Classification of Elements and Periodicity in Properties – Class 11 Chemistry Chapter 3 | NCERT Notes for JEE & NEET

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1. Introduction – Classification of Elements and Periodicity in Properties

🧠 Topics Covered:

• Introduction
• Need for Classification
• Historical Background

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🔷 1. Introduction

The discovery of more and more elements over time led to confusion in studying and understanding their properties. With over 100 known elements today, it became essential to organize them in a way that allowed scientists to predict, study, and understand their behaviors efficiently.

✅ Goal of classification:
To systematically organize the elements so that:

• Similar elements are grouped together.
• Chemical behavior can be predicted.
• Relationships among elements become clearer.

Before classification, studying elements was like memorizing random facts. The idea was to bring order to chaos by finding patterns in elemental properties.

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🔷 2. Need for Classification

Why classify elements?

Reason	Explanation
🌱 Growing number of elements	By the 18th century, nearly 60 elements were known; now over 118. A system was required to manage them.
🧪 Understanding chemical properties	Grouping elements with similar properties made it easier to predict their behavior.
🔍 Avoiding repetition	Without a pattern, each element had to be studied individually. Classification revealed periodic repetition in properties.
📚 Simplifying study	Classification helped in structuring chemistry textbooks and made learning more organized.
🔄 Revealing trends	Classification showed how properties like atomic size, ionization energy, etc., change periodically.

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Analogy:

Think of classification like arranging books in a library—grouping similar genres together so they’re easier to find and understand.

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🔷 3. Historical Background

Let’s take a look at how scientists attempted classification over time:

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🧪 Early Attempts

🔹 a) Dobereiner’s Triads (1817)

• Grouped elements in triads (groups of 3).
• Middle element had atomic mass = average of the other two.
• Example: Li (7), Na (23), K (39) ⇒ (7+39) / 2 = 23
• Limitation: Only worked for a few elements.

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🔹 b) Newlands’ Law of Octaves (1865)

• Arranged elements by increasing atomic mass.
• Every 8th element had similar properties (like musical notes).
• Worked well up to Calcium (Ca).
• Limitation: Failed for heavier elements, forced dissimilar elements together.

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⚗️ Major Breakthrough: Mendeleev’s Periodic Table (1869)

• Based on increasing atomic mass.
• Grouped elements with similar properties into columns (groups).
• Left gaps for undiscovered elements and predicted their properties (e.g., Gallium, Germanium).
• Corrected atomic masses of some elements.
• Limitation: Could not explain the position of isotopes and anomalies like Co-Ni.

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🔬 Modern Periodic Table (Henry Moseley, 1913)

• Introduced the concept of atomic number (Z) as the organizing principle.
• Modern Periodic Law:
  “The properties of elements are periodic functions of their atomic numbers.”
• Solved earlier problems:
  • No anomaly with Co/Ni or Te/I.
  • Isotopes placed correctly.
• Basis for today’s Modern Periodic Table.

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🧠 Summary Chart:

Scientist	Year	Basis	Successes	Failures
Dobereiner	1817	Atomic mass (Triads)	Atomic mass relationships	Limited applicability
Newlands	1865	Atomic mass (Octaves)	Repeating properties	Failed after Ca
Mendeleev	1869	Atomic mass (Table)	Predicted new elements	Isotope anomaly
Moseley	1913	Atomic number	Corrected earlier errors	—

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✅ Conclusion

The need for classification arises from the increasing number of elements and the desire to organize their properties meaningfully. Over time, scientists refined their understanding — from Dobereiner’s triads to Moseley’s atomic number, leading to the modern periodic table, which forms the backbone of modern chemistry.

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🧠 For JEE/NEET:

You should be able to:

• Explain why classification is needed.
• Compare historical attempts.
• Define the modern periodic law.
• Identify advantages and limitations of each classification system.

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2. 📘 Dobereiner’s Triads (1817)

🔷 1. Concept

Johann Wolfgang Dobereiner, a German chemist, was the first to classify elements into groups based on similarities in chemical properties.

He observed that certain groups of three elements (called triads) showed a unique pattern:

🧠 Dobereiner’s Law of Triads:
When elements are arranged in increasing order of their atomic masses, the atomic mass of the middle element is approximately the arithmetic mean of the atomic masses of the first and third elements.

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🔷 2. Examples of Dobereiner’s Triads

Here are some famous triads with their atomic masses and properties:

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✅ Triad 1: Alkali Metals

Element	Symbol	Atomic Mass
Lithium	Li	7
Sodium	Na	23
Potassium	K	39

Check: Mean of Li and K = (7+39) / 2 = 23 = Na (middle element)

🔹 Properties: All are soft, highly reactive metals, react with water to form hydroxides.

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✅ Triad 2: Alkaline Earth Metals

Element	Symbol	Atomic Mass
Calcium	Ca	40
Strontium	Sr	88
Barium	Ba	137

Check: (40+137) / 2 = 88.5 ≈ 88=Sr

🔹 Properties: All form +2 ions, react with acids to release H₂ gas.

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✅ Triad 3: Halogens

Element	Symbol	Atomic Mass
Chlorine	Cl	35.5
Bromine	Br	80
Iodine	I	127

Check: (35.5+127) / 2 = 81.25 ≈ 80 = Br

🔹 Properties: Non-metals, form diatomic molecules (Cl₂, Br₂, I₂), show similar reactivity with metals.

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🔷 3. Graphical Representation

📈 If you plot the atomic mass of the elements, the middle one lies between the two in a linear pattern.

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🔷 4. Successes of Dobereiner’s Triads

✅ Strengths
Introduced the idea of grouping elements based on chemical properties.
Helped in showing periodicity — a repeating pattern in properties.
Inspired future periodic classifications like Newlands’ Octaves and Mendeleev’s Table.

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🔴 5. Limitations of Dobereiner’s Triads

❌ Limitations
Only a few triads could be formed out of the known elements.
Did not include all known elements at the time.
Atomic mass of some elements did not fit the arithmetic mean rule.
Could not explain properties of transition elements or lanthanides/actinides.

For example, elements like Nitrogen, Phosphorus, Arsenic do not form a triad.

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🔷 6. Importance for JEE/NEET

• Often asked as a theoretical short-answer.
• May appear in assertion-reason questions.
• Understanding triads is key to grasping how periodic trends were discovered.

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🧠 Final Summary

Feature	Dobereiner’s Triads
Proposed by	Johann Dobereiner (1817)
Basis	Atomic mass and chemical similarity
Pattern	Middle element’s atomic mass ≈ average of the other two
Significance	Early step toward periodic classification
Limitations	Couldn’t explain all elements or predict new ones

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3. 📘 Newlands’ Law of Octaves (1865)

🔷 1. Introduction

As more elements were being discovered in the 19th century, chemists tried to find a logical way to organize them.
In 1865, John Newlands, an English chemist, proposed a new form of classification based on the idea of musical notes.

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🔷 2. Newlands’ Law of Octaves

🧠 Statement:
“When elements are arranged in increasing order of their atomic masses, every eighth element shows similar chemical properties to the first one, just like the eighth note in a musical octave.”

🎵 Musical Analogy:
In music, the eighth note resembles the first note in terms of tone:
Sa Re Ga Ma Pa Dha Ni Sa → 8th = same as 1st (repetition begins)

Newlands applied this idea to the periodic classification of elements.

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🔷 3. Arrangement of Elements

Newlands arranged the known elements at the time (only 56 elements) in a tabular form where each row had 7 elements, and the 8th element repeated the properties of the 1st.

🧾 Table Showing Newlands’ Octaves

Sa	Re	Ga	Ma	Pa	Dha	Ni
H	Li	Be	B	C	N	O
F	Na	Mg	Al	Si	P	S
Cl	K	Ca	—	—	—	—

Examples:

• H and F
• Li and Na
• Be and Mg
• B and Al
• C and Si
• N and P
• O and S

📌 Observation: Every eighth element had similar properties to the first one in the row.

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🔷 4. Successes / Merits

✅ Advantages of Newlands’ Law
First to show periodicity in properties.
Based on a logical sequence — atomic mass.
Correctly grouped some elements like Li–Na, Be–Mg, F–Cl.
Inspired later scientists like Mendeleev.

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🔴 5. Limitations of Newlands’ Law of Octaves

❌ Limitation	🔍 Explanation
Only worked till calcium (Ca)	Law failed for elements after Ca; chemical properties did not repeat reliably.
Dissimilar elements in same column	Example: Co & Ni grouped with halogens like F and Cl (which was chemically incorrect).
No room for undiscovered elements	Did not leave gaps; forced elements into the pattern.
Works only for lighter elements	Heavier elements didn’t follow the octave rule.
No distinction for noble gases	Noble gases weren’t discovered yet; inclusion would’ve disrupted the pattern.

