This chapter explains how the modern periodic law arranges all elements by atomic number into 7 periods, 18 groups and the s, p, d and f blocks, and how an element’s electronic configuration fixes its exact position. It then tracks the periodic trends that flow from this layout, namely atomic and ionic radii, ionization enthalpy, electron gain enthalpy, electronegativity, valence and the acidic or basic nature of oxides. For NEET it is a high-yield, almost fully fact-and-trend based chapter where most questions test the direction and the common exceptions of these periodic properties.
Table of Contents
- Key Concepts - Genesis of classification, periodic laws, modern table, blocks, periodic trends
- Weightage in Board & Entrance Exams
- Important Definitions
- Solved Examples
- Important Questions for Board Exams
- Quick Revision Points
Key Concepts
1. Genesis of Periodic Classification
Long before the modern table, chemists kept trying to group elements with similar properties so the subject would be easier to learn. Each attempt got a little closer to the truth.
- Döbereiner’s Triads (1829): elements arranged in groups of three where the atomic mass of the middle element was roughly the average of the other two (e.g. Li, Na, K). Failed because very few such triads existed.
- Newlands’ Law of Octaves (1865): when elements were arranged by increasing atomic mass, every eighth element had properties similar to the first - like the eighth note of a musical octave. Worked only up to calcium.
- Lothar Meyer (1869): plotted atomic volume against atomic mass and obtained a periodic curve, showing properties repeat at intervals.
2. Mendeleev’s Periodic Law
Dmitri Mendeleev (1869) gave the first widely accepted table. His periodic law states: the physical and chemical properties of elements are a periodic function of their atomic masses.
He arranged 63 known elements in order of increasing atomic mass into horizontal periods and vertical groups of similar elements.
Merits of Mendeleev’s Table
- He left gaps for undiscovered elements and predicted their properties (eka-aluminium = gallium, eka-silicon = germanium) with remarkable accuracy.
- He placed elements by property even when it broke the mass order (e.g. he predicted corrected atomic masses).
Defects of Mendeleev’s Table
- Anomalous pairs: some elements with higher atomic mass came before lower ones (Ar before K; Co before Ni).
- Position of isotopes could not be explained (same element, different masses).
- Hydrogen had no fixed position (resembles both Group 1 and Group 17).
3. Modern Periodic Law and Present Form of the Table
Henry Moseley (1913) showed from X-ray studies that atomic number (Z), not atomic mass, is the fundamental property of an element. This corrected all of Mendeleev’s anomalies.
Modern Periodic Law: the physical and chemical properties of elements are a periodic function of their atomic numbers.
Features of the Long Form Table
- 7 periods (horizontal rows) = number of principal energy shells; period number = value of n for the outermost shell.
- 18 groups (vertical columns) = elements with the same outer electronic configuration and similar properties.
- The number of elements in each period (2, 8, 8, 18, 18, 32, 32) is fixed by how many electrons fill the available subshells (2n² rule).
4. Electronic Configuration and Types of Elements (s, p, d, f Blocks)
The table is divided into four blocks based on the subshell into which the last (differentiating) electron enters.
| Block | Groups | Outer Configuration | Examples |
|---|---|---|---|
| s-block | 1 & 2 | ns¹⁻² | Alkali & alkaline earth metals (Na, Ca) |
| p-block | 13–18 | ns² np¹⁻⁶ | B, C, N, O, halogens, noble gases |
| d-block | 3–12 | (n−1)d¹⁻¹⁰ ns⁰⁻² | Transition metals (Fe, Cu, Zn) |
| f-block | placed below | (n−2)f¹⁻¹⁴ | Lanthanoids & actinoids |
- s- and p-block together form the representative (normal) elements.
- d-block = transition elements; f-block = inner transition elements.
- Group 18 = noble gases (ns² np⁶, fully filled, very stable).
Tip to locate an element: the period = highest value of n; for s/p-block the group = (outer electrons) or (10 + outer electrons) for p-block; for d-block group = (n−1)d + ns electrons.
5. Atomic and Ionic Radius
Atomic radius is the distance from the centre of the nucleus to the outermost shell of electrons. We use covalent, van der Waals, or metallic radii depending on the bonding.
- Across a period (left → right): atomic radius decreases, because nuclear charge increases while electrons enter the same shell, pulling them inward.
- Down a group (top → bottom): atomic radius increases, because a new shell is added in each period.
Ionic Radius
- A cation is smaller than its parent atom (it loses a shell / has higher effective nuclear charge per electron).
- An anion is larger than its parent atom (added electrons increase repulsion).
