NEET Inorganic Chemistry: Periodic Table Trends & Shortcuts

The periodic table is not just a chart — it’s the language of Inorganic Chemistry. For NEET aspirants, mastering the periodic trends is one of the smartest ways to improve speed and accuracy. Nearly every chapter in Inorganic Chemistry — including chemical bonding, s-, p-, and d-block elements — relies heavily on periodic properties. Understanding these patterns allows students to decode element behavior, determine reactivity, and predict bond types without brute-force memorization.

Many students approach Inorganic Chemistry with fear because it appears to be memory-intensive. But the truth is, once you crack the logic behind periodicity, the rest of Inorganic Chemistry becomes manageable. In this comprehensive guide from Deeksha Vedantu, we’ll dive deep into the essential periodic trends that appear in NEET exams, explore smart shortcuts, and give you memory hacks that will save you valuable exam time.

Chapter Overview: Periodic Table & Periodic Properties

The periodic table is your foundational tool. It helps you understand how and why elements behave a certain way. NEET often includes both direct and indirect questions from the chapter “Classification of Elements and Periodicity in Properties” from Class 11 NCERT.

Key properties tested in NEET include:

• Atomic and Ionic Radius
• Ionization Enthalpy
• Electron Gain Enthalpy
• Electronegativity
• Metallic and Non-Metallic Character
• Oxidation States

These properties don’t exist in isolation. They influence:

• Bond strength predictions in Chemical Bonding
• Reactivity orders in s- and p-block chapters
• Redox potential in Coordination Compounds
• Acid-base nature of oxides
• Trends in basic and acidic strength of hydrides and oxoacids

A solid grip on these properties helps in not just scoring well in periodicity-related questions but also boosts your performance in related chapters.

Key Periodic Trends Explained

Atomic Radius & Ionic Radius

Atomic radius is the distance from the nucleus to the outermost electron. Ionic radius is the size of an ion after loss or gain of electrons.

• Atomic radius decreases across a period due to increased nuclear charge pulling electrons closer.
• It increases down a group as new shells are added.
• Cations are smaller than their parent atoms due to increased nuclear pull.
• Anions are larger due to increased repulsion between added electrons.

Shortcut: ARAD — Atomic Radius Across Decreases

Watch Out For: Transition metals have irregular trends due to ineffective shielding by d-electrons. Lanthanide contraction also causes slight anomalies.

Ionization Enthalpy

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

• Increases across a period due to stronger nuclear attraction.
• Decreases down a group as atomic size increases and shielding weakens attraction.

Important Notes:

• Second and third ionization enthalpies are always higher.
• Elements with stable configurations (e.g., full or half-filled subshells) have unusually high values.

Exceptions:

• Boron < Beryllium
• Oxygen < Nitrogen

This is due to electron configuration stability and inter-electronic repulsion.

Electron Gain Enthalpy

Electron gain enthalpy is the energy released when an atom accepts an electron.

• Becomes more negative across a period (greater attraction for electrons).
• Becomes less negative down a group (weaker attraction).

Interesting Exception: Fluorine has lower (less negative) electron gain enthalpy than chlorine due to high electron-electron repulsion in its small 2p orbital.

Electronegativity

Electronegativity is the tendency of an atom to attract a bonding electron pair.

• Increases across a period due to increasing nuclear charge.
• Decreases down a group due to increased atomic size and shielding.

Fluorine tops the scale (Pauling scale value of 4.0).

Application: Determines bond polarity, acidity/basicity, and bond types (ionic/covalent).

Metallic & Non-Metallic Character

• Metals are electropositive; they lose electrons easily.
• Non-metals are electronegative; they gain electrons.

Trend:

• Metallic character increases down a group, decreases across a period.
• Non-metallic characters do the opposite.

Important Application:

• Metals form basic oxides.
• Non-metals form acidic oxides.
• Amphoteric oxides show intermediate behavior (e.g., Al₂O₃).

Oxidation States

• Main group elements show oxidation states based on their group number.
• Transition elements exhibit variable oxidation states due to d-electrons.

Shortcut for Main Group:

• Max oxidation state = Group number
• Min oxidation state = Group number – 8 (except alkali/alkaline earth metals)

Trick: Group 17 elements exhibit +1, +3, +5, +7 oxidation states (except fluorine, which shows only -1).

Tables for Quick Revision

Periodic Property Comparison