Periodic trends diagram for Class 11 Chemistry Chapter 3 showing atomic radius, ionization energy, electronegativity, and metallic character

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

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All The Topics Are-

S. No.Topic NameSubtopics Name
1Introduction– Need for classification- Historical background
2Dobereiner’s Triads– Concept- Examples- Limitations
3Newlands’ Law of Octaves– Law of Octaves- Examples- Limitations
4Mendeleev’s Periodic Table– Basis of classification- Merits- Limitations
5Modern Periodic Law & Table– Modern Periodic Law- Structure of modern table- Groups, Periods, Blocks- Types of elements
6Nomenclature (Z > 100)– IUPAC rules- Naming conventions
7Electronic Configuration & Periodic Table– Relation with group & period- Valence electrons- Block classification
8Periodic Trends in Properties
8.1➤ Atomic Radius– Covalent, Metallic, Van der Waals radius- Periodic and group trends
8.2➤ Ionic Radius– Cations and anions- Isoelectronic species
8.3➤ Ionization Enthalpy– 1st & 2nd IE- Factors- Exceptions
8.4➤ Electron Gain Enthalpy– Definition- Trends- Anomalies
8.5➤ Electronegativity– Pauling scale- Periodic and group trends
8.6➤ Valency– Relation with group number- Periodic trend
9Anomalous Trends & Exceptions– Irregularities- Diagonal relationship
10Periodicity and Reactivity– Metallic/non-metallic character- Reactivity trends- Oxidizing/reducing nature
11Applications of Periodic Table– Predicting bonding- Oxides/hydrides- Oxidation states
12Question Types & Practice– Assertion-Reason- MCQs- Numericals

1. Introduction – Classification of Elements and Periodicity in Properties

🧠 Topics Covered:

  • Introduction
  • Need for Classification
  • Historical Background

🔷 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.


🔷 2. Need for Classification

Why classify elements?

ReasonExplanation
🌱 Growing number of elementsBy the 18th century, nearly 60 elements were known; now over 118. A system was required to manage them.
🧪 Understanding chemical propertiesGrouping elements with similar properties made it easier to predict their behavior.
🔍 Avoiding repetitionWithout a pattern, each element had to be studied individually. Classification revealed periodic repetition in properties.
📚 Simplifying studyClassification helped in structuring chemistry textbooks and made learning more organized.
🔄 Revealing trendsClassification showed how properties like atomic size, ionization energy, etc., change periodically.

Analogy:

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


🔷 3. Historical Background

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


🧪 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.

🔹 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.

⚗️ 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.

🔬 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.

🧠 Summary Chart:

ScientistYearBasisSuccessesFailures
Dobereiner1817Atomic mass (Triads)Atomic mass relationshipsLimited applicability
Newlands1865Atomic mass (Octaves)Repeating propertiesFailed after Ca
Mendeleev1869Atomic mass (Table)Predicted new elementsIsotope anomaly
Moseley1913Atomic numberCorrected earlier errors

✅ 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.


🧠 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.

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.


🔷 2. Examples of Dobereiner’s Triads

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


Triad 1: Alkali Metals

ElementSymbolAtomic Mass
LithiumLi7
SodiumNa23
PotassiumK39

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.


Triad 2: Alkaline Earth Metals

ElementSymbolAtomic Mass
CalciumCa40
StrontiumSr88
BariumBa137

Check: (40+137) / 2 = 88.5 ≈ 88=Sr

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


Triad 3: Halogens

ElementSymbolAtomic Mass
ChlorineCl35.5
BromineBr80
IodineI127

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

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


🔷 3. Graphical Representation

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


🔷 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.

🔴 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.


🔷 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.

🧠 Final Summary

FeatureDobereiner’s Triads
Proposed byJohann Dobereiner (1817)
BasisAtomic mass and chemical similarity
PatternMiddle element’s atomic mass ≈ average of the other two
SignificanceEarly step toward periodic classification
LimitationsCouldn’t explain all elements or predict new ones

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.


🔷 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.


🔷 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

SaReGaMaPaDhaNi
HLiBeBCNO
FNaMgAlSiPS
ClKCa

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.


🔷 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.