- Isoelectronic species (same number of electrons): radius decreases as nuclear charge increases. Order: N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺.
6. Ionization Enthalpy (IE)
Ionization enthalpy is the minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.
M(g) + IE₁ → M⁺(g) + e⁻
- Across a period: IE generally increases (nuclear charge rises, size falls).
- Down a group: IE decreases (size increases, shielding increases).
- Successive ionization enthalpies always increase: IE₁ < IE₂ < IE₃.
Exceptions (very important for exams): Be has higher IE than B (stable fully filled 2s²), and N has higher IE than O (stable half-filled 2p³ configuration).
7. Electron Gain Enthalpy (ΔₑgH)
Electron gain enthalpy is the energy change when an electron is added to an isolated gaseous atom to form an anion.
X(g) + e⁻ → X⁻(g) + ΔₑgH
- More negative ΔₑgH means the atom releases more energy and accepts the electron more readily.
- Across a period: becomes more negative (atoms get smaller, more eager for electrons).
- Down a group: becomes less negative (size increases).
- Halogens have the most negative electron gain enthalpies.
Exception: ΔₑgH of chlorine is more negative than fluorine, because F is so small that incoming electrons face high repulsion in its compact 2p shell.
8. Electronegativity
Electronegativity is the tendency of an atom in a molecule to attract the shared pair of electrons towards itself. Unlike IE and ΔₑgH, it is a relative property with no units.
- Across a period: electronegativity increases (size decreases, nuclear pull increases).
- Down a group: electronegativity decreases.
- Fluorine is the most electronegative element (Pauling value 4.0).
Electronegativity decides bond polarity, the metallic/non-metallic character, and the acidic/basic nature of oxides.
9. Valence (Valency)
Valence is the combining capacity of an element, decided by the number of valence electrons.
- Across a period: valence towards hydrogen rises 1→4 then falls 4→1 (CH₄, NH₃, H₂O, HF).
- Down a group: valence usually stays constant (same outer configuration).
- For oxides, the highest valence often equals the group number for representative elements.
10. Periodic Trends in Chemical Properties
The physical trends above translate directly into chemical behaviour.
- Metallic character increases down a group and decreases across a period (metals lose electrons easily - low IE).
- Non-metallic character increases across a period and decreases down a group.
- Nature of oxides: metallic oxides are basic (Na₂O), non-metallic oxides are acidic (SO₃, Cl₂O₇), and oxides of borderline elements are amphoteric (Al₂O₃, ZnO).
- Reactivity of metals increases down a group; reactivity of non-metals decreases down a group.
[DIAGRAM: A periodic table with arrows - atomic radius and metallic character increasing down/left; ionization enthalpy, electronegativity and non-metallic character increasing up/right towards fluorine.]
11. Diagonal Relationship
Some elements of period 2 resemble the period-3 element placed diagonally to the right (Li–Mg, Be–Al, B–Si). This happens because the diagonal pair has similar charge/size ratios and similar electronegativities.
12. Nomenclature of Elements with Z > 100
To avoid naming disputes for newly synthesised superheavy elements, IUPAC gives a temporary systematic name based on the digits of the atomic number, using numerical roots, with the suffix -ium.
| Digit | 0 | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9 |
|---|---|---|---|---|---|---|---|---|---|---|
| Root | nil | un | bi | tri | quad | pent | hex | sept | oct | enn |
- Z = 101 → Unnilunium (Unu), Md
- Z = 104 → Unnilquadium (Unq), Rf
- Z = 120 → Unbinilium (Ubn)
Rule: write the roots for each digit in order, add -ium, and the symbol is the first letter of each root. If “i” of “bi”/”tri” is followed by “ium”, one “i” is dropped; “enn” + “nil” drops one “n”.
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Weightage in Board & Entrance Exams
| Exam | Typical Weightage | Most-Tested Areas |
|---|---|---|
| CBSE Board (Class 11) | 6–8 marks | Periodic trends, IE exceptions, electronic configuration, blocks |
| JEE Main / Advanced | 1–2 questions | Isoelectronic radii, IE order, electronegativity, nomenclature |
| NEET | 2–3 questions | Atomic/ionic radius, IE & ΔₑgH exceptions, metallic character |
[TABLE: Question-type split - VSA (1 mark): definitions, blocks, most electronegative element; SA (2–3 marks): trend reasoning, IE exceptions; LA (5 marks): compare across period/group, predict properties, nomenclature.]
