ChemistryNCERT Class 11

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Classification of Elements and Periodicity in Properties Notes

Study Notes

5 Topics19 Formulas24 PYQs36 Key Points

Chapter Overview Periodic Classification Modern Periodic Table Electronic Configuration s, p, d & f Block Elements Periodic Trends

Chapter at a Glance

Topics5

Key Points36

Formulas19

Diagrams13

NEET PYQs24

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Topics

Chapter Overview

Overview

This chapter explains how more than one hundred elements are arranged so that their properties can be predicted quickly. The journey begins with early attempts like Dobereiner’s triads, Newlands’ octaves and Mendeleev’s table, then reaches the modern periodic law based on atomic number. You learn how electronic configuration controls the position of an element, why elements are divided into s, p, d and f blocks, and how properties repeat periodically. NEET mainly tests trends in atomic size, ionic size, ionization enthalpy, electron gain enthalpy, electronegativity, valency, metallic character and oxidizing or reducing nature. The central idea is simple: periodic properties arise because valence-shell electronic configurations repeat at regular intervals.

Key Points6

1Periodic classification helps study many elements by grouping elements with similar properties.
2Atomic number is the fundamental basis of the modern periodic table because it fixes electronic configuration.
3The periodic table has 7 periods, 18 groups and four blocks based on the differentiating electron.
4Periodicity is mainly due to recurrence of similar valence-shell electronic configurations.
5Anomalies in trends usually arise from noble gas stability, half-filled or fully filled subshells, small size, high nuclear charge or poor shielding by d and f electrons.
6For NEET, trend questions often compare elements in the same period, same group or isoelectronic series.

Memory Tricks2

Trend Shortcut

FONClBrISCH: Fluorine is most electronegative; oxygen, nitrogen and chlorine often create exceptions in electron gain or ionization trends.

Periodic Table Direction Trick

Right and up generally means more non-metallic, more electronegative and higher ionization enthalpy; left and down generally means more metallic and larger size.

Examples2

Predicting an Unknown Element

If an element has valence configuration ns2np5, it belongs to Group 17, is a halogen, is highly electronegative and tends to form a -1 ion.

Everyday Periodicity

Sodium and potassium are both soft, reactive metals because both have one valence electron. Their similar outer configuration explains similar chemical behavior.

Reference Tables2

One-Page Chapter Map

Area	Main idea	High-yield NEET focus
Periodic Classification	Historical grouping of elements	Triads, octaves, Mendeleev merits and limitations
Modern Periodic Table	Arrangement by atomic number	Groups, periods, position and nomenclature
Electronic Configuration	Filling of electrons in orbitals	Aufbau, Pauli, Hund, valence shell, exceptions
Blocks	Classification by differentiating electron	s, p, d, f block features
Periodic Trends	Regular variation of properties	Radius, IE, electron gain enthalpy, electronegativity, valency

Major Trend Summary

Property	Across a period	Down a group	Reason
Atomic radius	Decreases	Increases	Zeff increases across; shells increase down
Ionization enthalpy	Generally increases	Generally decreases	Electron held more strongly across; farther from nucleus down
Electron gain enthalpy	Generally becomes more negative	Generally less negative	Small atoms attract added electron better; exceptions exist
Electronegativity	Increases	Decreases	Attraction for shared electron pair rises with Zeff
Metallic character	Decreases	Increases	Electron loss becomes harder across and easier down

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Common Mistakes2

Atomic mass versus atomic number

Do not use atomic mass to arrange the modern periodic table. Modern classification is based on atomic number because atomic number determines electronic configuration.

Ignoring exceptions

NEET often asks exceptions such as Be > B and N > O in first ionization enthalpy, and Cl having more negative electron gain enthalpy than F.

Formula Cards3

Effective Nuclear Charge

Approximate net positive charge felt by an electron after shielding by other electrons. Higher Zeff generally means smaller size and higher ionization enthalpy.

Variables

Zeff=

effective nuclear charge experienced by an electron

Z=

actual atomic number or nuclear charge

σ=

shielding or screening constant

Moseley Relation

Moseley showed that X-ray frequency depends on atomic number, supporting atomic number as the basis of classification.