🔴 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 columnExample: Co & Ni grouped with halogens like F and Cl (which was chemically incorrect).
No room for undiscovered elementsDid not leave gaps; forced elements into the pattern.
Works only for lighter elementsHeavier elements didn’t follow the octave rule.
No distinction for noble gasesNoble gases weren’t discovered yet; inclusion would’ve disrupted the pattern.

🔷 6. Comparison with Dobereiner’s Triads

FeatureDobereiner’s TriadsNewlands’ Octaves
Year18171865
PatternAtomic mass avg. in triadsEvery 8th element similar
Number of elements coveredVery fewUp to Calcium
Major limitationVery limited triadsFailed for heavier elements

🔷 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.

🧠 Final Summary

AspectNewlands’ Law of Octaves
Proposed byJohn Newlands (1865)
BasisIncreasing atomic mass
Key ideaEvery 8th element shows similar properties
Number of elements explainedUp to Calcium
SignificanceFirst idea of periodicity
LimitationFailed beyond Ca, no gaps, dissimilar elements in groups

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.


🔷 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.


🧱 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.

⚙️ Features of Mendeleev’s Table:

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

🔷 2. Merits of Mendeleev’s Periodic Table

Mendeleev’s work was revolutionary and had several advantages:

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

🔷 3. Examples of Mendeleev’s Predictions

Predicted ElementActual ElementProperties predicted
Eka-siliconGermanium (Ge)Density, atomic mass, oxide formula
Eka-aluminiumGallium (Ga)Melting point, reactivity
Eka-boronScandium (Sc)Oxidation state, density

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


🔴 4. Limitations of Mendeleev’s Periodic Table

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

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

🔷 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.

🔷 6. Comparison: Mendeleev vs Modern Table

FeatureMendeleev’s TableModern Table
BasisAtomic MassAtomic Number
Periods67
Groups818
GapsYes (for new elements)No gaps
IsotopesNot explainedExplained
Noble gasesAbsent initiallyPresent in Group 18

🧠 Final Summary

📌 Aspect🔍 Mendeleev’s Table
Proposed byDmitri Mendeleev (1869)
Classification BasisIncreasing atomic mass
Structure8 groups, 6 periods
Key LawPeriodic Law (properties repeat with atomic mass)
StrengthPrediction of elements, systematic grouping
WeaknessAnomalies, isotopes, noble gases excluded
LegacyBasis for modern periodic table

🎯 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

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.

🔷 2. STRUCTURE OF THE MODERN PERIODIC TABLE

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

📊 Key Features:

FeatureDescription
🔢 BasisAtomic number (Z)
🔄 PeriodicityDue to repetition of electronic configurations
📏 Total Elements118 known elements (as of 2025)
🧬 Division7 periods and 18 groups
🧱 Blockss-block, p-block, d-block, f-block
🌈 Zig-zag lineSeparates metals from nonmetals

🔷 3. PERIODS IN MODERN PERIODIC TABLE

📐 Horizontal Rows = Periods

There are 7 periods in the modern periodic table:

PeriodNo. of ElementsStarts WithEnds With
1st2H (Z=1)He (Z=2)
2nd8Li (Z=3)Ne (Z=10)
3rd8Na (Z=11)Ar (Z=18)
4th18K (Z=19)Kr (Z=36)
5th18Rb (Z=37)Xe (Z=54)
6th32Cs (Z=55)Rn (Z=86)
7th32 (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.

🔷 4. GROUPS IN MODERN PERIODIC TABLE

📏 Vertical Columns = Groups

There are 18 groups in total:

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

💡 Same group elements have:

  • Same valence shell configuration
  • Similar chemical properties

🔷 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:

BlockElementsLast e⁻ entersGroups Covered
s-blockAlkali & Alkaline Earth Metals + Hes-orbitalGroups 1 & 2
p-blockNon-metals, Halogens, Noble Gases, Metalloidsp-orbitalGroups 13–18
d-blockTransition Metalsd-orbitalGroups 3–12
f-blockLanthanides & Actinidesf-orbitalNot in main table (placed separately)

🔷 6. TYPES OF ELEMENTS IN MODERN PERIODIC TABLE

Based on position & properties, elements are categorized as:


🟥 A. Metals

FeatureDescription
LocationLeft and center of the table
Number≈ 78% of all elements
PropertiesMalleable, ductile, conductors, electropositive, lose e⁻ easily
Blockss, d, f, and some p-block