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🔷 6. Comparison with Dobereiner’s Triads

Feature	Dobereiner’s Triads	Newlands’ Octaves
Year	1817	1865
Pattern	Atomic mass avg. in triads	Every 8th element similar
Number of elements covered	Very few	Up to Calcium
Major limitation	Very limited triads	Failed for heavier elements

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🔷 7. JEE/NEET Relevance

🧪 Mostly appears in theory-based or assertion-reason type questions.
You should be able to:

• State the law of octaves.
• Identify the correct 8th element pairs.
• Point out limitations.
• Compare with other classification attempts.

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🧠 Final Summary

Aspect	Newlands’ Law of Octaves
Proposed by	John Newlands (1865)
Basis	Increasing atomic mass
Key idea	Every 8th element shows similar properties
Number of elements explained	Up to Calcium
Significance	First idea of periodicity
Limitation	Failed beyond Ca, no gaps, dissimilar elements in groups

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4. 📘 Mendeleev’s Periodic Table (1869)

Dmitri Ivanovich Mendeleev, a Russian chemist, made the most successful attempt to classify elements in a systematic manner before the modern periodic table.

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🔷 1. Basis of Classification

🧠 Mendeleev’s Periodic Law (1869):

“The properties of elements are a periodic function of their atomic masses.”

In other words, when elements are arranged in increasing order of atomic mass, their chemical and physical properties repeat periodically.

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🧱 Structure of the Table:

• Elements were arranged in horizontal rows called periods and vertical columns called groups.
• Elements with similar properties were placed in the same group.
• Mendeleev’s table had:
  • 8 groups (I to VIII), where group VIII had 3 elements placed together (triads).
  • 6 periods.

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⚙️ Features of Mendeleev’s Table:

Feature	Description
📊 Rows (Periods)	Horizontal rows where atomic masses increase from left to right
📈 Columns (Groups)	Vertical columns where elements show similar chemical properties
❓ Gaps left intentionally	For undiscovered elements like gallium, scandium, and germanium
🧮 Order based on atomic mass	Not atomic number (which was not discovered at that time)

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🔷 2. Merits of Mendeleev’s Periodic Table

Mendeleev’s work was revolutionary and had several advantages:

✅ Merit	🔍 Explanation
🧪 Systematic study	For the first time, elements were arranged in a logical pattern that could be studied systematically.
🧭 Prediction of new elements	Left gaps for undiscovered elements and predicted their properties with surprising accuracy. For example: Eka-aluminium (→ Gallium), Eka-silicon (→ Germanium).
🧬 Correction of atomic masses	Rectified incorrect atomic masses of elements (e.g., Beryllium, Indium).
⚛️ Grouping of similar elements	Elements with similar properties like halogens (F, Cl, Br, I) and alkali metals (Li, Na, K, Rb) were placed together.
🔁 Periodicity observed	Recognized and formalized the periodic repetition of chemical properties.

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🔷 3. Examples of Mendeleev’s Predictions

Predicted Element	Actual Element	Properties predicted
Eka-silicon	Germanium (Ge)	Density, atomic mass, oxide formula
Eka-aluminium	Gallium (Ga)	Melting point, reactivity
Eka-boron	Scandium (Sc)	Oxidation state, density

This accurate prediction of properties gave huge credibility to Mendeleev’s work.

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🔴 4. Limitations of Mendeleev’s Periodic Table

Despite its brilliance, Mendeleev’s table had several shortcomings:

❌ Limitation	🔍 Explanation
⚠️ Position of isotopes not explained	Isotopes have same atomic number but different atomic masses — the table didn’t accommodate them.
❌ Anomalous pairs	Some elements were reversed to maintain group properties, violating atomic mass order. Example: Co (58.9) placed before Ni (58.7).
🧩 No place for noble gases	Noble gases were discovered later (after 1894), so not included in the original table.
🧪 Dissimilar elements grouped together	Some elements with different properties ended up in the same group (e.g., Cu, Ag, Au with alkali metals).
📉 No explanation for periodicity	Periodic trends were observed but not explained (because the concept of atomic number was not yet known).
💥 Lanthanides and Actinides not addressed	These elements didn’t fit well into the periodic table structure.

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🔷 5. Legacy and Importance

Despite its flaws, Mendeleev’s Periodic Table laid the foundation of modern chemistry:

• It led to the development of the Modern Periodic Table by Henry Moseley (1913), based on atomic number.
• Helped in systematizing chemical knowledge.
• Influenced the discovery and study of new elements.

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🔷 6. Comparison: Mendeleev vs Modern Table

Feature	Mendeleev’s Table	Modern Table
Basis	Atomic Mass	Atomic Number
Periods	6	7
Groups	8	18
Gaps	Yes (for new elements)	No gaps
Isotopes	Not explained	Explained
Noble gases	Absent initially	Present in Group 18

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🧠 Final Summary

📌 Aspect	🔍 Mendeleev’s Table
Proposed by	Dmitri Mendeleev (1869)
Classification Basis	Increasing atomic mass
Structure	8 groups, 6 periods
Key Law	Periodic Law (properties repeat with atomic mass)
Strength	Prediction of elements, systematic grouping
Weakness	Anomalies, isotopes, noble gases excluded
Legacy	Basis for modern periodic table

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🎯 JEE/NEET Tip:

You must be able to:

• State Mendeleev’s Periodic Law
• Give examples of predicted elements
• List its merits and limitations
• Compare it with the modern periodic table

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5. Modern Periodic Law & Modeden Periodic Table

🔷 1. MODERN PERIODIC LAW

🧠 Statement of Modern Periodic Law:

“The physical and chemical properties of elements are a periodic function of their atomic numbers.”
— Proposed by Henry Moseley in 1913

📌 Explanation:

• Mendeleev arranged elements based on atomic mass, but some anomalies remained (e.g., Co & Ni, Ar & K).
• Moseley discovered that atomic number (Z), not atomic mass, is the true fundamental property of an element.
• When elements are arranged in order of increasing atomic number, their properties repeat periodically.

🔁 Periodic Repetition:

• Properties like valency, electronegativity, metallic/non-metallic character, atomic/ionic size, ionization enthalpy, etc., repeat at regular intervals.

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🔷 2. STRUCTURE OF THE MODERN PERIODIC TABLE

The Modern Periodic Table is the tabular arrangement of elements based on increasing atomic number.

📊 Key Features:

Feature	Description
🔢 Basis	Atomic number (Z)
🔄 Periodicity	Due to repetition of electronic configurations
📏 Total Elements	118 known elements (as of 2025)
🧬 Division	7 periods and 18 groups
🧱 Blocks	s-block, p-block, d-block, f-block
🌈 Zig-zag line	Separates metals from nonmetals

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🔷 3. PERIODS IN MODERN PERIODIC TABLE

📐 Horizontal Rows = Periods

There are 7 periods in the modern periodic table:

Period	No. of Elements	Starts With	Ends With
1st	2	H (Z=1)	He (Z=2)
2nd	8	Li (Z=3)	Ne (Z=10)
3rd	8	Na (Z=11)	Ar (Z=18)
4th	18	K (Z=19)	Kr (Z=36)
5th	18	Rb (Z=37)	Xe (Z=54)
6th	32	Cs (Z=55)	Rn (Z=86)
7th	32 (incomplete)	Fr (Z=87)	Og (Z=118)

📌 Note:

• 6th & 7th periods include lanthanides and actinides.
• Number of elements per period depends on number of available orbitals.

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🔷 4. GROUPS IN MODERN PERIODIC TABLE

📏 Vertical Columns = Groups

There are 18 groups in total:

Group No.	Type of Elements
1	Alkali Metals
2	Alkaline Earth Metals
13–18	p-block elements
3–12	Transition elements (d-block)
Bottom two rows	Inner Transition elements (f-block) – Lanthanides & Actinides

💡 Same group elements have:

• Same valence shell configuration
• Similar chemical properties

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🔷 5. BLOCKS OF THE PERIODIC TABLE

The periodic table is divided into 4 blocks based on the type of orbital that receives the last electron.