Variables

ν=

frequency of characteristic X-rays

Z=

atomic number

a, b=

constants for a given spectral series

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Diagrams3

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Periodic Classification

Overview

Periodic classification was needed because the number of known elements increased rapidly and memorizing each element separately became impossible. Scientists tried to arrange elements so that similarities in physical and chemical properties became visible. Dobereiner grouped elements into triads, where the atomic mass of the middle element was approximately the average of the other two. Newlands arranged elements by increasing atomic mass and observed repetition after every eighth element, called the law of octaves. Mendeleev gave the first highly successful periodic table based mainly on atomic mass and chemical properties. He left gaps for undiscovered elements and predicted their properties, but his table could not properly explain isotopes, anomalous pairs and the position of hydrogen.

Key Points6

1The need for classification arose from the increasing number of elements and their compounds.
2Early classifications used atomic mass because atomic number had not yet been established.
3Dobereiner’s triads showed that atomic masses and chemical properties were related.
4Newlands introduced the idea of periodic repetition of properties.
5Mendeleev’s greatest success was predicting undiscovered elements and correcting atomic masses.
6The failure of atomic mass as a fundamental basis led to the modern periodic table based on atomic number.

Memory Tricks2

Order of Early Classification

D-N-M-M: Dobereiner, Newlands, Mendeleev, Moseley. Think 'Do Not Memorize Mass' because the final correct basis became atomic number.

Mendeleev’s Eka Names

Eka means one place below. Eka-aluminium became gallium, eka-silicon became germanium, and eka-boron became scandium.

Examples2

Dobereiner Triad Example

For chlorine, bromine and iodine: atomic masses are about 35.5, 79.9 and 126.9. The average of chlorine and iodine is close to bromine.

Mendeleev Prediction Example

Mendeleev predicted eka-silicon with properties close to germanium before germanium was discovered, showing the predictive power of his table.

Reference Tables2

Timeline of Classification

Scientist	Basis	Main idea	Limitation
Dobereiner	Atomic mass and properties	Triads of three similar elements	Only a few triads were known
Newlands	Increasing atomic mass	Law of octaves	Failed beyond calcium and did not suit transition elements
Mendeleev	Atomic mass plus chemical properties	Periodic table with gaps	Could not explain isotopes and anomalous pairs
Moseley	Atomic number	Modern periodic law foundation	Required knowledge of subatomic structure

Merits and Limitations of Mendeleev’s Periodic Table

Merits	Limitations
Left gaps for undiscovered elements	No fixed position for hydrogen
Predicted properties of eka-Al, eka-Si and eka-B	Isotopes could not be accommodated
Corrected atomic masses of some elements	Anomalous pairs violated increasing atomic mass order
Grouped elements with similar formulae of oxides and hydrides	Rare earth elements were not properly placed

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Common Mistakes2

Saying Mendeleev ignored properties

Mendeleev arranged mostly by atomic mass but sometimes reversed order to keep elements with similar chemical properties together.

Applying Newlands to all elements

Newlands’ law of octaves failed for heavier elements and transition elements; it is not a universal periodic law.

Formula Cards2

Dobereiner Triad Average

In a valid triad, the middle element has properties intermediate between the other two, and its atomic mass is nearly the average of the remaining two.

Variables

mass of first=

atomic mass of the lightest element in the triad

mass of third=

atomic mass of the heaviest element in the triad

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Diagrams3

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Modern Periodic Table

Overview

The modern periodic table is based on the modern periodic law: physical and chemical properties of elements are periodic functions of their atomic numbers. Since atomic number determines electronic configuration, the table becomes a systematic map of electron arrangement and properties. The long form periodic table has 7 periods and 18 groups. Periods are horizontal rows and generally indicate the highest occupied shell, while groups are vertical columns containing elements with similar valence-shell configurations. Elements are also positioned into s, p, d and f blocks. For elements with atomic number above 100, IUPAC temporary nomenclature uses numerical roots such as nil, un, bi and tri. NEET questions commonly ask group, period, block and valence configuration from atomic number.

Key Points6

1Atomic number removes the isotope problem because isotopes have the same atomic number and occupy the same position.
2The period number equals the highest principal quantum number present in electronic configuration.
3Group number for s-block equals the number of valence electrons; for p-block, group number equals 10 plus valence electrons.
4d-block elements occupy Groups 3-12 and are placed between s- and p-blocks.
5f-block elements are shown separately to keep the table compact, but they belong to periods 6 and 7.
6Hydrogen has a special position because it resembles both alkali metals and halogens in different ways.

Memory Tricks2

p-Block Group Trick

For p-block, count ns + np valence electrons and add 10. Example ns2np5 has 7 valence electrons, so group = 17.