🟨 B. Non-Metals

FeatureDescription
LocationUpper right corner
PropertiesBrittle, poor conductors, electronegative, gain e⁻
ExamplesH, C, N, O, F, P, S, Cl, Br, I

🟩 C. Metalloids

FeatureDescription
LocationAlong the “stair-step” zig-zag line
PropertiesIntermediate between metals and non-metals
ExamplesB, Si, Ge, As, Sb, Te

🟦 D. Noble Gases

FeatureDescription
LocationGroup 18
PropertiesStable, inert, monoatomic gases
ExamplesHe, Ne, Ar, Kr, Xe, Rn

🟪 E. Transition Elements

FeatureDescription
Locationd-block (Groups 3–12)
PropertiesVariable oxidation states, colored compounds, form complexes

🟫 F. Inner Transition Elements

FeatureDescription
Locationf-block (bottom of table)
SeriesLanthanides (Z=58–71), Actinides (Z=90–103)
PropertiesMostly radioactive (especially actinides), large atomic size

🔷 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.


🧠 FINAL SUMMARY

🔑 Concept📘 Modern Periodic Table
Proposed byHenry Moseley (1913)
Based onAtomic number
No. of periods7
No. of groups18
Classifications, p, d, f blocks
Types of elementsMetals, Nonmetals, Metalloids, Noble Gases, Transition & Inner Transition Elements
Periodicity due toRepetition of similar valence shell configurations

🎯 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).

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.

🔷 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.


📌 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).

🔷 Table: Roots for Each Digit

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

🧩 Special Guidelines:

SituationRule
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 allowedExcept “ii” (remove one “i”)
SymbolFormed by taking first letter of each root, capitalizing the first letter (e.g., Unq for Unnilquadium)

🔷 Examples with Explanation


✅ 1. Element 104

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

🧾 Now permanently named as Rutherfordium (Rf).


✅ 2. Element 105

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

🧾 Now known as Dubnium (Db).


✅ 3. Element 106

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

🧾 Now named Seaborgium (Sg).


✅ 4. Element 118

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

🧾 Now named Oganesson (Og).


✅ 5. Element 109

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

🧾 Now named Meitnerium (Mt).


✅ 6. Element 113

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

🧾 Now named Nihonium (Nh).


🔷 Final Names vs. Systematic Names

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

🔷 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.

🧠 Final Summary

FeatureDetail
Applies toElements with Z > 100
Based onAtomic number digits
UsesLatin/Greek roots for each digit
Ends with“-ium”
SymbolFirst letters of the roots
PurposeTemporary, systematic names until official names are assigned

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

🔷 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 TypeGroup 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

📌 B. Period Number and Electronic Configuration

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

ElementConfigurationPeriod
Na (11)3s¹3
Cl (17)3s² 3p⁵3
K (19)4s¹4
Fe (26)3d⁶ 4s²4

🔷 3. Valence Electrons

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

  • Group number
  • Valency
  • Reactivity
BlockValence Electrons
s-blockns¹–²
p-blockns² 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

🔷 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

📘 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

📘 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

📘 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

🔷 5. Trends Based on Electronic Configuration

TrendExplanation
ValencyBased on number of valence electrons
Atomic RadiusIncreases down group, decreases across period
Ionization EnergyDecreases down group, increases across period
ElectronegativityDecreases down group, increases across period
Metallic characterIncreases down group, decreases across period

🔷 6. Summary Table

ConceptExplanationExample
Group NumberRelated to number of valence electronsMg (3s²) → Group 2
Period NumberBased on highest n valueCl (3s²3p⁵) → Period 3
Valence ElectronsOuter shell electronsO (2s² 2p⁴) → 6 valence electrons
BlockBased on orbital receiving last electronFe (3d⁶ 4s²) → d-block
Type of Elements = active metals, p = diverse, d = transition, f = inner transition

🔷 Visual Classification (Quick Reference)

BlockGroupsLast Electron EntersExampleCharacteristics
s1-2s-orbitalNa (3s¹)Soft, reactive metals
p13-18p-orbitalO (2p⁴)Nonmetals, metalloids
d3-12d-orbitalFe (3d⁶)Transition metals
ff-orbitalU (5f³)Rare earth metals

🎯 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

🧠 Final Summary

ConceptKey Points
Electronic ConfigurationDetermines position in periodic table
GroupBased on valence electrons
PeriodBased on principal quantum number (n)
BlockBased on type of orbital receiving last electron
Periodic TrendsArise 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.


