Classification of Elements and Periodic Table
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1. Why Elements Are Classified
A systematic arrangement of elements in which recurring similarities in electronic structure and properties become visible.
Classification reduces a large number of elements into meaningful groups. Elements in the same group often show similar chemical behavior because they have related valence-shell electronic configurations.
Diagram 1: Electronic origin of periodicity
2. Mendeleev’s Periodic Law — Historical Foundation
The physical and chemical properties of elements are periodic functions of their atomic weights (atomic masses), according to the historical formulation.
In 1869 Dmitri Mendeleev organized the known elements mainly in increasing atomic mass while placing chemically similar elements together. His system was historically important because it revealed recurring patterns and allowed prediction of undiscovered elements.
3. Main Features of Mendeleev’s Periodic Table
The uploaded notes describe the historical table using the older nomenclature of its time:
- Elements were primarily arranged according to increasing atomic weight.
- Horizontal rows were called periods.
- Vertical columns were called groups.
- Many groups were divided into A and B subgroups.
- Noble gases were later placed in a separate zero group in the old notation.
- Short and long periods were recognized.
4. Defects / Limitations of Mendeleev’s Periodic Table
Pages 3–4 of the uploaded PDF list several classical limitations.
- Position of hydrogen: hydrogen resembles both alkali metals and halogens, so its placement was ambiguous.
- Anomalous atomic-mass pairs: strict increasing-mass order could not explain pairs such as Ar–K and Co–Ni.
- Isotopes: isotopes have different masses but identical chemical identity, so mass-based placement is problematic.
- Lanthanides and actinides: their position was not integrated naturally into the main historical table.
- Some similar/dissimilar groupings: older subgroup arrangements could place chemically unlike elements near one another or separate related ones.
- Cause of periodicity: atomic mass did not explain the fundamental reason for recurring properties.
Diagram 2: Atomic-number ordering solves key mass-order problems
5. Merits of Mendeleev’s Periodic Table
- Provided a systematic study of known elements.
- Grouped chemically similar elements.
- Left gaps for undiscovered elements and enabled successful predictions of their properties.
- Helped correct some doubtful atomic masses.
- Established the idea that elemental properties recur periodically.
6. Modern Periodic Law
The physical and chemical properties of elements are periodic functions of their atomic numbers.
Atomic number is more fundamental than atomic mass because it identifies nuclear charge and determines the number of electrons in a neutral atom. Electron configuration, in turn, controls chemical behavior.
7. Periodicity and Cause of Periodicity
The recurrence of similar properties at regular intervals when elements are arranged in increasing atomic number.
The fundamental cause is the recurrence of similar valence-shell electronic configurations. For example, Group 1 elements have the general outer configuration ns¹ and therefore show similar valency and chemistry.
| Group | General valence configuration | Typical behavior |
|---|---|---|
| 1 | ns¹ | Electropositive metals; often +1 ions |
| 2 | ns² | Often +2 ions |
| 17 | ns²np⁵ | Strong tendency to gain/share one electron |
| 18 | ns²np⁶ (He: 1s²) | Very stable valence shell |
8. Modern Periodic Table: Periods and Groups
A horizontal row of the periodic table.
A vertical column of the periodic table.
- The modern periodic table has 7 periods.
- It has 18 groups.
- Elements in the same group commonly have related valence-shell configurations.
- The period number usually indicates the highest occupied principal shell in the ground-state configuration.
| Period | Number of elements | General description |
|---|---|---|
| 1 | 2 | Shortest period |
| 2 | 8 | Short period |
| 3 | 8 | Short period |
| 4 | 18 | Long period |
| 5 | 18 | Long period |
| 6 | 32 | Includes lanthanides |
| 7 | 32 when complete | Includes actinides; contains synthetic/transuranium elements |
9. Classification into s, p, d and f Blocks
The region of the periodic table identified by the type of subshell receiving the differentiating (last-added) electron in the ground-state configuration.
Diagram 3: Location of s-, p-, d- and f-block elements
10. s-Block Elements
- Differentiating electron enters an s subshell.
