periodic Table Complete Guide

The Periodic Table is one of the most fundamental tools in Chemistry, especially for JEE aspirants. It organizes all known chemical elements in a systematic way based on their atomic number, electronic configuration, and recurring chemical properties. Understanding the periodic table thoroughly helps you predict element behavior, trends, and reactions, which is vital for scoring well in the JEE Chemistry section.

1. History of the Periodic Table

The periodic table's evolution is a story of scientific discovery:

• Dobereiner’s Triads (1817): Grouped elements in sets of three with similar properties and average atomic masses.
• Newlands’ Law of Octaves (1865): Arranged elements by increasing atomic weight, noting properties repeated every 8 elements.
• Mendeleev’s Periodic Table (1869): Created the first widely accepted table arranging elements by atomic weight and predicted undiscovered elements.
• Moseley’s Work (1913): Showed atomic number (not atomic weight) is the correct basis for element arrangement.
• Modern Periodic Table: Elements arranged by increasing atomic number and electron configurations.

2. Structure of the Periodic Table

The modern periodic table is organized into rows called periods and columns called groups.

Periods

- There are 7 periods corresponding to the principal quantum number \(n = 1\) to 7.
- Properties change gradually across a period.
- Each period starts with an alkali metal and ends with a noble gas (except the first).

Groups

- There are 18 groups in total.
- Elements in the same group have similar valence electron configurations and exhibit similar chemical properties.
- Important groups for JEE:

• Group 1: Alkali metals (except Hydrogen)
• Group 2: Alkaline earth metals
• Group 17: Halogens
• Group 18: Noble gases

Blocks

Elements are also classified into s, p, d, and f blocks based on their valence electron orbital:

• s-block: Groups 1 and 2 + Helium
• p-block: Groups 13 to 18
• d-block: Transition metals (Groups 3 to 12)
• f-block: Lanthanides and Actinides (Inner transition metals)

3. Electronic Configuration and Periodicity

The position of an element in the periodic table is closely related to its electronic configuration.

- Elements in the same group have the same number of valence electrons.
- Elements in the same period have the same number of shells.
- Periodicity of properties arises due to the repeating pattern of valence electron configurations.

General electronic configuration

• s-block: \(ns^1 - ns^2\)
• p-block: \(ns^2 np^{1-6}\)
• d-block: \((n-1)d^{1-10} ns^{0-2}\)
• f-block: \(4f^{1-14}\) or \(5f^{1-14}\)

4. Periodic Trends

Certain properties of elements show regular trends as you move across periods or down groups:

4.1 Atomic Radius

Atomic radius is the distance from nucleus to the outermost electron.

• Across a period: Decreases from left to right due to increasing nuclear charge pulling electrons closer.
• Down a group: Increases due to addition of electron shells.

4.2 Ionization Energy (IE)

Ionization energy is the energy required to remove an electron from a gaseous atom.

• Across a period: Increases due to higher nuclear charge and smaller radius.
• Down a group: Decreases due to larger atomic size and shielding.
• First IE corresponds to removing the first electron, second IE is removing the second electron and so on.

4.3 Electron Affinity (EA)

Electron affinity measures the energy change when an atom gains an electron.

• Generally becomes more negative (more energy released) across a period.
• Less variation down a group.

4.4 Electronegativity

Electronegativity is the ability of an atom to attract shared electrons in a covalent bond.

• Increases across a period and decreases down a group.
• Fluorine has the highest electronegativity.

4.5 Metallic and Non-Metallic Character

• Metallic character decreases across a period and increases down a group.
• Non-metallic character shows the opposite trend.

5. Group-wise Properties

5.1 Group 1: Alkali Metals

• Have one valence electron (\(ns^1\)).
• Very reactive metals, react vigorously with water producing hydroxides and hydrogen gas.
• Have low ionization energy and metallic character.
• Form ionic compounds, usually +1 oxidation state.
• Examples: Li, Na, K, Rb, Cs.

5.2 Group 2: Alkaline Earth Metals

• Two valence electrons (\(ns^2\)).
• Less reactive than alkali metals but still reactive.
• Form oxides and hydroxides.
• Examples: Be, Mg, Ca, Sr, Ba.

5.3 Group 17: Halogens

• Have seven valence electrons (\(ns^2 np^5\)).
• Highly reactive non-metals; exist as diatomic molecules (e.g., \(F_2\), \(Cl_2\)).
• Show oxidizing behavior; form salts with metals.
• Oxidation state: -1 in compounds.

5.4 Group 18: Noble Gases

• Full valence shell (\(ns^2 np^6\)), except helium (1s2).
• Very stable and inert under normal conditions.
• Used in lighting, welding, and as inert environments.
• Examples: He, Ne, Ar, Kr, Xe, Rn.

6. Transition Elements (d-Block)

Transition metals have partially filled d orbitals. They show multiple oxidation states and form colored compounds.

• Good conductors of heat and electricity.
• Form complex compounds with ligands.
• Exhibit magnetic properties.
• Examples: Fe, Cu, Ni, Zn.

7. Lanthanides and Actinides (f-Block)

These are inner transition metals, filling 4f and 5f orbitals respectively.

• Lanthanides: Known as rare earth elements, show similar chemical properties.
• Actinides: Mostly radioactive, include elements like Uranium and Plutonium.

8. Important Periodic Table Concepts for JEE

8.1 Effective Nuclear Charge (\(Z_{\text{eff}}\))

The net positive charge experienced by valence electrons after accounting for shielding by inner electrons.

This explains trends in atomic radius, ionization energy, and electronegativity.

8.2 Shielding Effect

Inner shell electrons shield valence electrons from the full nuclear charge, reducing effective attraction.

8.3 Diagonal Relationship

Elements diagonally adjacent in the periodic table (e.g., Li and Mg) exhibit similar properties due to comparable charge density and size.

8.4 Anomalous Electron Configurations

Some elements show exceptions to expected configurations for extra stability (half-filled or fully filled d or f subshells). Examples: Cr = [Ar] 3d5 4s1, Cu = [Ar] 3d10 4s1.

8.5 Atomic and Ionic Radii Trends

Cations are smaller than parent atoms; anions are larger due to electron-electron repulsions.

8.6 Oxidation States

Elements show typical oxidation states depending on group and electronic configuration. Transition metals have multiple oxidation states.

9. Periodic Table Practice Questions

1. Arrange the following elements in order of increasing atomic radius: Na, Mg, Al, Si.
2. Which element has the highest ionization energy among F, O, N, and C?
3. Explain why fluorine has a higher electronegativity than chlorine.
4. Predict the formula of the compound formed between Group 2 and Group 17 elements.
5. Why do transition metals exhibit variable oxidation states?

10. Tips for JEE Chemistry Preparation on Periodic Table

• Memorize key group properties and common oxidation states.
• Understand periodic trends conceptually rather than rote learning.
• Practice questions related to electronic configuration and anomaly exceptions.
• Focus on f-block elements' chemistry as it is frequently asked.
• Revise important periodic table laws and historical facts for theory questions.

Mastering the periodic table concepts will give you a significant advantage in the JEE Chemistry exam. Focus on understanding the logic behind periodicity and group properties and practice application-based problems regularly.