🧠 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

📊 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

🔷 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

🔷 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.

🔷 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.

🔷 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).

🔷 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


🔷 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.

🔷 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)

🔷 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).

📌 Summary Table of Periodic Trends

PropertyAcross Period (→)Down Group (↓)
Atomic RadiusDecreasesIncreases
Ionic RadiusDecreasesIncreases
Ionization EnthalpyIncreasesDecreases
Electron Gain EnthalpyMore NegativeLess Negative
ElectronegativityIncreasesDecreases
Metallic CharacterDecreasesIncreases
Non-Metallic CharacterIncreasesDecreases
ValencyIncreases then DecreasesConstant
Reactivity (Metals)DecreasesIncreases
Reactivity (Non-metals)IncreasesDecreases

💡 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.

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:


🧪 Types of Atomic Radii

TypeDefinitionApplicable ToSymbol
Covalent RadiusHalf the distance between the nuclei of two covalently bonded atoms of the same elementNon-metals (e.g. H₂, Cl₂, O₂)rcov
Metallic RadiusHalf the distance between nuclei of two adjacent atoms in the metallic latticeMetals (e.g. Na, Fe, Cu)rmetal
Van der Waals RadiusHalf the distance between two non-bonded atoms in neighboring moleculesNoble gases or unbonded atoms (e.g. Ar, Ne)rvdW

🔹 Comparison of Radii (Generally):

rvdW​ > rmetal ​> rcov​

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


📏 Typical Examples

ElementCovalent Radius (pm)Metallic Radius (pm)Van der Waals Radius (pm)
Hydrogen (H)37120
Chlorine (Cl)99180
Sodium (Na)186
Argon (Ar)188

📈 Trends in Atomic Radius


🔸 Across a Period (Left → Right)

ObservationExplanation
Atomic radius decreasesAs 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) increasesShielding remains nearly constant, but the pull increases.

🔍 Example:

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


🔸 Down a Group (Top → Bottom)

ObservationExplanation
Atomic radius increasesNew 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


🔄 Comparison: Cation vs Anion Radius

SpeciesSize Compared to Parent AtomReason
Cation (+)SmallerLoss of electrons → Reduced repulsion & increased Z_eff
Anion (−)LargerGain of electrons → Increased repulsion & reduced Z_eff

🔍 Example:

O > O2 Na > Na+ Cl < Cl


🔍 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.

🧠 Memory Trick for Trends

DirectionAtomic Radius
Across a period →Decreases ⬇️
Down a group ⬇️Increases ⬆️

🔬 Visual Summary

Group↓         Period →
↑ →
| ↑ radius ↓ radius


🧪 JEE/NEET-Level MCQ Patterns

TypeExample
Order-basedArrange the elements in increasing atomic radius
Assertion-ReasonA: Atomic radius of Na > Cl. R: Na has lower nuclear charge.
ConceptualWhy is Cl⁻ larger than Cl?

📌 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

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).


🧊 Types of Ions

Type of IonFormed byChargeSize Compared to AtomExample
CationLoss of electronsPositive (+)Smaller than parent atomNa⁺ < Na
AnionGain of electronsNegative (−)Larger than parent atomCl⁻ > Cl

🔹 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


👉 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


🔷 Order of Size:

For the same element:

Anion > Atom > Cation


💠 Isoelectronic Species

🔹 Definition:

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

SpeciesAtomic Number (Z)ElectronsType
N³⁻710Anion
O²⁻810Anion
F⁻910Anion
Ne1010Atom
Na⁺1110Cation
Mg²⁺1210Cation
Al³⁺1310Cation

🔹 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

🔬 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


🎯 JEE/NEET Important Points

ConceptExpected Question Type
Cation vs anion sizeConceptual MCQs
Isoelectronic seriesArrange in increasing/decreasing radius
Size of specific ionsAssertion–Reason
Trend in groups/periodsFill in the blanks, True/False