Important Definitions
| Term | Definition |
|---|---|
| Modern periodic law | Properties of elements are a periodic function of their atomic numbers |
| Period | Horizontal row; equals the number of the outermost principal shell (n) |
| Group | Vertical column of elements with the same outer electronic configuration |
| Atomic radius | Distance from the nucleus to the outermost electron shell |
| Isoelectronic species | Atoms/ions with the same number of electrons (e.g. O²⁻, F⁻, Na⁺) |
| Ionization enthalpy | Energy to remove the outermost electron from a gaseous atom |
| Electron gain enthalpy | Energy change when an electron is added to a gaseous atom |
| Electronegativity | Tendency of an atom to attract a shared electron pair in a bond |
| Representative elements | s- and p-block elements (Groups 1, 2 and 13–18) |
| Diagonal relationship | Similarity between an element and the one diagonally right below it |
Solved Examples
Example 1
An element has the electronic configuration 1s² 2s² 2p⁶ 3s² 3p³. Identify its block, group, and period.
Answer: Last electron enters the 3p subshell → p-block. Highest n = 3 → Period 3. Outer electrons = 2 + 3 = 5 → group = 10 + 5 = Group 15. The element is phosphorus (Z = 15).
Example 2
Arrange the isoelectronic species N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺ in increasing order of ionic radius.
Answer: All have 10 electrons; radius decreases as nuclear charge rises. Increasing order: Mg²⁺ < Na⁺ < F⁻ < O²⁻ < N³⁻.
Example 3
Why is the first ionization enthalpy of nitrogen greater than that of oxygen?
Answer: Nitrogen has a stable, exactly half-filled 2p³ configuration which is hard to disturb, so removing an electron needs more energy. Oxygen (2p⁴) loses one electron to reach the stable half-filled state, so its IE₁ is lower.
Example 4
The electron gain enthalpy of chlorine is more negative than that of fluorine. Explain.
Answer: Fluorine is very small, so its compact 2p shell causes strong electron–electron repulsion when an extra electron is added, reducing the energy released. Chlorine is larger (3p shell), repulsion is less, so its ΔₑgH is more negative.
Example 5
Which of Be and B has the higher first ionization enthalpy, and why?
Answer: Beryllium. Be has a stable fully filled 2s² configuration, so its outer electron is tightly held. Boron’s single 2p electron is at higher energy and easier to remove, so B has a lower IE₁.
Example 6
Write the IUPAC systematic name and symbol for the element with atomic number 105.
Answer: Digits 1-0-5 → un-nil-pent + ium = Unnilpentium, symbol Unp (now named Dubnium, Db).
Important Questions for Board Exams
1-Mark Questions (VSA)
- State the modern periodic law.
- Name the most electronegative element in the periodic table.
- Which block do transition elements belong to?
- Why is the size of a cation smaller than its parent atom?
- Write the IUPAC name of the element with Z = 120.
2–3-Mark Questions (SA)
- Why does atomic radius decrease across a period but increase down a group? Explain.
- Explain why the first ionization enthalpy of N is greater than that of O.
- Arrange O²⁻, F⁻, Na⁺, Mg²⁺ in increasing order of size and justify.
- Distinguish between ionization enthalpy and electron gain enthalpy with one example each.
5-Mark Questions (LA)
- Discuss the variation of atomic radius, ionization enthalpy, electron gain enthalpy, and electronegativity across a period and down a group, with reasons.
- What are the defects of Mendeleev’s periodic table, and how does the modern periodic law remove them?
- Explain the trend in metallic and non-metallic character and the acidic/basic nature of oxides across period 3.
Quick Revision Points
- Döbereiner triads → Newlands octaves → Mendeleev (atomic mass) → Modern law (atomic number)
- Modern periodic law: properties are a periodic function of atomic number Z
- Long form: 7 periods, 18 groups; period = outermost shell n
- Blocks: s (Gp 1–2), p (Gp 13–18), d (Gp 3–12, transition), f (inner transition)
- Atomic radius: decreases across a period, increases down a group
- Cation < atom < anion; isoelectronic radius falls as nuclear charge rises
- IE: increases across a period, decreases down a group; exceptions Be > B, N > O
- ΔₑgH: most negative for halogens; exception Cl more negative than F
- Electronegativity: increases across a period, decreases down a group; F is highest (4.0)
- Metallic character ↓ across, ↑ down; oxides: metal = basic, non-metal = acidic, borderline = amphoteric
- Z > 100 nomenclature: nil-un-bi-tri-quad-pent-hex-sept-oct-enn + ium
Next Chapter: Chapter 4 - Chemical Bonding and Molecular Structure
Class 11 Chemistry · Chapter 3 – swipe through all 9 cards to understand the whole chapter.
The Modern Periodic Law
Element properties repeat in a regular cycle as atomic number rises — that one idea organises the whole chapter.