IUPAC Digits Trick

0-9 roots: nil, un, bi, tri, quad, pent, hex, sept, oct, enn. Remember 'Naughty Unicorns Bite Triangular Quad Penta Hexa Sept Oct Envelopes'.

Examples2

Position of Magnesium

Mg has Z = 12 and configuration 1s2 2s2 2p6 3s2. Highest n is 3, so period 3; ns2 means Group 2; last electron enters s, so s-block.

Temporary Name Example

Element 104 was temporarily named un-nil-quad-ium from digits 1, 0 and 4 before receiving its permanent name rutherfordium.

Reference Tables2

Periods in the Long Form Periodic Table

Period	Shell filled	Number of elements	Ends with
1	K shell	2	He
2	L shell	8	Ne
3	M shell	8	Ar
4	N shell	18	Kr
5	O shell	18	Xe
6	P shell	32	Rn
7	Q shell	Incomplete/32 possible	Og

IUPAC Temporary Nomenclature Roots

Digit	Root	Example use
0	nil	110: un-un-nil-ium
1	un	101: un-nil-un-ium
2	bi	102: un-nil-bi-ium
3	tri	103: un-nil-tri-ium
4	quad	104: un-nil-quad-ium
5	pent	105: un-nil-pent-ium
6	hex	106: un-nil-hex-ium
7	sept	107: un-nil-sept-ium
8	oct	108: un-nil-oct-ium
9	enn	109: un-nil-enn-ium

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Common Mistakes2

Confusing period and group

Period is horizontal and linked to highest n. Group is vertical and linked to valence-shell similarity.

Placing helium in p-block

Helium is placed in Group 18 due to noble gas properties, but its configuration is 1s2, so it is an s-block element by configuration.

Formula Cards3

Period Number from Configuration

The largest principal quantum number present tells the period of the element.

Variables

n=

principal quantum number of the outermost occupied shell

p-Block Group Number

For Groups 13 to 18, group number can be calculated from valence electrons in ns and np orbitals.

Variables

valence electrons=

electrons present in ns and np subshells of the valence shell

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Diagrams3

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Electronic Configuration

Overview

Electronic configuration describes how electrons are distributed among shells, subshells and orbitals. It is the bridge between atomic structure and the periodic table because the position and properties of an element depend mainly on its valence-shell configuration. Electrons fill orbitals according to the Aufbau principle, which follows increasing energy based on the n + l rule. Pauli’s exclusion principle states that an orbital can hold a maximum of two electrons with opposite spins. Hund’s rule says that electrons occupy degenerate orbitals singly with parallel spins before pairing. Valence electrons decide group, block, valency, bonding behavior and chemical reactivity. Exceptions such as chromium and copper occur because half-filled and fully filled subshells give extra stability.

Key Points6

1Electronic configuration explains why properties repeat periodically.
2The filling order begins 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s.
3s subshell has 1 orbital, p has 3, d has 5 and f has 7 orbitals.
4Maximum electrons: s = 2, p = 6, d = 10, f = 14.
5Valence electrons are easiest to use for predicting group, valency and bonding.
6Half-filled and fully filled subshells are unusually stable due to symmetry and exchange energy.

Memory Tricks2

Filling Order Mnemonic

Remember: 1s 2s 2p 3s 3p 4s 3d 4p 5s. Say 'S S P S P S D P S' while drawing diagonal arrows.

Pauli and Hund Shortcut

Pauli means 'Pair opposite'; Hund means 'Half-fill before pairing'.

Examples2

Electronic Configuration of Chlorine

Cl has Z = 17: 1s2 2s2 2p6 3s2 3p5. It has 7 valence electrons and belongs to Group 17.

Why Chromium is Exceptional

Chromium becomes [Ar] 3d5 4s1 instead of [Ar] 3d4 4s2 because a half-filled d subshell gives extra stability.

Reference Tables2

Subshell Capacity and Orbitals

Subshell	l value	Number of orbitals	Maximum electrons
s	0	1	2
p	1	3	6
d	2	5	10
f	3	7	14

Common Electronic Configuration Exceptions

Element	Expected configuration	Actual configuration	Reason
Cr, Z=24	[Ar] 3d4 4s2	[Ar] 3d5 4s1	half-filled d subshell stability
Cu, Z=29	[Ar] 3d9 4s2	[Ar] 3d10 4s1	fully filled d subshell stability
Mo, Z=42	[Kr] 4d4 5s2	[Kr] 4d5 5s1	half-filled d subshell stability
Ag, Z=47	[Kr] 4d9 5s2	[Kr] 4d10 5s1	fully filled d subshell stability

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Common Mistakes2

Filling 3d before 4s in neutral atoms

In neutral atoms, 4s fills before 3d because 4s is lower in energy during filling. However, during ion formation from transition metals, 4s electrons are removed first.