🧠 Quick Tricks to Remember

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

✅ Summary Table

ConceptKey Point
Ionic radiusSize of ion (cation or anion)
Cation (M⁺)Smaller than atom
Anion (X⁻)Bigger than atom
IsoelectronicSame electrons, different protons
Trend in series↑Z → ↓radius

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)


🔹 Types of Ionization Enthalpy:

TypeSymbolReactionDescription
First Ionization EnthalpyIE1M(g) → ​​M(g) + e−  Energy to remove first electron
Second Ionization EnthalpyIE2M+(g) → M2+(g)+ eEnergy to remove second electron
Third Ionization EnthalpyIE3M+2(g) → M3(g) + eEnergy 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).


📈 Factors Affecting Ionization Enthalpy

FactorEffect on IEExplanation
Atomic Size↓ Size → ↑ IESmaller atoms hold electrons more tightly
Nuclear Charge (Z)↑ Z → ↑ IEMore protons pull electrons more strongly
Electron Shielding↑ Shielding → ↓ IEInner electrons repel outer electrons
Penetration of Orbitalss > p > d > fs-orbitals are closer to nucleus
Stable ConfigurationsMore stable → ↑ IEHalf/full-filled orbitals are stable and resist electron loss

🧠 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

Exceptions in Ionization Enthalpy Trends

✅ 1. Be vs B

ElementIE₁ (kJ/mol)Explanation
Be (1s² 2s²)900Stable filled 2s orbital
B (1s² 2s² 2p¹)800Easier to remove 2p electron (less penetrating)

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


✅ 2. N vs O

ElementIE₁ (kJ/mol)Explanation
N (1s² 2s² 2p³)1402Half-filled stable p-orbital (↑ stability)
O (1s² 2s² 2p⁴)1314Extra electron causes repulsion, easier to remove

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


🔷 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.

📊 Trend Table (Sample First Ionization Enthalpies)

ElementAtomic No.ConfigurationIE₁ (kJ/mol)
Li31s² 2s¹520
Be41s² 2s²900
B51s² 2s² 2p¹800
C61s² 2s² 2p²1086
N71s² 2s² 2p³1402
O81s² 2s² 2p⁴1314
F91s² 2s² 2p⁵1681
Ne101s² 2s² 2p⁶2080

🎯 JEE/NEET Tips:

Question TypeHow it appears
Arrange in order of IEBased on configuration or trend
Reason-based (Assertion)e.g., Be has higher IE than B
Conceptual MCQWhich has lowest IE, or exception questions
Graph/Trend analysisCompare IE₁ vs Atomic number chart

📝 Summary Notes

ConceptKey Idea
IE DefinitionEnergy to remove electron from isolated gaseous atom
IE OrderIE₁ < IE₂ < IE₃
Trend in PeriodIncreases left to right
Trend in GroupDecreases top to bottom
ExceptionsBe>B, N>O
Influencing FactorsAtomic size, Z, Shielding, Electron configuration

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 negativeExothermic (favorable).
  • If energy is absorbed, ∆H is positiveEndothermic (unfavorable).

🔹 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


🔹 Successive Electron Gain Enthalpies

1st EGE: X(g) + e →X(g) + ΔH1​

2nd EGE: X(g) + e → X2−(g) + ΔH2

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

🧠 Factors Affecting Electron Gain Enthalpy

FactorEffectExplanation
Atomic Size↓ Size → ↑ EA (more negative)Smaller atoms attract added electron more strongly
Nuclear Charge (Z)↑ Z → ↑ EAStronger attraction of nucleus to incoming electron
Electronic ConfigurationStable config → ↓ EANoble gases, Be, N resist extra electrons
Shielding Effect↑ Shielding → ↓ EAWeakens the nuclear pull on added electron

📈 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 ElementsEA₁ 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 17EA₁ Trend
F < Cl < Br < I < At

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


Anomalies / Exceptions

Fluorine vs Chlorine

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

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


Be, Mg, N, Ne (Group 2, 15, 18)

ElementEGEReason
Be, MgPositive / near zeroFull s-orbital (2s², 3s²) – stable
NLow EAHalf-filled p-orbitals → Extra electron causes repulsion
Ne, Ar, KrPositiveFull octet, no tendency to accept electrons