🧱 BLOCK-WISE CLASSIFICATION:

Block	Elements	Last e⁻ enters	Groups Covered
s-block	Alkali & Alkaline Earth Metals + He	s-orbital	Groups 1 & 2
p-block	Non-metals, Halogens, Noble Gases, Metalloids	p-orbital	Groups 13–18
d-block	Transition Metals	d-orbital	Groups 3–12
f-block	Lanthanides & Actinides	f-orbital	Not in main table (placed separately)

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🔷 6. TYPES OF ELEMENTS IN MODERN PERIODIC TABLE

Based on position & properties, elements are categorized as:

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🟥 A. Metals

Feature	Description
Location	Left and center of the table
Number	≈ 78% of all elements
Properties	Malleable, ductile, conductors, electropositive, lose e⁻ easily
Blocks	s, d, f, and some p-block

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🟨 B. Non-Metals

Feature	Description
Location	Upper right corner
Properties	Brittle, poor conductors, electronegative, gain e⁻
Examples	H, C, N, O, F, P, S, Cl, Br, I

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🟩 C. Metalloids

Feature	Description
Location	Along the “stair-step” zig-zag line
Properties	Intermediate between metals and non-metals
Examples	B, Si, Ge, As, Sb, Te

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🟦 D. Noble Gases

Feature	Description
Location	Group 18
Properties	Stable, inert, monoatomic gases
Examples	He, Ne, Ar, Kr, Xe, Rn

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🟪 E. Transition Elements

Feature	Description
Location	d-block (Groups 3–12)
Properties	Variable oxidation states, colored compounds, form complexes

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🟫 F. Inner Transition Elements

Feature	Description
Location	f-block (bottom of table)
Series	Lanthanides (Z=58–71), Actinides (Z=90–103)
Properties	Mostly radioactive (especially actinides), large atomic size

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🔷 7. PERIODICITY AND TRENDS (Brief Mention)

The modern table allows for understanding trends like:

• Atomic radius
• Ionization enthalpy
• Electron affinity
• Electronegativity
• Metallic/Non-metallic character

These trends repeat periodically across periods and down groups due to repetition of electronic configurations.

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🧠 FINAL SUMMARY

🔑 Concept	📘 Modern Periodic Table
Proposed by	Henry Moseley (1913)
Based on	Atomic number
No. of periods	7
No. of groups	18
Classification	s, p, d, f blocks
Types of elements	Metals, Nonmetals, Metalloids, Noble Gases, Transition & Inner Transition Elements
Periodicity due to	Repetition of similar valence shell configurations

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🎯 JEE/NEET Tip:

• Questions often test your knowledge of periodicity, group/period positions, and electronic configuration.
• Be comfortable identifying block and group from electronic configuration.
• Understand properties across periods and down groups (e.g., metallic character increases down a group).

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6. Nomenclature of Elements with Atomic Number > 100

🔷 Introduction

When elements with atomic number greater than 100 were discovered, they were not named immediately. To avoid confusion and ensure uniformity, IUPAC (International Union of Pure and Applied Chemistry) devised a systematic nomenclature.

✅ Why is this needed?

• Many heavy elements are synthetic and short-lived.
• Their discovery often led to naming disputes.
• Temporary naming ensures neutrality and clarity until a permanent name is decided.

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🔷 IUPAC Rules for Naming Elements (Z > 100)

🧠 IUPAC Nomenclature Rule (1978):

For elements with Z > 100, names are based on Latin/Greek roots of their atomic number digits, followed by -ium as the suffix.

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📌 Rule Format:

Name = Root1 + Root2 + Root3 + … + “ium”

Where:

• Each digit in the atomic number is replaced with a specific root word.
• The roots are combined in the order of digits.
• Ends with “-ium” regardless of the element being a metal or not.
• Some letter combinations are modified to avoid awkward repetition (explained below).

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🔷 Table: Roots for Each Digit

Digit	Root
0	nil
1	un
2	bi
3	tri
4	quad
5	pent
6	hex
7	sept
8	oct
9	enn

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🧩 Special Guidelines:

Situation	Rule
If the last root ends in “i” and is followed by “ium”	Drop the extra “i” (e.g., biium → bium)
If “enn” is followed by “nil”	Keep both (e.g., 9 & 0 → ennnil, not “enil”)
Double letters allowed	Except “ii” (remove one “i”)
Symbol	Formed by taking first letter of each root, capitalizing the first letter (e.g., Unq for Unnilquadium)

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🔷 Examples with Explanation

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✅ 1. Element 104

• Digits: 1 – 0 – 4
• Roots: un – nil – quad
• Name: Unnilquadium
• Symbol: Unq

🧾 Now permanently named as Rutherfordium (Rf).

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✅ 2. Element 105

• Digits: 1 – 0 – 5
• Roots: un – nil – pent
• Name: Unnilpentium
• Symbol: Unp

🧾 Now known as Dubnium (Db).

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✅ 3. Element 106

• Digits: 1 – 0 – 6
• Roots: un – nil – hex
• Name: Unnilhexium
• Symbol: Unh

🧾 Now named Seaborgium (Sg).

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✅ 4. Element 118

• Digits: 1 – 1 – 8
• Roots: un – un – oct
• Name: Ununoctium
• Symbol: Uuo

🧾 Now named Oganesson (Og).

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✅ 5. Element 109

• Digits: 1 – 0 – 9
• Roots: un – nil – enn
• Name: Unnilennium
• Symbol: Une

🧾 Now named Meitnerium (Mt).

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✅ 6. Element 113

• Digits: 1 – 1 – 3
• Roots: un – un – tri
• Name: Ununtrium
• Symbol: Uut

🧾 Now named Nihonium (Nh).

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🔷 Final Names vs. Systematic Names

Atomic Number	Systematic Name	Permanent IUPAC Name
104	Unnilquadium (Unq)	Rutherfordium (Rf)
105	Unnilpentium (Unp)	Dubnium (Db)
106	Unnilhexium (Unh)	Seaborgium (Sg)
107	Unnilseptium (Uns)	Bohrium (Bh)
108	Unniloctium (Uno)	Hassium (Hs)
109	Unnilennium (Une)	Meitnerium (Mt)
110	Ununnilium (Uun)	Darmstadtium (Ds)
111	Unununium (Uuu)	Roentgenium (Rg)
112	Ununbium (Uub)	Copernicium (Cn)
113	Ununtrium (Uut)	Nihonium (Nh)
114	Ununquadium (Uuq)	Flerovium (Fl)
115	Ununpentium (Uup)	Moscovium (Mc)
116	Ununhexium (Uuh)	Livermorium (Lv)
117	Ununseptium (Uus)	Tennessine (Ts)
118	Ununoctium (Uuo)	Oganesson (Og)

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🔷 JEE / NEET Relevance

• Important for theoretical MCQs, especially in JEE.
• Often asked in match-the-following questions.
• You may be asked to decode the temporary name or symbol.

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🧠 Final Summary

Feature	Detail
Applies to	Elements with Z > 100
Based on	Atomic number digits
Uses	Latin/Greek roots for each digit
Ends with	“-ium”
Symbol	First letters of the roots
Purpose	Temporary, systematic names until official names are assigned

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7. Electronic Configuration & the Periodic Table

🔷 1. Introduction

The electronic configuration of an element refers to the distribution of electrons in the various atomic orbitals.

It forms the core basis of the Modern Periodic Table, which is arranged according to increasing atomic number (Z) and electronic configuration.

✅ Why is it important?

• Explains the position of elements in the periodic table
• Determines group, period, and block
• Predicts valency, chemical reactivity, and periodic trends

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🔷 2. Relation with Group and Period

📌 A. Group Number and Electronic Configuration

The group number is linked to the number of valence electrons (electrons in the outermost shell).

Element Type	Group Number Rule
Main group (s & p block)	Number of valence electrons (or valence electrons + 10 for groups 13–18)
Transition elements (d block)	(n-1)d and ns electrons contribute
Inner transition (f block)	Electrons enter (n-2)f subshell; group number is not strictly defined

👉 Examples:

• Na (Z = 11): 1s² 2s² 2p⁶ 3s¹ → Valence electron = 1 → Group 1
• Mg (Z = 12): 1s² 2s² 2p⁶ 3s² → Group 2
• Cl (Z = 17): 3s² 3p⁵ → Group 17
• Fe (Z = 26): [Ar] 3d⁶ 4s² → Group 8

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📌 B. Period Number and Electronic Configuration

The period number corresponds to the highest principal quantum number (n) occupied by electrons.

Element	Configuration	Period
Na (11)	3s¹	3
Cl (17)	3s² 3p⁵	3
K (19)	4s¹	4
Fe (26)	3d⁶ 4s²	4

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🔷 3. Valence Electrons

Valence electrons are the outermost electrons involved in bonding. They determine:

• Group number
• Valency
• Reactivity

Block	Valence Electrons
s-block	ns¹–²
p-block	ns² np¹–⁶
d-block	(n-1)d¹–¹⁰ ns⁰–²
f-block	(n-2)f¹–¹⁴ (variable)

Example:

• Carbon (Z = 6) → 1s² 2s² 2p² → Valence electrons = 4 → Valency = 4
• Oxygen (Z = 8) → 1s² 2s² 2p⁴ → Valence electrons = 6 → Valency = 2

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🔷 4. Block Classification (s, p, d, f)

The periodic table is divided into 4 blocks based on the type of orbital in which the last electron enters.

📘 A. s-block Elements

• Groups 1 and 2
• General configuration: ns¹–²
• Soft, highly reactive metals (alkali & alkaline earth metals)

Examples:

• Li: 1s² 2s¹ → s-block
• Ca: [Ar] 4s² → s-block

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📘 B. p-block Elements

• Groups 13 to 18
• General configuration: ns² np¹–⁶
• Includes metals, metalloids, nonmetals, and noble gases

Examples:

• N: 1s² 2s² 2p³ → p-block
• Cl: [Ne] 3s² 3p⁵ → p-block

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📘 C. d-block Elements (Transition Metals)

• Groups 3 to 12
• General configuration: (n-1)d¹–¹⁰ ns⁰–²
• Show variable valency, form colored compounds, and are often good catalysts.