Pairing too early

For p, d and f orbitals, do not pair electrons until all degenerate orbitals have one electron each.

Formula Cards3

Aufbau n + l Rule

Orbitals with smaller n + l value fill first; if two orbitals have the same n + l, the one with smaller n fills first.

Variables

n=

principal quantum number

l=

azimuthal quantum number of subshell: s=0, p=1, d=2, f=3

Maximum Electrons in Subshell

A subshell with azimuthal quantum number l contains 2l + 1 orbitals, each holding two electrons.

Variables

l=

azimuthal quantum number

2l + 1=

number of orbitals in the subshell

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Diagrams3

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s, p, d & f Block Elements

Overview

Elements are classified into s, p, d and f blocks according to the subshell into which the differentiating electron enters. This block classification connects electronic configuration with chemical behavior. s-block elements have their last electron in an s orbital and include Groups 1 and 2, except helium’s special placement. p-block elements have their last electron in a p orbital and include Groups 13 to 18; they contain metals, non-metals and metalloids. d-block elements are transition elements where the penultimate d subshell is being filled, commonly showing variable oxidation states and colored compounds. f-block elements are inner transition elements involving 4f and 5f filling, including lanthanoids and actinoids. NEET often asks block, general configuration and characteristic trends.

Key Points6

1Blocks reflect the type of valence or differentiating orbital, not simply table color or position.
2s-block elements lose electrons easily and form mainly ionic compounds.
3p-block elements show greater variation in bonding, oxidation states and metallic character.
4d-block elements are called transition elements because they lie between s- and p-blocks and have partially filled d subshells in atoms or ions.
5f-block elements are placed separately to avoid an excessively wide periodic table.
6The block arrangement explains broad similarities in physical and chemical properties.

Memory Tricks2

Block General Configuration Trick

s is simple ns; p is ns np; d is one shell behind; f is two shells behind.

Location Trick

s is left, p is right, d is middle, f is footnote. This layout itself helps remember the blocks.

Examples2

Identify the Block

Element with configuration [Ne] 3s2 3p3 has last electron in p subshell, so it is a p-block element.

d-Block Behavior

Iron forms Fe2+ and Fe3+ because d-block elements can lose ns and some (n-1)d electrons, leading to variable oxidation states.

Reference Tables2

Block Comparison

Block	Groups/Series	General configuration	Common nature
s-block	Groups 1-2	ns1-2	Reactive electropositive metals
p-block	Groups 13-18	ns2np1-6	Metals, metalloids, non-metals, noble gases
d-block	Groups 3-12	(n-1)d1-10ns0-2	Transition metals
f-block	Lanthanoids and actinoids	(n-2)f1-14(n-1)d0-1ns2	Inner transition elements

General Characteristics of Each Block

Property	s-block	p-block	d-block	f-block
Bonding	Mostly ionic	Ionic and covalent	Metallic, ionic, covalent, complex	Mostly metallic and ionic
Oxidation states	Usually fixed +1 or +2	Variable, includes negative states	Variable	Commonly +3, actinoids more variable
Metallic character	Very high	Decreases across period	High	High
Compounds	Often colorless salts	Wide variety	Often colored	Often colored/paramagnetic

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Common Mistakes2

Calling all d-block elements transition elements without condition

A transition element should have partially filled d subshell in atom or common ion. Zn, Cd and Hg are d-block but not typical transition elements.

Forgetting helium’s configuration

Helium is placed with noble gases because of chemical inertness, but by electronic configuration it belongs to s-block.

Formula Cards4

s-Block General Configuration

The last electron enters the outermost s subshell. Group 1 is ns1 and Group 2 is ns2.

Variables

n=

principal quantum number of the valence shell

p-Block General Configuration

The last electron enters a p subshell; p-block includes Groups 13 to 18.

Variables

n=

principal quantum number of valence shell

np=

p subshell of the outermost shell

d-Block General Configuration

The differentiating electron enters the penultimate d subshell; many d-block elements are transition elements.