🔎 Numerical Data (1st EA of Some Elements)

ElementConfigurationEA (kJ/mol)
Li1s² 2s¹–60
Be1s² 2s²~0
B1s² 2s² 2p¹–27
C1s² 2s² 2p²–122
N1s² 2s² 2p³–7
O1s² 2s² 2p⁴–141
F1s² 2s² 2p⁵–328
Cl[Ne] 3s² 3p⁵–349
Ar[Ne] 3s² 3p⁶~0

📊 Summary of Periodic Trends

TrendAcross a PeriodDown a Group
Electron Gain EnthalpyMore negativeLess negative
Most Negative EAGroup 17 (Cl)Decreases from Cl → I

🔥 JEE/NEET Focus Points

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

📝 Key Takeaways

PointDetail
EGE DefinitionEnergy change when 1 electron added to gaseous atom
Exothermic EAReleases energy → EA is negative
Most Negative EAChlorine (Cl)
Positive EANoble gases, Group 2, Nitrogen
2nd EAAlways positive due to repulsion
F vs ClCl has more negative EA due to lower e⁻–e⁻ repulsion
Trend Across PeriodEA becomes more negative
Trend Down GroupEA becomes less negative

🔁 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!


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.

🔹 Difference from Related Terms

PropertyApplies ToDescribesMeasured in
Electron Gain Enthalpy (EGE)Isolated atomEnergy released on adding an electronkJ/mol
ElectronegativityBonded atomPull of electrons in a bondNo unit (relative)

🔬 Pauling Scale of Electronegativity


👨‍🔬 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= (EAB​− (EAA​⋅EBB)1/2)1/2

Where:

  • χ = Electronegativity
  • EAB​​ = Bond energy of A–B bond
  • EAA, EBB​​​ = Bond energies of A–A and B–B

✅ Example Values:

ElementPauling EN
F4.0 (Highest)
O3.5
N3.0
Cl3.0
C2.5
H2.1
Na0.9
Cs0.7 (Lowest)

📈 Periodic Trends of Electronegativity


🔹 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)


🔹 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)


🔹 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)

🧠 Factors Affecting Electronegativity

FactorEffect on ENReason
Atomic Radius↓ radius → ↑ ENCloser nucleus, stronger pull
Nuclear Charge↑ charge → ↑ ENMore protons pull more strongly
Shielding Effect↑ shielding → ↓ ENInner electrons reduce effective pull
HybridizationMore s-character → ↑ ENs-orbitals are closer to nucleus

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


⚠️ Key Exceptions & Points

CaseExplanation
Noble gasesNot assigned EN in some scales (don’t usually form bonds)
HydrogenEN = 2.1, behaves variably based on bonding partner
Transition ElementsEN does not change uniformly across a period due to d-electron shielding
Lanthanide ContractionCauses higher-than-expected EN in later lanthanides

📊 Comparison of Common EN Scales

ScaleProposed byBasisMax Value
PaulingLinus PaulingBond energy4.0
MullikenRobert MullikenAvg of IE & EGEVaries
Allred-RochowAllred & RochowEffective nuclear charge & radiusVaries
SandersonR.T. SandersonStability ratioVaries

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


💡 Applications of Electronegativity

ApplicationDescription
Bond PolarityGreater difference in EN → more polar bond
Ionic vs Covalent CharacterLarge ΔEN → ionic bond; small ΔEN → covalent
Acidic/Basic NatureHigher EN → more acidic oxides (non-metals), Lower EN → basic oxides (metals)
Reactivity TrendsHalogens (high EN) → reactive nonmetals, Alkali metals (low EN) → reactive metals

🧪 Trick to Remember Electronegativity Order (for Period 2)

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


🧮 NEET/JEE Question Types

Question StyleExample
Order of ENArrange O, N, F in increasing EN
Predict bond characterWhich is more ionic: NaCl or MgCl₂?
Reason-basedWhy EN of N is less than O?
ConceptualWhy is EN not defined for noble gases?

📝 Quick Summary

ConceptTrend
Highest ENFluorine (F) = 4.0
Lowest ENCaesium (Cs) = 0.7
Across PeriodIncreases (L → R)
Down GroupDecreases
Depends onSize ↓, Z ↑, Shielding ↓
Used toPredict 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).


🔹 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

🔷 Types of Valency

TypeDefinitionExample
Electropositive ValencyNo. of electrons lostNa (Z=11): 1s²2s²2p⁶3s¹ → loses 1e⁻ → Valency = 1
Electronegative ValencyNo. of electrons gainedCl (Z=17): 1s²2s²2p⁶3s²3p⁵ → gains 1e⁻ → Valency = 1
CovalencyNo. of electrons sharedC (Z=6): shares 4e⁻ (2s²2p²) → Valency = 4

🔷 Valency and Group Number

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

🧠 Mnemonic (For p-block):

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


🔹 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

🔹 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

🔷 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 2LiBeBCNOFNe
Valency12343210

2️⃣ Down a Group (↓)

  • Valency remains constant
  • Because elements in a group have same number of valence electrons
  • Example:
Group 1LiNaKRbCs
Valency11111

🧠 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)

🔥 JEE & NEET Focused Questions on Valency

TypeExamples
Arrange elements by valency across a periodLi < Be < B < C…
Which element has variable valency?Fe, Cu, Cr
Noble gas with non-zero valency?Xe
Identify valency from configuration2s²2p³ → Valency = 3
Match group number to valencyGroup 16 → Valency = 2

📊 Quick Summary Table: Periodic Table Valency Trend

GroupValence ElectronsTypical ValencyType
1 (IA)11Electropositive
2 (IIA)22Electropositive
13 (IIIA)33Electropositive
14 (IVA)44Covalent
15 (VA)53Electronegative
16 (VIA)62Electronegative
17 (VIIA)71Electronegative
18 (VIIIA)80Inert

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.

9. Anomalous Trends & Exceptions in Periodic Properties

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


✳️ 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)

🔷 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

PropertyLithium (Li)Sodium (Na) & othersCause of Difference
Ionic RadiusSmallLargerSmall atomic size
Polarizing PowerHighLowerHigh charge density
ReactivityLess reactive with air/waterMore reactiveForms protective oxide layer
Nature of CompoundMore 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

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

🧠 Why Does Diagonal Relationship Occur?

Because of similarities in:

FactorExplanation
Ionic/Atomic RadiusNearly equal radii despite being in different groups
ElectronegativityClose values → Similar chemical bonding
Charge Density (Z/r)Similar polarizing power
Ionization EnthalpyClose values due to balance of size & nuclear charge

🧪 Detailed Examples

Li vs Mg

PropertyLiMg
Ionic Radius76 pm72 pm
Forms NitridesLi₃NMg₃N₂
Forms OxidesLi₂O (not peroxide)MgO
SolubilityLi₂CO₃ decomposes on heatingMgCO₃ also decomposes

Be vs Al

PropertyBeAl
Amphoteric natureBe(OH)₂Al(OH)₃
Covalent ChlorideBeCl₂AlCl₃
Tendency to form complex ions[BeF₄]²⁻[AlF₆]³⁻

B vs Si

PropertyBSi
SemiconductorYesYes
Acidic oxidesB₂O₃SiO₂
Forms hydridesB₂H₆SiH₄

🔷 3. Other Anomalous Trends in Periodic Properties


✅ 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)

✅ b) Electron Gain Enthalpy Exception

ElementEG EnthalpyReason
F < ClF is smaller → high electron repulsionFluorine’s small atomic size leads to tightly packed electrons, causing strong repulsion when an extra electron is added.
O < SSame reason: smaller O leads to less favorable electron gainSmaller atoms like F and O have high electron density, causing strong repulsion that reduces electron gain enthalpy.

✅ 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

✅ 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

🔷 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)|


📌 Summary Table: Anomalous Behaviors

AnomalyExampleReason
Li ≠ other alkali metalsLi forms Li₃NSmall size, high polarizing power
Be ≠ other alkaline earth metalsBeCl₂ covalentSmall size, high IE
Be ≈ AlAmphoteric natureDiagonal relationship
B ≈ SiCovalent oxidesDiagonal relationship
IE anomalyBe > B, N > OStable electronic configuration
Electron gain enthalpyCl > FF is too small, high repulsion
Lanthanide contractionZr ≈ HfPoor shielding by 4f

🧠 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

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

🧲 Reactivity and Periodicity

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

For MetalsReactivity ∝ Ease of losing electrons
For Non-MetalsReactivity ∝ Ease of gaining electrons

🔷 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)

🔄 Periodic Trends

DirectionMetallic CharacterNon-Metallic Character
Across a Period (→)DecreasesIncreases
Down a Group (↓)IncreasesDecreases

🔍 Reason:

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

🧪 Examples:

ElementMetallic Character
Na > Mg > AlMore metallic (Group 1 > 2 > 13)
P < S < ClIncreasing non-metallic character
Li < Na < K < Rb < CsIncreasing metallic character

🔷 2. Reactivity Trends in Groups and Periods

🔹 For Metals (like Group 1 and 2)

TrendReactivity
Down the groupIncreases
Across a periodDecreases

✅ Why?

  • Larger atoms → lower ionization energy → easily lose electrons

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


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

TrendReactivity
Down the groupDecreases
Across a periodIncreases

✅ Why?

  • Smaller atoms with high nuclear charge attract electrons better

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


🔷 3. Oxidizing and Reducing Nature


🔴 Oxidizing Agents

Substances that gain electrons themselves and oxidize others.

  • Typically non-metals (e.g. halogens)
  • Higher electron affinity = stronger oxidizing agent
ExampleReaction
F₂ + 2e⁻ → 2F⁻Fluorine is a strong oxidizing agent

🔺 Trend in Oxidizing Power (Group 17): F2​>Cl2​>Br2​>I2​


🔵 Reducing Agents

Substances that lose electrons themselves and reduce others.

  • Typically metals
  • Lower ionization energy = stronger reducing agent
ExampleReaction
Na → Na⁺ + e⁻Sodium is a strong reducing agent

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


🔄 Summary Table: Periodicity of Key Properties

PropertyAcross a Period (→)Down a Group (↓)
Atomic SizeDecreasesIncreases
Ionization EnthalpyIncreasesDecreases
ElectronegativityIncreasesDecreases
Metallic CharacterDecreasesIncreases
Non-Metallic CharacterIncreasesDecreases
Metal ReactivityDecreasesIncreases
Non-Metal ReactivityIncreasesDecreases
Oxidizing Power (Non-Metals)IncreasesDecreases
Reducing Power (Metals)DecreasesIncreases

📌 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

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.


🔷 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 BondElectronegativity DifferenceExampleWhere Found
IonicLarge difference (≥1.7)NaClMetal + Non-metal
CovalentSmall difference (<1.7)H₂O, Cl₂Non-metal + Non-metal
MetallicBetween metalsFe, CuBetween metals (free electrons)

📌 Use Periodic Table to:

  • Identify metals vs non-metals
  • Predict bonding type:
    • Left + RightIonic
    • Right + RightCovalent
    • Left + LeftMetallic

🔷 2. Nature of Oxides and Hydrides

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


🔹 Oxides:

Type of ElementNature of OxideExample
Metals (Groups 1, 2)BasicNa₂O, CaO
Metalloids (Group 13, 14)AmphotericAl₂O₃, SnO
Non-metals (Group 15–17)AcidicSO₂, P₂O₅

🔁 Trend:

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

🔹 Hydrides:

Type of ElementNature of HydrideExample
Group 1 & 2 metalsIonic (saline)NaH, CaH₂
p-Block non-metalsCovalentCH₄, NH₃, HCl
Transition metalsInterstitialTiH₂

🧠 Hydrides’ nature helps infer:

  • Bonding type
  • Reducing ability
  • Thermal stability

🔷 3. Predicting Oxidation States

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

📌 Trends in Oxidation States:

GroupGeneral 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 metalsVariable oxidation states (due to d-electrons)

🔁 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

🧠 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⁺

🔷 Summary Table

ApplicationWhat to PredictHow Periodic Table Helps
BondingIonic, covalent, metallicElectronegativity, position (metal/non-metal)
OxidesAcidic/basic/amphotericMetallic character → nature of oxide
HydridesIonic, covalent, interstitialGroup type + bonding nature
Oxidation StatesValency & possible chargesGroup number, inert pair effect, block type

🧪 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.

Thank You

🥰Learn Sufficient Notes🥰

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