Examples:

• Fe: [Ar] 3d⁶ 4s² → d-block
• Zn: [Ar] 3d¹⁰ 4s² → d-block

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📘 D. f-block Elements (Inner Transition Elements)

• Lanthanides: Z = 58 to 71
• Actinides: Z = 90 to 103
• General configuration: (n-2)f¹–¹⁴ (n-1)d⁰–¹ ns²

Examples:

• Ce: [Xe] 4f¹ 5d¹ 6s² → f-block
• U: [Rn] 5f³ 6d¹ 7s² → f-block

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🔷 5. Trends Based on Electronic Configuration

Trend	Explanation
Valency	Based on number of valence electrons
Atomic Radius	Increases down group, decreases across period
Ionization Energy	Decreases down group, increases across period
Electronegativity	Decreases down group, increases across period
Metallic character	Increases down group, decreases across period

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🔷 6. Summary Table

Concept	Explanation	Example
Group Number	Related to number of valence electrons	Mg (3s²) → Group 2
Period Number	Based on highest n value	Cl (3s²3p⁵) → Period 3
Valence Electrons	Outer shell electrons	O (2s² 2p⁴) → 6 valence electrons
Block	Based on orbital receiving last electron	Fe (3d⁶ 4s²) → d-block
Type of Element	s = active metals, p = diverse, d = transition, f = inner transition

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🔷 Visual Classification (Quick Reference)

Block	Groups	Last Electron Enters	Example	Characteristics
s	1-2	s-orbital	Na (3s¹)	Soft, reactive metals
p	13-18	p-orbital	O (2p⁴)	Nonmetals, metalloids
d	3-12	d-orbital	Fe (3d⁶)	Transition metals
f	–	f-orbital	U (5f³)	Rare earth metals

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🎯 JEE/NEET Tips

• Know configurations up to Z = 30 (and d-block till Z = 56)
• Use the (n + l) rule to write configurations
• Expect MCQs on identifying period/group from configuration
• Understand why transition and inner transition elements behave differently

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🧠 Final Summary

Concept	Key Points
Electronic Configuration	Determines position in periodic table
Group	Based on valence electrons
Period	Based on principal quantum number (n)
Block	Based on type of orbital receiving last electron
Periodic Trends	Arise due to similar configurations across periods/groups

8. Periodic Trends in Properties

🔷 Introduction to Periodic Trends

When elements are arranged in the Modern Periodic Table (in increasing order of atomic number), they show regular patterns or trends in certain physical and chemical properties. These trends are called:

✅ Periodic Properties – Properties that repeat at regular intervals (i.e., show periodicity) when elements are arranged by increasing atomic number.

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🧠 Why Do Trends Occur?

The periodic trends occur mainly due to:

• Effective Nuclear Charge (Z_eff)
• Atomic Radius
• Number of Shells
• Electron Configuration
• Shielding or Screening Effect

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📊 List of Periodic Properties Showing Trends

S. No.	Periodic Property
1.	Atomic Radius / Atomic Size
2.	Ionic Radius
3.	Ionization Enthalpy (Ionization Energy)
4.	Electron Gain Enthalpy (Electron Affinity)
5.	Electronegativity
6.	Metallic and Non-metallic Character
7.	Valency
8.	Chemical Reactivity

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🔷 1. Atomic Radius / Atomic Size

Definition: Distance from the nucleus to the outermost shell (valence shell) of an atom.

🔁 Trend Across a Period (→ left to right):

• Decreases
• Reason: Increase in effective nuclear charge pulls electrons closer.

🔃 Trend Down a Group (↓ top to bottom):

• Increases
• Reason: Addition of new shells increases atomic size.

🧪 Example:

• Period: Li > Be > B > C > N > O > F
• Group: F < Cl < Br < I

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🔷 2. Ionic Radius

Definition: Radius of an ion (cation or anion).

✅ Cations (+ve ions):

• Smaller than parent atom (loss of shell).
• Example: Na⁺ < Na

❌ Anions (−ve ions):

• Larger than parent atom (electron repulsion).
• Example: Cl⁻ > Cl

🧠 Trend:

• Similar to atomic size → decreases across a period, increases down a group.

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🔷 3. Ionization Enthalpy (IE)

Definition: The minimum energy required to remove the outermost electron from a gaseous atom.

M _(g)  → M⁺_(g) + e⁻

🔁 Across a Period:

• Increases → Due to higher Z_eff and smaller size.

🔃 Down a Group:

• Decreases → Electron is farther from the nucleus and easier to remove.

🧪 Exceptions:

• Be > B, N > O due to stable electronic configurations.

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🔷 4. Electron Gain Enthalpy (EGE)

Definition: Enthalpy change when an electron is added to a neutral gaseous atom.

X_(g) +e⁻ → X⁻_(g)

🔁 Across a Period:

• Becomes more negative → Atoms more eager to gain electrons.

🔃 Down a Group:

• Becomes less negative → Atoms less willing to gain electrons.

⚠️ Note:

• Group 17 (halogens) have highest electron gain enthalpy.
• Noble gases have positive electron gain enthalpy (do not accept electrons).

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🔷 5. Electronegativity

Definition: The tendency of an atom to attract a shared pair of electrons in a chemical bond.

• Scale: Pauling Scale (F = 4.0, highest)

🔁 Across a Period:

• Increases (higher Z_eff, stronger attraction)

🔃 Down a Group:

• Decreases (more shielding, less pull)

🧪 Order (Across Period 2):

Li < Be < B < C < N < O < F

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🔷 6. Metallic and Non-Metallic Character

✅ Metallic Character:

• Tendency to lose electrons.
• Increases down a group
• Decreases across a period

❌ Non-Metallic Character:

• Tendency to gain electrons.
• Increases across a period
• Decreases down a group

🧪 Examples:

• Na is more metallic than Mg.
• Cl is more non-metallic than P.

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🔷 7. Valency

Definition: Number of electrons lost, gained or shared to complete an octet.

🔁 Across a Period:

• Increases from 1 to 4, then decreases back to 1.

🔃 Down a Group:

• Remains same (same number of valence electrons)

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🔷 8. Chemical Reactivity

• Metals: Reactivity increases down a group, decreases across a period.
• Non-metals: Reactivity decreases down a group, increases across a period.

🧪 Examples:

• Alkali metals (Group 1): Reactivity increases down the group (Li < Na < K).
• Halogens (Group 17): Reactivity decreases down the group (F > Cl > Br > I).

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📌 Summary Table of Periodic Trends

Property	Across Period (→)	Down Group (↓)
Atomic Radius	Decreases	Increases
Ionic Radius	Decreases	Increases
Ionization Enthalpy	Increases	Decreases
Electron Gain Enthalpy	More Negative	Less Negative
Electronegativity	Increases	Decreases
Metallic Character	Decreases	Increases
Non-Metallic Character	Increases	Decreases
Valency	Increases then Decreases	Constant
Reactivity (Metals)	Decreases	Increases
Reactivity (Non-metals)	Increases	Decreases

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💡 Important MCQ Tips (JEE/NEET)

• Ionization energy trend questions are common.
• Know exceptions in electron gain enthalpy (Be, N, Ne).
• Be careful with atomic/ionic size order-based problems.
• Remember periodic trends for p-block and s-block elements.

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8.1 ➤ Atomic Radius

🔹 Definition:

Atomic radius is the distance from the nucleus of an atom to the outermost shell of electrons (valence shell).

However, since the electron cloud is diffuse and probabilistic, we don’t measure atomic radius directly. Instead, we define it using different bonding contexts:

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🧪 Types of Atomic Radii

Type	Definition	Applicable To	Symbol
Covalent Radius	Half the distance between the nuclei of two covalently bonded atoms of the same element	Non-metals (e.g. H₂, Cl₂, O₂)	rcov
Metallic Radius	Half the distance between nuclei of two adjacent atoms in the metallic lattice	Metals (e.g. Na, Fe, Cu)	rmetal​
Van der Waals Radius	Half the distance between two non-bonded atoms in neighboring molecules	Noble gases or unbonded atoms (e.g. Ar, Ne)	rvdW​

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🔹 Comparison of Radii (Generally):

r_vdW​ > r_metal ​> r_cov​

This is because Van der Waals radii involve atoms that are not bonded, so they are further apart.