Variables

n=

outermost shell number

(n - 1)d=

d subshell of the penultimate shell

f-Block General Configuration

The differentiating electron enters an inner f subshell, giving lanthanoids and actinoids.

Variables

(n - 2)f=

f subshell two shells inside the valence shell

ns2=

outer s subshell generally containing two electrons

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Diagrams3

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Periodic Trends

10m

Overview

Periodic trends are regular variations in properties caused by changing effective nuclear charge, number of shells and shielding. Across a period, electrons enter the same shell while nuclear charge increases, so effective nuclear charge increases and atomic radius generally decreases. Down a group, new shells are added and shielding increases, so size increases. Ionization enthalpy generally increases across a period and decreases down a group, with important exceptions such as Be-B and N-O. Electron gain enthalpy generally becomes more negative across a period, but chlorine is more negative than fluorine due to lower electron-electron repulsion. Electronegativity increases across and decreases down. Valency, metallic character, non-metallic character and oxidizing or reducing nature follow from electron loss or gain tendencies.

Key Points6

1Atomic radius includes covalent radius, metallic radius and van der Waals radius depending on bonding situation.
2Ionic radius depends on charge and electron count: more positive charge means smaller size, more negative charge means larger size.
3Ionization enthalpy is affected by size, nuclear charge, shielding, penetration and stable electronic configurations.
4Electron gain enthalpy is more favorable for atoms that can achieve stable configuration by gaining an electron.
5Valency in representative elements often increases from 1 to 4 and then decreases from 4 to 0 across a period.
6Non-metals are better oxidizing agents because they gain electrons; metals are better reducing agents because they lose electrons.

Memory Tricks3

Cation-Anion Size Trick

CATions are PAWsitive and PAW smaller; ANions are negative and expand. Same electrons? More protons means smaller ion.

Ionization Exception Trick

Be-B and N-O are the classic first IE dips. Remember: 'Be stable s2, N stable p3'.

Electron Gain Trick

For halogens, remember 'Cl beats F' in electron gain enthalpy because fluorine is too small and crowded.

Examples3

Comparing Atomic Radius

Na is larger than Mg because both are in period 3, but Mg has higher nuclear charge pulling electrons closer.

Comparing Reducing Nature

Potassium is a stronger reducing agent than sodium because it loses its valence electron more easily due to larger size and lower ionization enthalpy.

Comparing Electronegativity

Fluorine is more electronegative than chlorine because fluorine is smaller and attracts the shared electron pair more strongly.

Reference Tables2

Periodic Trend Comparison

Property	Across period	Down group	Important exceptions
Atomic radius	Decreases	Increases	Noble gas radii are van der Waals radii and appear large
Ionic radius	Cations shrink, anions expand; isoelectronic size falls with Z	Usually increases	O2− > F− > Na+ > Mg2+ > Al3+
Ionization enthalpy	Generally increases	Generally decreases	Be > B, N > O, Mg > Al, P > S
Electron gain enthalpy	Generally more negative	Generally less negative	Cl more negative than F; noble gases unfavorable
Electronegativity	Increases	Decreases	F highest
Metallic character	Decreases	Increases	Cs/Fr region most metallic

High-Yield NEET Exceptions

Trend	Expected	Observed	Reason
First IE: Be and B	B higher than Be	Be > B	Removing 2p electron from B is easier than removing 2s electron from Be
First IE: N and O	O higher than N	N > O	N has stable half-filled 2p3; O has paired electron repulsion
Electron gain: F and Cl	F more negative	Cl more negative than F	F is very small; added electron faces greater repulsion
Atomic radius: noble gases	Should be small	Reported radii are large	Usually van der Waals radii, not covalent radii

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Common Mistakes3

Confusing electron gain enthalpy and electronegativity

Electron gain enthalpy is for adding an electron to an isolated gaseous atom. Electronegativity is attraction for a shared pair in a bond.

Using one radius type for all elements

Atomic radius may mean covalent, metallic or van der Waals radius depending on the element and bonding situation. Noble gases are often compared using van der Waals radii.

Forgetting isoelectronic rule

In O2−, F−, Na+, Mg2+ and Al3+, all have 10 electrons. Size decreases as nuclear charge increases: O2− > F− > Na+ > Mg2+ > Al3+.

Formula Cards4

Effective Nuclear Charge

Higher effective nuclear charge pulls electrons closer, decreasing radius and increasing ionization enthalpy.