Mendeleev used atomic MASS; Moseley proved Z is the true key.
- Z (proton count), not mass, fixes an element’s place
- Trap: ‘atomic mass’ answer belongs to Mendeleev, not the modern law
- Moseley used X-ray frequencies: √ν = a(Z − b)
Periods, Groups & Blocks
The long-form table has 7 rows, 18 columns, and 4 blocks named after the subshell being filled.
f-block (lanthanoids/actinoids) parked below the main table.
- Period no. = principal shell n; group = similar valence config
- Blocks = s, p, d, f (last subshell filled)
- Period 6 has 32 (don’t forget the 14 lanthanoids)
Configuration → Position
Read an element’s block, group and period straight off its electronic configuration.
Period = highest n; block = subshell getting the last electron.
- s-block group = number of ns electrons (1 or 2)
- Cl (…3s23p5): valence 7 → group 10+7 = 17, period 3
- Don’t count a buried, filled 3d10 as valence — it’s core
Atomic & Ionic Radii
Size shrinks across a period and grows down a group — and most other trends just follow size.
Cation < parent atom; anion > parent atom.
- Across: same shell + more protons → tighter, smaller
- Na⁺ < Na (lost a shell); Cl⁻ > Cl (added repulsion)
- Ga ≈ Al, even slightly smaller (poor 3d shielding)
Isoelectronic Species
Same electron count, different proton count — more protons squeeze the cloud smaller.
Compare PROTONS, not electrons (electrons are equal).
- Protons: O 8, F 9, Na 11, Mg 12, Al 13
- Fewer protons → weakest pull → biggest size
- Reverse the list for increasing order of size
Ionization Enthalpy
Energy to pull the loosest electron off an isolated gaseous atom — always endothermic.
Across → IE ↑ (smaller, higher Zeff); down ↓ IE ↓. Units kJ/mol.
- Dips: B < Be and O < N (half/full-filled stability of Be 2s2, N 2p3)
- Same dips repeat in period 3: Al < Mg, S < P
- A sudden big jump in successive IE reveals the group number
Electron Gain Enthalpy
Enthalpy change when a gaseous atom gains an electron — usually negative (energy released).
Noble gases: large POSITIVE values (electron must start a new shell).
- More negative across a period, less negative down a group
- Cl beats F: F’s tiny 2p shell crowds the incoming electron
- Halogens have the most negative ΔegH
Electronegativity
An atom’s tendency, inside a bond, to attract the shared electron pair — a relative, unitless number.
Bigger EN difference → more polar bond (e.g. H–F more polar than H–Cl).
- Fluorine is the most electronegative element (Pauling 4.0)
- Rises across (Zeff ↑), falls down a group (bigger atoms)
- F has highest EN, but Cl has the more negative ΔegH — don’t mix them
Chemical Trends & Oxides
Valence, metallic character and oxide nature all march in periodic step — the most-tested payoff.
Acidic character of oxides increases across a period.
- Amphoteric oxides: Al2O3, ZnO, BeO (react with acid AND base)
- Hydride valence (higher p-block) = 8 − outer electrons → N is 3 in NH3
- Diagonal pairs: Li–Mg, Be–Al, B–Si (similar charge/size ratio)
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Frequently Asked Questions
The modern periodic law states that the physical and chemical properties of elements are a periodic function of their atomic numbers. This corrected Mendeleev’s earlier law, which was based on atomic mass, after Moseley showed that atomic number (Z), not mass, is the fundamental property of an element.
The period equals the highest principal quantum number (n) in the configuration, and the block is the subshell that receives the last electron (s, p, d or f). For groups, an s-block group equals the number of valence electrons, while a p-block group equals 10 plus the number of valence electrons; for example chlorine ending in 3s2 3p5 has 7 valence electrons, so it is group 17, period 3, p-block.
Across a period atomic radius decreases and ionization enthalpy increases, because the nuclear charge rises while electrons fill the same shell, raising the effective nuclear charge. Down a group atomic radius increases and ionization enthalpy decreases, because new shells are added and shielding grows.
Yes, Classification of Elements and Periodicity is part of the NEET Class 11 Chemistry syllabus and is regularly examined, usually contributing about 1 to 2 questions per paper. It is considered scoring because most questions test periodic trends, exceptions and the position of an element rather than lengthy calculations.
Ionization enthalpy is the energy needed to remove the most loosely bound electron from a gaseous atom and is always positive (endothermic). Electron gain enthalpy is the enthalpy change when a gaseous atom gains an electron and is usually negative (energy released), being most negative for the halogens.