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📏 Typical Examples

Element	Covalent Radius (pm)	Metallic Radius (pm)	Van der Waals Radius (pm)
Hydrogen (H)	37	—	120
Chlorine (Cl)	99	—	180
Sodium (Na)	—	186	—
Argon (Ar)	—	—	188

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📈 Trends in Atomic Radius

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🔸 Across a Period (Left → Right)

Observation	Explanation
Atomic radius decreases	As you move across a period, the atomic number increases, increasing nuclear charge (Z). Electrons are added to the same shell, so they are pulled closer to the nucleus.
Effective nuclear charge (Z_eff) increases	Shielding remains nearly constant, but the pull increases.

🔍 Example:

Atomic radius trend: Na > Mg > Al > Si > P > S > Cl

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🔸 Down a Group (Top → Bottom)

Observation	Explanation
Atomic radius increases	New energy levels (shells) are added. Even though nuclear charge increases, the shielding effect dominates and electrons are further from the nucleus.

🔍 Example:

Atomic radius trend : F < Cl < Br <I

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🔄 Comparison: Cation vs Anion Radius

Species	Size Compared to Parent Atom	Reason
Cation (+)	Smaller	Loss of electrons → Reduced repulsion & increased Z_eff
Anion (−)	Larger	Gain of electrons → Increased repulsion & reduced Z_eff

🔍 Example:

O > O₂⁻ Na > Na⁺ Cl < Cl⁻

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🔍 Special Cases & Exceptions

• Transition elements: Decrease in radius across the period is less significant due to d-electron shielding.
• Lanthanide contraction: Steady decrease in atomic radius across lanthanides due to poor shielding by f-electrons.

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🧠 Memory Trick for Trends

Direction	Atomic Radius
Across a period →	Decreases ⬇️
Down a group ⬇️	Increases ⬆️

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🔬 Visual Summary

Group↓         Period →
   ↑                 →
   |   ↑ radius       ↓ radius
   ↓

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🧪 JEE/NEET-Level MCQ Patterns

Type	Example
Order-based	Arrange the elements in increasing atomic radius
Assertion-Reason	A: Atomic radius of Na > Cl. R: Na has lower nuclear charge.
Conceptual	Why is Cl⁻ larger than Cl?

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📌 Conclusion

Key Points to Remember
Atomic radius is not fixed, depends on bonding/measurement
Types: Covalent < Metallic < Van der Waals
Trends: Decreases across period, increases down group
Cations are smaller, anions are larger than their neutral atoms

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8.2 ➤ Ionic Radius

🔹 Definition:

Ionic radius is the effective distance from the nucleus of an ion to the outermost shell (where the electrons are likely to be found).

It refers to the size of an ion — either positively charged (cation) or negatively charged (anion).

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🧊 Types of Ions

Type of Ion	Formed by	Charge	Size Compared to Atom	Example
Cation	Loss of electrons	Positive (+)	Smaller than parent atom	Na⁺ < Na
Anion	Gain of electrons	Negative (−)	Larger than parent atom	Cl⁻ > Cl

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🔹 Why Does Size Change in Ions?

👉 Cations (Positive Ions)

• Fewer electrons → reduced electron-electron repulsion
• Same nuclear charge pulls electrons closer
• Size shrinks

Example:
Na (Z = 11): 1s² 2s² 2p⁶ 3s¹
Na⁺: 1s² 2s² 2p⁶ (lost 1 electron)
🔻 Radius decreases

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👉 Anions (Negative Ions)

• More electrons → increased electron-electron repulsion
• Same nucleus cannot hold extra electrons tightly
• Size increases

Example:
Cl (Z = 17): 1s² 2s² 2p⁶ 3s² 3p⁵
Cl⁻: 1s² 2s² 2p⁶ 3s² 3p⁶ (gained 1 e⁻)
🔺 Radius increases

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🔷 Order of Size:

For the same element:

Anion > Atom > Cation

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💠 Isoelectronic Species

🔹 Definition:

Species (atoms or ions) having the same number of electrons but different nuclear charges.

Species	Atomic Number (Z)	Electrons	Type
N³⁻	7	10	Anion
O²⁻	8	10	Anion
F⁻	9	10	Anion
Ne	10	10	Atom
Na⁺	11	10	Cation
Mg²⁺	12	10	Cation
Al³⁺	13	10	Cation

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🔹 Trend in Isoelectronic Series:

As nuclear charge (Z) increases, size (ionic radius) decreases because the same number of electrons are pulled more tightly.

📉 Order of ionic size in isoelectronic series:

N³⁻ > O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺ > Al³⁺

🔍 Why?

• All have 10 electrons
• But increasing Z pulls electrons tighter → smaller radius

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🔬 Graphical Visualization of Trends

Ionic Radius Across Period (in Isoelectronic species)
📉 Decreases with increasing atomic number

Ionic Radius Down a Group
📈 Increases due to addition of new electron shells

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🎯 JEE/NEET Important Points

Concept	Expected Question Type
Cation vs anion size	Conceptual MCQs
Isoelectronic series	Arrange in increasing/decreasing radius
Size of specific ions	Assertion–Reason
Trend in groups/periods	Fill in the blanks, True/False

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🧠 Quick Tricks to Remember

• Cation = compact (think: cat shrinks)
• Anion = airy, bigger
• In isoelectronic series, more protons → smaller ion

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✅ Summary Table

Concept	Key Point
Ionic radius	Size of ion (cation or anion)
Cation (M⁺)	Smaller than atom
Anion (X⁻)	Bigger than atom
Isoelectronic	Same electrons, different protons
Trend in series	↑Z → ↓radius

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8.3 ➤ Ionization Enthalpy (IE)

📘 Definition:

Ionization Enthalpy is the minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.

General Reaction: M_(g) → ​​M⁺_(g) + e⁻   (IE₁)

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🔹 Types of Ionization Enthalpy:

Type	Symbol	Reaction	Description
First Ionization Enthalpy	IE1	M(g) → ​​M(g) + e−	Energy to remove first electron
Second Ionization Enthalpy	IE2	M+(g) → M2+(g)+ e−	Energy to remove second electron
Third Ionization Enthalpy	IE3	M+2(g) → M3(g) + e−	Energy to remove third electron

🔺 Note:
Always:

IE₁ < IE₂ < IE₃, Because after removing one electron, the remaining electrons experience more nuclear attraction (due to increased effective nuclear charge).

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📈 Factors Affecting Ionization Enthalpy

Factor	Effect on IE	Explanation
Atomic Size	↓ Size → ↑ IE	Smaller atoms hold electrons more tightly
Nuclear Charge (Z)	↑ Z → ↑ IE	More protons pull electrons more strongly
Electron Shielding	↑ Shielding → ↓ IE	Inner electrons repel outer electrons
Penetration of Orbitals	s > p > d > f	s-orbitals are closer to nucleus
Stable Configurations	More stable → ↑ IE	Half/full-filled orbitals are stable and resist electron loss

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🧠 Periodic Trends in Ionization Enthalpy

🔹 Across a Period (→):

• IE increases
• Reason: ↓ Atomic size, ↑ Nuclear charge

🔹 Down a Group (↓):

• IE decreases
• Reason: ↑ Atomic size, ↑ Shielding effect, ↓ Nuclear attraction on valence electron

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❗ Exceptions in Ionization Enthalpy Trends

✅ 1. Be vs B

Element	IE₁ (kJ/mol)	Explanation
Be (1s² 2s²)	900	Stable filled 2s orbital
B (1s² 2s² 2p¹)	800	Easier to remove 2p electron (less penetrating)

✅ Be > B (Though B is to the right, Be has higher IE)

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✅ 2. N vs O

Element	IE₁ (kJ/mol)	Explanation
N (1s² 2s² 2p³)	1402	Half-filled stable p-orbital (↑ stability)
O (1s² 2s² 2p⁴)	1314	Extra electron causes repulsion, easier to remove

✅ N > O (Due to half-filled p-orbital stability)

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🔷 Important Points to Remember

• Noble gases have very high IE due to stable octet.
• Alkali metals have lowest IE in their respective periods.
• Transition metals show variable IE due to electron removal from d-orbitals.

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📊 Trend Table (Sample First Ionization Enthalpies)

Element	Atomic No.	Configuration	IE₁ (kJ/mol)
Li	3	1s² 2s¹	520
Be	4	1s² 2s²	900
B	5	1s² 2s² 2p¹	800
C	6	1s² 2s² 2p²	1086
N	7	1s² 2s² 2p³	1402
O	8	1s² 2s² 2p⁴	1314
F	9	1s² 2s² 2p⁵	1681
Ne	10	1s² 2s² 2p⁶	2080

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🎯 JEE/NEET Tips:

Question Type	How it appears
Arrange in order of IE	Based on configuration or trend
Reason-based (Assertion)	e.g., Be has higher IE than B
Conceptual MCQ	Which has lowest IE, or exception questions
Graph/Trend analysis	Compare IE₁ vs Atomic number chart

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📝 Summary Notes

Concept	Key Idea
IE Definition	Energy to remove electron from isolated gaseous atom
IE Order	IE₁ < IE₂ < IE₃
Trend in Period	Increases left to right
Trend in Group	Decreases top to bottom
Exceptions	Be>B, N>O
Influencing Factors	Atomic size, Z, Shielding, Electron configuration

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8.4 ➤ Electron Gain Enthalpy (EGE / EA)

📘 Definition

Electron Gain Enthalpy is the amount of energy released or absorbed when an electron is added to a neutral isolated gaseous atom to form a gaseous anion.