Variables

Zeff=

net nuclear charge felt by an electron

Z=

atomic number

σ=

shielding constant

First Ionization Enthalpy

Energy required to remove the most loosely bound electron from one mole of isolated gaseous atoms.

Variables

X(g)=

isolated gaseous atom

ΔiH1=

first ionization enthalpy

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Diagrams4

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Formula Sheet

Effective Nuclear Charge

Approximate net positive charge felt by an electron after shielding by other electrons. Higher Zeff generally means smaller size and higher ionization enthalpy.

Variables

Zeff=

effective nuclear charge experienced by an electron

Z=

actual atomic number or nuclear charge

σ=

shielding or screening constant

Moseley Relation

Moseley showed that X-ray frequency depends on atomic number, supporting atomic number as the basis of classification.

Variables

ν=

frequency of characteristic X-rays

Z=

atomic number

a, b=

constants for a given spectral series

Maximum Electrons in a Shell

Gives the maximum number of electrons that can be accommodated in a shell with principal quantum number n.

Variables

n=

principal quantum number of the shell

Dobereiner Triad Average

In a valid triad, the middle element has properties intermediate between the other two, and its atomic mass is nearly the average of the remaining two.

Variables

mass of first=

atomic mass of the lightest element in the triad

mass of third=

atomic mass of the heaviest element in the triad

Newlands Octave Relation

When elements were arranged by increasing atomic mass, Newlands observed that every eighth element showed similar properties, like musical octaves.

Variables

1st element=

an element in Newlands' arrangement

8th element=

the element appearing after seven positions

Period Number from Configuration

The largest principal quantum number present tells the period of the element.

Variables

n=

principal quantum number of the outermost occupied shell

p-Block Group Number

For Groups 13 to 18, group number can be calculated from valence electrons in ns and np orbitals.

Variables

valence electrons=

electrons present in ns and np subshells of the valence shell

IUPAC Temporary Name Digit Roots

Used to name elements with atomic number greater than 100 before permanent names are assigned.

Variables

digit=

each digit of the atomic number written as a root

Aufbau n + l Rule

Orbitals with smaller n + l value fill first; if two orbitals have the same n + l, the one with smaller n fills first.

Variables

n=

principal quantum number

l=

azimuthal quantum number of subshell: s=0, p=1, d=2, f=3

Maximum Electrons in Subshell

A subshell with azimuthal quantum number l contains 2l + 1 orbitals, each holding two electrons.

Variables

l=

azimuthal quantum number

2l + 1=

number of orbitals in the subshell

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Modern periodic law: properties are periodic functions of atomic number, not atomic mass.

Period number equals the highest principal quantum number of the valence shell.

Group number is strongly linked with valence electrons, especially for s- and p-block elements.

s-block: Groups 1-2; p-block: Groups 13-18; d-block: Groups 3-12; f-block: inner transition elements.

Across a period, effective nuclear charge generally increases and atomic radius decreases.

Down a group, shells increase, so atomic and ionic radii generally increase.

Ionization enthalpy and electronegativity generally increase across a period and decrease down a group.

Metallic character increases down a group and decreases across a period.

Periodic classification helps study many elements by grouping elements with similar properties.

Atomic number is the fundamental basis of the modern periodic table because it fixes electronic configuration.

Classification reduces memorization by grouping elements with similar properties.

Dobereiner’s triads worked only for a few groups such as Li, Na, K and Cl, Br, I.

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NEET PYQs — Classification of Elements and Periodicity in Properties

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Showing 3 of 24 questions. Sign up to practice all with answers, explanations, and AI help.

NEET 2026Set 11MediumQ1

Identify the incorrect statement from the following:

NEET 2026Set 11EasyQ2

The correct order of increasing metallic character of Na, Be, P, Mg and Si is:

NEET 2025Set 45MediumQ3

Which of the following statements are true? A. Unlike Ga that has a very high melting point, Cs has a very low melting point. B. On Pauling scale, the electronegativity values of N and Cl are not the same. C. $\mathrm{Ar}$, $\mathrm{K^+}$, $\mathrm{Cl^-}$, $\mathrm{Ca^{2+}}$, and $\mathrm{S^{2-}}$ are all isoelectronic species. D. The correct order of the first ionization enthalpies of Na, Mg, Al, and Si is: $$\mathrm{Si > Al > Mg > Na}$$ E. The atomic radius of Cs is greater than that of Li and Rb. Choose the correct answer from the options given below: