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 Name | Subtopics Name |
---|---|---|
1 | Introduction | – Need for classification- Historical background |
2 | Dobereiner’s Triads | – Concept- Examples- Limitations |
3 | Newlands’ Law of Octaves | – Law of Octaves- Examples- Limitations |
4 | Mendeleev’s Periodic Table | – Basis of classification- Merits- Limitations |
5 | Modern Periodic Law & Table | – Modern Periodic Law- Structure of modern table- Groups, Periods, Blocks- Types of elements |
6 | Nomenclature (Z > 100) | – IUPAC rules- Naming conventions |
7 | Electronic Configuration & Periodic Table | – Relation with group & period- Valence electrons- Block classification |
8 | Periodic 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 |
9 | Anomalous Trends & Exceptions | – Irregularities- Diagonal relationship |
10 | Periodicity and Reactivity | – Metallic/non-metallic character- Reactivity trends- Oxidizing/reducing nature |
11 | Applications of Periodic Table | – Predicting bonding- Oxides/hydrides- Oxidation states |
12 | Question 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?
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. |
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:
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 | — |
✅ 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
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.
✅ 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.
✅ 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.

🔷 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
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 |
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
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.
🔷 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 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. |
🔷 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 |
🔷 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
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 |
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:
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) |
🔷 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. |
🔷 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.
🔴 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. |
🔷 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
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 |
🧠 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 |
🎯 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:
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 |
🔷 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.
🔷 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
🔷 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) |
🔷 6. TYPES OF ELEMENTS IN MODERN PERIODIC TABLE
Based on position & properties, elements are categorized as:
🟥 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 |
🟨 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 |
🟩 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 |
🟦 D. Noble Gases
Feature | Description |
---|---|
Location | Group 18 |
Properties | Stable, inert, monoatomic gases |
Examples | He, Ne, Ar, Kr, Xe, Rn |
🟪 E. Transition Elements
Feature | Description |
---|---|
Location | d-block (Groups 3–12) |
Properties | Variable oxidation states, colored compounds, form complexes |
🟫 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 |
🔷 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 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 |
🎯 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
Digit | Root |
---|---|
0 | nil |
1 | un |
2 | bi |
3 | tri |
4 | quad |
5 | pent |
6 | hex |
7 | sept |
8 | oct |
9 | enn |
🧩 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) |
🔷 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 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) |
🔷 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
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 |
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 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
📌 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 |
🔷 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
🔷 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
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 |
🔷 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 |
🔷 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 |
🎯 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
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.
🧠 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
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 |
💡 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
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 |
🔹 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
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 |
📈 Trends in Atomic Radius
🔸 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
🔸 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
🔄 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−
🔍 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
Direction | Atomic Radius |
---|---|
Across a period → | Decreases ⬇️ |
Down a group ⬇️ | Increases ⬆️ |
🔬 Visual Summary
Group↓ Period →
↑ →
| ↑ radius ↓ radius
↓
🧪 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? |
📌 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 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 |
🔹 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.
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 |
🔹 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
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 |
🧠 Quick Tricks to Remember
- Cation = compact (think: cat shrinks)
- Anion = airy, bigger
- In isoelectronic series, more protons → smaller ion
✅ 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 |
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:
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).
📈 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 |
🧠 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
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)
✅ 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)
🔷 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)
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 |
🎯 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 |
📝 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 |
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).
🔹 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
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 |
📈 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)
❗ 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.
✅ 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 |
🔎 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 |
📊 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 |
🔥 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 |
📝 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 |
🔁 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
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) |
🔬 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:
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) |
📈 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
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)
⚠️ 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 |
📊 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.
💡 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 |
🧪 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 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? |
📝 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).
🔹 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
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 |
🔷 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)
🔹 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 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 |
🧠 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
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 |
📊 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.
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
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
🔷 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) |
🧠 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 |
🧪 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 |
🔷 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
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. |
✅ 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
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 |
🧠 Tips for JEE/NEET:
- Memorize diagonal pairs (Li–Mg, Be–Al, B–Si)
- Focus on exceptions in IE, EA, EN
- Use concepts like shielding, penetration, and stability
- 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 Metals | Reactivity ∝ Ease of losing electrons |
---|---|
For Non-Metals | Reactivity ∝ 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
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
🧪 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 |
🔷 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)
🔹 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)
🔷 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
Example | Reaction |
---|---|
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
Example | Reaction |
---|---|
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
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 |
📌 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 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
🔷 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 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)
🔹 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
🔷 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) |
🔁 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
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 |
🧪 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