General Reaction: X_(g) + e⁻ → X⁻_(g) + ΔH

• If energy is released, ∆H is negative → Exothermic (favorable).
• If energy is absorbed, ∆H is positive → Endothermic (unfavorable).

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🔹 Electron Gain Enthalpy vs Electron Affinity

• Electron Gain Enthalpy: Thermodynamic term (can be positive or negative).
• Electron Affinity: Generally used to mean the negative of EGE.

For most cases in chemistry, Electron Affinity = – Electron Gain Enthalpy

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🔹 Successive Electron Gain Enthalpies

1st EGE: X_(g) + e⁻ →X⁻_(g) + ΔH_1​​

2nd EGE: X^–_(g) + e⁻ → X²⁻_(g) + ΔH₂

• ∆H₂ is always positive (endothermic) due to repulsion between added electron and already negatively charged ion.

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🧠 Factors Affecting Electron Gain Enthalpy

Factor	Effect	Explanation
Atomic Size	↓ Size → ↑ EA (more negative)	Smaller atoms attract added electron more strongly
Nuclear Charge (Z)	↑ Z → ↑ EA	Stronger attraction of nucleus to incoming electron
Electronic Configuration	Stable config → ↓ EA	Noble gases, Be, N resist extra electrons
Shielding Effect	↑ Shielding → ↓ EA	Weakens the nuclear pull on added electron

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📈 Periodic Trends of Electron Gain Enthalpy

🔹 Across a Period (→)

• Electron gain enthalpy becomes more negative (more energy released).
• Reason: Decreasing atomic size and increasing nuclear charge.

Period 2 Elements	EA₁ Trend
Li < B < C < O < F

🔹 Down a Group (↓)

• Electron gain enthalpy becomes less negative (less energy released).
• Reason: Increasing atomic size and shielding outweighs nuclear charge.

Group 17	EA₁ Trend
F < Cl < Br < I < At

✅ Cl has more negative EGE than F despite being below it! (explained below)

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❗ Anomalies / Exceptions

✅ Fluorine vs Chlorine

Element	EA (kJ/mol)	Why?
F	–328	Very small size → High electron–electron repulsion
Cl	–349	Optimal size → Greater energy release

Cl has the most negative EGE in the periodic table, not F.

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✅ Be, Mg, N, Ne (Group 2, 15, 18)

Element	EGE	Reason
Be, Mg	Positive / near zero	Full s-orbital (2s², 3s²) – stable
N	Low EA	Half-filled p-orbitals → Extra electron causes repulsion
Ne, Ar, Kr	Positive	Full octet, no tendency to accept electrons

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🔎 Numerical Data (1st EA of Some Elements)

Element	Configuration	EA (kJ/mol)
Li	1s² 2s¹	–60
Be	1s² 2s²	~0
B	1s² 2s² 2p¹	–27
C	1s² 2s² 2p²	–122
N	1s² 2s² 2p³	–7
O	1s² 2s² 2p⁴	–141
F	1s² 2s² 2p⁵	–328
Cl	[Ne] 3s² 3p⁵	–349
Ar	[Ne] 3s² 3p⁶	~0

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📊 Summary of Periodic Trends

Trend	Across a Period	Down a Group
Electron Gain Enthalpy	More negative	Less negative
Most Negative EA	Group 17 (Cl)	Decreases from Cl → I

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🔥 JEE/NEET Focus Points

Expected Question Type	What to Master
Arrange elements by EA	Based on size and configuration
Exceptions (F vs Cl, N vs O)	Must be memorized and understood
Assertion–Reason type	e.g., “N has low EA due to half-filled p-orbital”
Conceptual MCQs	Positive vs negative EA, stability reasons

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📝 Key Takeaways

Point	Detail
EGE Definition	Energy change when 1 electron added to gaseous atom
Exothermic EA	Releases energy → EA is negative
Most Negative EA	Chlorine (Cl)
Positive EA	Noble gases, Group 2, Nitrogen
2nd EA	Always positive due to repulsion
F vs Cl	Cl has more negative EA due to lower e⁻–e⁻ repulsion
Trend Across Period	EA becomes more negative
Trend Down Group	EA becomes less negative

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🔁 Revision Mnemonic:

“Small Size, Big Bite!”
→ Smaller atoms bite the incoming electron harder, releasing more energy (negative EA)
But if too small (like F), crowding reduces EA!

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8.5 ➤ Electronegativity (EN)

📘 Definition:

Electronegativity is the tendency of an atom to attract the shared pair of electrons towards itself in a chemical bond.

• It’s not measurable directly (unlike ionization enthalpy or electron gain enthalpy).
• It is a relative property, meaning values are assigned by comparison.

🧪 Important:

• Applies only in bonded atoms (not isolated atoms).
• The concept is crucial in predicting bond polarity, reactivity, and molecular behavior.

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🔹 Difference from Related Terms

Property	Applies To	Describes	Measured in
Electron Gain Enthalpy (EGE)	Isolated atom	Energy released on adding an electron	kJ/mol
Electronegativity	Bonded atom	Pull of electrons in a bond	No unit (relative)

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🔬 Pauling Scale of Electronegativity

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👨‍🔬 Proposed by: Linus Pauling (1932)

• He calculated electronegativity differences based on bond energies.
• The scale is dimensionless (no units).
• The difference in bond energy between heteronuclear and homonuclear molecules was used.

📐 Formula (conceptual):

χA​−χB= (E_AB​− (E_AA​⋅E_BB)^1/2)^1/2​

Where:

• χ = Electronegativity
• E_AB​​ = Bond energy of A–B bond
• E_AA, E_BB​​​ = Bond energies of A–A and B–B

✅ Example Values:

Element	Pauling EN
F	4.0 (Highest)
O	3.5
N	3.0
Cl	3.0
C	2.5
H	2.1
Na	0.9
Cs	0.7 (Lowest)

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📈 Periodic Trends of Electronegativity

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🔹 Across a Period (→)

• Increases left to right.
• Why?
  • Nuclear charge ↑
  • Atomic size ↓
  • Shielding remains relatively constant

➡️ Result: Stronger pull on bonding electrons

Example:
Li (1.0) < Be (1.5) < B (2.0) < C (2.5) < N (3.0) < O (3.5) < F (4.0)

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🔹 Down a Group (↓)

• Decreases top to bottom.
• Why?
  • Atomic size ↑
  • Shielding effect ↑
  • Outer electrons are farther away from nucleus

Example:
F (4.0) > Cl (3.0) > Br (2.8) > I (2.5)

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🔹 Diagonal Relationship

• EN remains nearly constant diagonally across the periodic table.
• Example:
  Li (1.0) and Mg (1.2)
  Be (1.5) and Al (1.5)

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🧠 Factors Affecting Electronegativity

Factor	Effect on EN	Reason
Atomic Radius	↓ radius → ↑ EN	Closer nucleus, stronger pull
Nuclear Charge	↑ charge → ↑ EN	More protons pull more strongly
Shielding Effect	↑ shielding → ↓ EN	Inner electrons reduce effective pull
Hybridization	More s-character → ↑ EN	s-orbitals are closer to nucleus

🔹 Example: sp (50% s) > sp² (33% s) > sp³ (25% s)

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⚠️ Key Exceptions & Points

Case	Explanation
Noble gases	Not assigned EN in some scales (don’t usually form bonds)
Hydrogen	EN = 2.1, behaves variably based on bonding partner
Transition Elements	EN does not change uniformly across a period due to d-electron shielding
Lanthanide Contraction	Causes higher-than-expected EN in later lanthanides

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📊 Comparison of Common EN Scales

Scale	Proposed by	Basis	Max Value
Pauling	Linus Pauling	Bond energy	4.0
Mulliken	Robert Mulliken	Avg of IE & EGE	Varies
Allred-Rochow	Allred & Rochow	Effective nuclear charge & radius	Varies
Sanderson	R.T. Sanderson	Stability ratio	Varies

➡️ Pauling scale is the most commonly used and accepted in exams.

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💡 Applications of Electronegativity

Application	Description
Bond Polarity	Greater difference in EN → more polar bond
Ionic vs Covalent Character	Large ΔEN → ionic bond; small ΔEN → covalent
Acidic/Basic Nature	Higher EN → more acidic oxides (non-metals), Lower EN → basic oxides (metals)
Reactivity Trends	Halogens (high EN) → reactive nonmetals, Alkali metals (low EN) → reactive metals

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🧪 Trick to Remember Electronegativity Order (for Period 2)

Hi Be Boring Cat Nap Our Friend
H < Be < B < C < N < O < F

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🧮 NEET/JEE Question Types

Question Style	Example
Order of EN	Arrange O, N, F in increasing EN
Predict bond character	Which is more ionic: NaCl or MgCl₂?
Reason-based	Why EN of N is less than O?
Conceptual	Why is EN not defined for noble gases?

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📝 Quick Summary

Concept	Trend
Highest EN	Fluorine (F) = 4.0
Lowest EN	Caesium (Cs) = 0.7
Across Period	Increases (L → R)
Down Group	Decreases
Depends on	Size ↓, Z ↑, Shielding ↓
Used to	Predict bond polarity, reactivity

8.6 ➤ Valency (Also spelled as Valence)

🔷 Definition

Valency is the combining capacity of an element, i.e., the number of electrons an atom loses, gains, or shares to achieve a stable noble gas configuration (usually an octet or duplet).

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🔹 Valency Based on Octet Rule

Atoms try to achieve 8 electrons in the outer shell (octet) (or 2 electrons for hydrogen and helium, i.e., duplet).

• Metals (e.g., Na, Mg): Lose electrons → Electropositive valency
• Nonmetals (e.g., Cl, O): Gain/share electrons → Electronegative valency

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🔷 Types of Valency

Type	Definition	Example
Electropositive Valency	No. of electrons lost	Na (Z=11): 1s²2s²2p⁶3s¹ → loses 1e⁻ → Valency = 1
Electronegative Valency	No. of electrons gained	Cl (Z=17): 1s²2s²2p⁶3s²3p⁵ → gains 1e⁻ → Valency = 1
Covalency	No. of electrons shared	C (Z=6): shares 4e⁻ (2s²2p²) → Valency = 4

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🔷 Valency and Group Number

Group	Valency	How?
Group 1 (IA)	1	Lose 1 electron (ns¹)
Group 2 (IIA)	2	Lose 2 electrons (ns²)
Group 13 (IIIA)	3	Lose 3 electrons (ns²np¹)
Group 14 (IVA)	4	Share 4 electrons (ns²np²)
Group 15 (VA)	3	Gain/share 3 electrons (ns²np³)
Group 16 (VIA)	2	Gain/share 2 electrons (ns²np⁴)
Group 17 (VIIA)	1	Gain/share 1 electron (ns²np⁵)
Group 18 (VIIIA)	0	Octet complete (ns²np⁶)

🧠 Mnemonic (For p-block):

Valency = 8 – group number (for nonmetals in groups 15–17)
Valency = group number (for groups 1–4)

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🔹 For Transition Elements (d-block)

• Show variable valency due to involvement of both (n-1)d and ns electrons.
• Example:
  • Fe: Valency = 2, 3 → [Ar] 3d⁶4s²
  • Cu: Valency = 1, 2 → [Ar] 3d¹⁰4s¹

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🔹 For f-block Elements (Lanthanides/Actinides)

• Often have valency = 3
• May show variable valency due to (n-2)f, (n-1)d, and ns electrons

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🔷 Periodic Trends in Valency

1️⃣ Across a Period (→)

• Valency first increases from 1 to 4, then decreases back to 0
• Pattern: 1, 2, 3, 4, 3, 2, 1, 0
• Reason: Elements tend to gain or lose electrons to achieve an octet

Period 2	Li	Be	B	C	N	O	F	Ne
Valency	1	2	3	4	3	2	1	0

2️⃣ Down a Group (↓)

• Valency remains constant
• Because elements in a group have same number of valence electrons
• Example:

Group 1	Li	Na	K	Rb	Cs
Valency	1	1	1	1	1

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🧠 Special Concepts & Exceptions

✅ Variable Valency

• Exhibited by elements with d and f orbitals
• Cause: Small energy difference between ns and (n–1)d orbitals
• Ex: Fe²⁺, Fe³⁺; Cu¹⁺, Cu²⁺

✅ Zero Valency

• Noble gases (He, Ne, Ar…) have complete octet → Valency = 0
• Exception: Some noble gases form compounds (e.g., XeF₂) under special conditions

✅ Hydrogen

• Valency = 1 (can act like Group 1 or 17 element)
• Forms both H⁺ and H⁻ depending on the compound (e.g., HCl vs. NaH)

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🔥 JEE & NEET Focused Questions on Valency

Type	Examples
Arrange elements by valency across a period	Li < Be < B < C…
Which element has variable valency?	Fe, Cu, Cr
Noble gas with non-zero valency?	Xe
Identify valency from configuration	2s²2p³ → Valency = 3
Match group number to valency	Group 16 → Valency = 2

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📊 Quick Summary Table: Periodic Table Valency Trend

Group	Valence Electrons	Typical Valency	Type
1 (IA)	1	1	Electropositive
2 (IIA)	2	2	Electropositive
13 (IIIA)	3	3	Electropositive
14 (IVA)	4	4	Covalent
15 (VA)	5	3	Electronegative
16 (VIA)	6	2	Electronegative
17 (VIIA)	7	1	Electronegative
18 (VIIIA)	8	0	Inert

✅ Conclusion

• Valency is crucial for predicting chemical bonding, compound formation, and reactivity.
• A deep understanding of electron configuration helps to predict valency and identify exceptions.
• Mastery of valency trends is essential for solving structure-based, formula-based, and reaction-based MCQs in JEE/NEET.

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9. Anomalous Trends & Exceptions in Periodic Properties

– Including Irregularities & Diagonal Relationship (NCERT + Advanced Level)

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✳️ What are Anomalous Trends?

Anomalies are deviations from expected periodic trends caused by factors like:

• Small atomic/ionic size
• High electronegativity
• Unique electron configuration
• Poor shielding (especially in d- & f-blocks)

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🔷 1. Anomalies in First Element of Each Group (Period 2)

The first element of Groups 1 to 17 shows anomalous behavior compared to the rest of its group.

✴️ Example: Li vs Other Alkali Metals

Property	Lithium (Li)	Sodium (Na) & others	Cause of Difference
Ionic Radius	Small	Larger	Small atomic size
Polarizing Power	High	Lower	High charge density
Reactivity	Less reactive with air/water	More reactive	Forms protective oxide layer
Nature of Compound	More covalent (LiCl)	More ionic (NaCl)	Fajan’s Rule

🔍 Reason:

• Li, Be, B, C, N, O, F have small sizes, high electronegativity, and high IE
• Their valence shells are in the second shell (n = 2) → No d-orbitals → No expansion of octet

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🔷 2. Diagonal Relationship (Unique Anomaly)

Diagonal Relationship: A phenomenon where the first element of a group (Period 2) shows similarities with the second element of the next group (Period 3) diagonally placed in the periodic table.

🧪 Examples of Diagonal Pairs

Period 2 Element	Diagonally Related (Period 3)
Li (Group 1)	Mg (Group 2)
Be (Group 2)	Al (Group 13)
B (Group 13)	Si (Group 14)

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🧠 Why Does Diagonal Relationship Occur?

Because of similarities in:

Factor	Explanation
Ionic/Atomic Radius	Nearly equal radii despite being in different groups
Electronegativity	Close values → Similar chemical bonding
Charge Density (Z/r)	Similar polarizing power
Ionization Enthalpy	Close values due to balance of size & nuclear charge

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🧪 Detailed Examples

✅ Li vs Mg

Property	Li	Mg
Ionic Radius	76 pm	72 pm
Forms Nitrides	Li₃N	Mg₃N₂
Forms Oxides	Li₂O (not peroxide)	MgO
Solubility	Li₂CO₃ decomposes on heating	MgCO₃ also decomposes

✅ Be vs Al

Property	Be	Al
Amphoteric nature	Be(OH)₂	Al(OH)₃
Covalent Chloride	BeCl₂	AlCl₃
Tendency to form complex ions	[BeF₄]²⁻	[AlF₆]³⁻

✅ B vs Si

Property	B	Si
Semiconductor	Yes	Yes
Acidic oxides	B₂O₃	SiO₂
Forms hydrides	B₂H₆	SiH₄

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🔷 3. Other Anomalous Trends in Periodic Properties

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✅ a) Ionization Enthalpy Exception

• Be > B
  Be = 1s² 2s² (stable full 2s)
  B = 1s² 2s² 2p¹ (easier to remove 2p electron)
• N > O
  N = 2p³ (half-filled stable)
  O = 2p⁴ (extra repulsion)

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✅ b) Electron Gain Enthalpy Exception

Element	EG Enthalpy	Reason
F < Cl	F is smaller → high electron repulsion	Fluorine’s small atomic size leads to tightly packed electrons, causing strong repulsion when an extra electron is added.
O < S	Same reason: smaller O leads to less favorable electron gain	Smaller atoms like F and O have high electron density, causing strong repulsion that reduces electron gain enthalpy.

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✅ c) Electronegativity Irregularities

• F is most electronegative
• O > Cl, even though Cl is to the right
  → O is smaller and forms stronger polar covalent bonds

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✅ d) Atomic/Ionic Radius

• Transition Elements: Show anomalous decrease in atomic size due to poor shielding by d-electrons
• Lanthanide Contraction:
  • Across lanthanides: atomic size decreases slowly
  • Causes Zr and Hf to have nearly same size

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🔷 4. Lanthanide Contraction (Another Anomaly)

Gradual decrease in atomic and ionic size from La (Z=57) to Lu (Z=71)

| Cause | Poor shielding by 4f electrons |
| Result | Elements after lanthanides (Hf, Ta) have similar sizes to those before them (Zr, Nb)|

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📌 Summary Table: Anomalous Behaviors

Anomaly	Example	Reason
Li ≠ other alkali metals	Li forms Li₃N	Small size, high polarizing power
Be ≠ other alkaline earth metals	BeCl₂ covalent	Small size, high IE
Be ≈ Al	Amphoteric nature	Diagonal relationship
B ≈ Si	Covalent oxides	Diagonal relationship
IE anomaly	Be > B, N > O	Stable electronic configuration
Electron gain enthalpy	Cl > F	F is too small, high repulsion
Lanthanide contraction	Zr ≈ Hf	Poor shielding by 4f

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🧠 Tips for JEE/NEET:

1. Memorize diagonal pairs (Li–Mg, Be–Al, B–Si)
2. Focus on exceptions in IE, EA, EN
3. Use concepts like shielding, penetration, and stability
4. Expect MCQs or Assertion-Reason questions from anomalies

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10. Periodicity and Reactivity of Elements

📘 What is Periodicity ?

Periodicity is the repetition of chemical and physical properties of elements at regular intervals when arranged by increasing atomic number.

These repeating trends help us predict:

• Reactivity
• Metallic and non-metallic character
• Oxidizing/reducing nature

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🧲 Reactivity and Periodicity

Reactivity refers to how easily an element loses or gains electrons to form chemical bonds.

For Metals	Reactivity ∝ Ease of losing electrons
For Non-Metals	Reactivity ∝ Ease of gaining electrons

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🔷 1. Metallic and Non-Metallic Character

📌 Metallic Character

Tendency of an atom to lose electrons and form positive ions (cations)

• Shows electropositive nature
• Forms basic oxides
• Found mostly on left side and center of periodic table

📌 Non-Metallic Character

Tendency of an atom to gain electrons and form negative ions (anions)

• Shows electronegative nature
• Forms acidic oxides
• Found on right side of periodic table (except noble gases)

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🔄 Periodic Trends

Direction	Metallic Character	Non-Metallic Character
Across a Period (→)	Decreases	Increases
Down a Group (↓)	Increases	Decreases

🔍 Reason:

• Across a period: Atomic size ↓, Ionization enthalpy ↑ → Harder to lose electrons
• Down a group: Atomic size ↑, Ionization enthalpy ↓ → Easier to lose electrons

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🧪 Examples:

Element	Metallic Character
Na > Mg > Al	More metallic (Group 1 > 2 > 13)
P < S < Cl	Increasing non-metallic character
Li < Na < K < Rb < Cs	Increasing metallic character

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🔷 2. Reactivity Trends in Groups and Periods

🔹 For Metals (like Group 1 and 2)

Trend	Reactivity
Down the group	Increases
Across a period	Decreases

✅ Why?

• Larger atoms → lower ionization energy → easily lose electrons

🧪 Example:
K > Na > Li (more reactive down the group)

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🔹 For Non-Metals (like Group 16 and 17)

Trend	Reactivity
Down the group	Decreases
Across a period	Increases

✅ Why?

• Smaller atoms with high nuclear charge attract electrons better

🧪 Example:
F > Cl > Br > I (reactivity decreases down Group 17)

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🔷 3. Oxidizing and Reducing Nature

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🔴 Oxidizing Agents

Substances that gain electrons themselves and oxidize others.

• Typically non-metals (e.g. halogens)
• Higher electron affinity = stronger oxidizing agent

Example	Reaction
F₂ + 2e⁻ → 2F⁻	Fluorine is a strong oxidizing agent

🔺 Trend in Oxidizing Power (Group 17): F₂​>Cl₂​>Br₂​>I_2​

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🔵 Reducing Agents

Substances that lose electrons themselves and reduce others.

• Typically metals
• Lower ionization energy = stronger reducing agent

Example	Reaction
Na → Na⁺ + e⁻	Sodium is a strong reducing agent

🔻 Trend in Reducing Power (Group 1): Cs>Rb>K>Na>Li

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🔄 Summary Table: Periodicity of Key Properties

Property	Across a Period (→)	Down a Group (↓)
Atomic Size	Decreases	Increases
Ionization Enthalpy	Increases	Decreases
Electronegativity	Increases	Decreases
Metallic Character	Decreases	Increases
Non-Metallic Character	Increases	Decreases
Metal Reactivity	Decreases	Increases
Non-Metal Reactivity	Increases	Decreases
Oxidizing Power (Non-Metals)	Increases	Decreases
Reducing Power (Metals)	Decreases	Increases

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📌 JEE/NEET Key Takeaways:

• Metals lose electrons easily → good reducing agents
• Non-metals gain electrons easily → good oxidizing agents
• F is the strongest oxidizing agent
• Li is the strongest reducing agent in aqueous solution due to high hydration energy
• Watch for trends across periods and groups – most questions test exceptions or comparisons

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11. Applications of the Periodic Table

The Periodic Table is a powerful tool that helps chemists predict and understand the chemical behavior of elements based on their position in the table.

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🔷 1. Predicting Type of Bonding

The type of bond formed between atoms depends on their electronegativity difference and positions in the periodic table.

🧪 Bond Types:

Type of Bond	Electronegativity Difference	Example	Where Found
Ionic	Large difference (≥1.7)	NaCl	Metal + Non-metal
Covalent	Small difference (<1.7)	H₂O, Cl₂	Non-metal + Non-metal
Metallic	Between metals	Fe, Cu	Between metals (free electrons)

📌 Use Periodic Table to:

• Identify metals vs non-metals
• Predict bonding type:
  • Left + Right → Ionic
  • Right + Right → Covalent
  • Left + Left → Metallic

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🔷 2. Nature of Oxides and Hydrides

Based on an element’s position, we can predict the nature of its oxides and hydrides.

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🔹 Oxides:

Type of Element	Nature of Oxide	Example
Metals (Groups 1, 2)	Basic	Na₂O, CaO
Metalloids (Group 13, 14)	Amphoteric	Al₂O₃, SnO
Non-metals (Group 15–17)	Acidic	SO₂, P₂O₅

🔁 Trend:

• Across period (left → right): Basic → Amphoteric → Acidic
• Down a group: Basicity increases (metallic nature increases)

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🔹 Hydrides:

Type of Element	Nature of Hydride	Example
Group 1 & 2 metals	Ionic (saline)	NaH, CaH₂
p-Block non-metals	Covalent	CH₄, NH₃, HCl
Transition metals	Interstitial	TiH₂

🧠 Hydrides’ nature helps infer:

• Bonding type
• Reducing ability
• Thermal stability

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🔷 3. Predicting Oxidation States

Oxidation state = apparent charge of atom after gaining/losing electrons.

📌 Trends in Oxidation States:

Group	General Oxidation States
1	+1
2	+2
13	+3, sometimes +1 (inert pair effect)
14	+4, +2 (inert pair effect in heavier elements)
15	-3, +3, +5
16	-2, +4, +6
17	-1, +1, +3, +5, +7
Transition metals	Variable oxidation states (due to d-electrons)

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🔁 Periodic Trends:

• s- and p-block elements: predictable oxidation states based on group number
• d-block (transition elements): multiple oxidation states due to similar energies of ns and (n–1)d electrons
• f-block (lanthanides/actinides): mostly +3, but variable for actinides

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🧠 Inert Pair Effect:

• Seen in Group 13, 14, 15 (heavier elements)
• Tendency to show lower oxidation states than expected due to reluctance of s-electrons to participate in bonding
• Example: Tl⁺ more stable than Tl³⁺, Pb²⁺ more stable than Pb⁴⁺

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🔷 Summary Table

Application	What to Predict	How Periodic Table Helps
Bonding	Ionic, covalent, metallic	Electronegativity, position (metal/non-metal)
Oxides	Acidic/basic/amphoteric	Metallic character → nature of oxide
Hydrides	Ionic, covalent, interstitial	Group type + bonding nature
Oxidation States	Valency & possible charges	Group number, inert pair effect, block type

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🧪 JEE/NEET Tips:

• JEE: Focus on exceptions in oxidation states (e.g. transition metals, inert pair effect).
• NEET: Focus on general trends and predictability in bonding & oxide nature.
• Always link periodic trends (atomic size, IE, electronegativity) to chemical behavior.

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Thank You

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