- General outer configuration: ns¹–².
- Mainly Groups 1 and 2; hydrogen and helium have s-subshell configurations although their chemical placement is special.
- Most are electropositive metals with relatively low ionization energies.
- Group 1: alkali metals; Group 2: alkaline-earth metals.
11. p-Block Elements
- Differentiating electron enters a p subshell.
- General outer configuration: ns²np¹–⁶.
- Located mainly in Groups 13–18.
- Contains metals, metalloids and non-metals.
- Includes halogens and noble gases.
- s- and p-block elements are commonly called representative elements.
12. d-Block Elements
- Differentiating electron enters a d subshell of the penultimate shell.
- General pattern: (n−1)d¹–¹⁰ ns⁰–².
- Located mainly in Groups 3–12.
- Often called transition elements when they satisfy the transition-element definition.
- Many show variable oxidation states, complex formation and colored ions/compounds.
13. f-Block Elements
- Differentiating electron enters an f subshell.
- General pattern involves filling of (n−2)f orbitals.
- Shown in two rows below the main body of the table for compactness.
- Lanthanides and actinides are commonly called inner-transition elements.
14. Nuclear Charge and Effective Nuclear Charge
The total positive charge of the nucleus, determined by the number of protons. In magnitude it is proportional to atomic number Z.
The net positive attraction experienced by an electron in a multi-electron atom after accounting qualitatively for shielding by other electrons.
Inner-shell electrons partly screen outer electrons from the full nuclear charge. Across a period, nuclear charge increases while shielding does not increase enough to cancel it, so effective nuclear attraction generally increases.
Diagram 4: Qualitative shielding effect
15. Atomic Radius / Atomic Size
A measure of atomic size. Because an electron cloud has no sharp outer boundary, atomic radius is defined operationally from distances between nuclei.
Common Radius Definitions
| Radius | Definition |
|---|---|
| Covalent radius | Half the internuclear distance between two identical covalently bonded atoms. |
| Metallic radius | Half the internuclear distance between nearest identical metal atoms in a metallic crystal. |
| van der Waals radius | Half the distance between nuclei of two nearest non-bonded identical atoms. |
Trend Across a Period
Atomic radius generally decreases from left to right because electrons are added to the same principal shell while effective nuclear charge increases.
Trend Down a Group
Atomic radius generally increases down a group because a new principal shell is added at each successive period and shielding increases.
Diagram 5: General atomic-size trend
16. Ionic Radius
A measure of the effective size of an ion in an ionic crystal or related environment.
Cations
A cation is generally smaller than its parent atom because electron loss reduces electron–electron repulsion and may remove the outermost shell.
Anions
An anion is generally larger than its parent atom because added electrons increase electron–electron repulsion while nuclear charge is unchanged.
Diagram 6: Effect of electron loss or gain on size
17. Isoelectronic Species
Atoms or ions having the same number of electrons but different nuclear charges.
Within an isoelectronic series, radius generally decreases as nuclear charge increases.
All species have 10 electrons, but the number of protons rises from O to Al. The same electron cloud is therefore pulled more strongly by the increasingly positive nucleus.
18. Valency and Periodicity
The combining capacity of an atom, usually related to the number of electrons lost, gained or shared to reach a stable valence-shell arrangement.
Across a main-group period, simple valency commonly increases from 1 to 4 and then decreases toward zero if expressed in the older elementary combining-capacity convention. Modern chemistry often discusses oxidation states and bonding instead of a single fixed valency.
| Second-period element | Li | Be | B | C | N | O | F | Ne |
|---|---|---|---|---|---|---|---|---|
| Simple valency pattern | 1 | 2 | 3 | 4 | 3 | 2 | 1 | 0 |
19. Ionization Energy
The minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.
General Trend
- Across a period: generally increases.
- Down a group: generally decreases.
Across a period, effective nuclear charge rises and atomic size generally falls. Down a group, the valence electron is farther from the nucleus and more shielded.
Diagram 7: General ionization-energy trend
20. Factors Affecting Ionization Energy
20.1 Atomic Size
As atomic size increases, the valence electron is farther from the nucleus, so it is generally easier to remove.
20.2 Nuclear / Effective Nuclear Charge
Greater effective nuclear attraction holds the valence electron more strongly and generally raises ionization energy.
20.3 Shielding Effect
Greater shielding by inner electrons reduces the attraction experienced by an outer electron, generally lowering ionization energy.
20.4 Electronic Configuration
Half-filled and fully filled subshells can have extra stability. This explains well-known exceptions such as:
B begins removing an electron from a higher-energy 2p subshell, while Be has a filled 2s² subshell. Oxygen has one paired 2p electron, increasing repulsion compared with the half-filled 2p³ arrangement of nitrogen.
21. Successive Ionization Energies
The energies required to remove electrons one after another from the same species in the gas phase.
After each electron is removed, the remaining electrons experience greater attraction per electron. A very large jump after a certain ionization can indicate that removal has reached a stable inner shell.
22. Electron Affinity
The energy change associated with adding an electron to an isolated gaseous atom to form an anion.
Older school notes often define electron affinity as the “energy released” when an electron is added. Modern thermochemical sign conventions distinguish the numerical electron affinity from electron-gain enthalpy; for NEB conceptual answers, state the definition clearly and focus on the trend.
General Trend
- Across a period, the tendency to accept an electron generally increases.
- Down a group, it generally decreases as size and shielding increase.
- There are important irregularities because subshell stability, repulsion and atomic size matter.
Factors
- Atomic size.
- Effective nuclear charge.
- Electronic configuration.
- Electron–electron repulsion.
23. Electronegativity
The relative tendency of an atom in a chemical bond to attract the shared electron pair toward itself.
General Trend
- Across a period: generally increases from left to right.
- Down a group: generally decreases.
- Fluorine is the most electronegative element on the commonly used Pauling scale.
Factors Affecting Electronegativity
- Atomic size.
- Effective nuclear charge.
- Shielding.
- Oxidation state and bonding environment.
Diagram 8: Electronegativity rises generally toward the upper-right
24. Metallic and Non-metallic Character
The tendency of an atom to lose electrons and form positive ions, accompanied by typical metallic physical and chemical behavior.
General Trend
- Across a period: metallic character generally decreases; non-metallic character increases.
- Down a group: metallic character generally increases.
The trend follows the increasing or decreasing ease of electron loss. Larger atoms with lower ionization energy tend to be more metallic.
Diagram 9: General metallic-character trend
25. Master Periodic-Trend Summary
| Property | Across a period → | Down a group ↓ | Main reason |
|---|---|---|---|
| Atomic radius | Generally decreases | Generally increases | Zeff vs number of shells |
| Ionic radius | Depends on ion type; generally decreases within comparable/isoelectronic sets as Z rises | Generally increases in same-charge group ions | Shells and nuclear attraction |
| Ionization energy | Generally increases | Generally decreases | Effective nuclear attraction and size |
| Electron affinity tendency | Generally increases, irregularly | Generally decreases, with exceptions | Attraction, size, configuration |
| Electronegativity | Generally increases | Generally decreases | Bonding-electron attraction |
| Metallic character | Generally decreases | Generally increases | Ease of electron loss |
Across a period: radius ↓, ionization energy ↑, electronegativity ↑, metallic character ↓.
Down a group: radius ↑, ionization energy ↓, electronegativity ↓, metallic character ↑.
Diagram 10: Quick direction map for major trends
26. Worked Examples
Element: Na, electronic configuration 1s² 2s² 2p⁶ 3s¹.
The differentiating electron enters an s subshell, so Na belongs to the s block.
Element: Cl, Z = 17 → [Ne] 3s²3p⁵.
Highest principal shell n = 3 → Period 3. Valence configuration ns²np⁵ → Group 17.
Arrange Na, Mg, Al in decreasing atomic radius.
All are in Period 3. Effective nuclear charge generally increases from left to right, so:
Na > Mg > AlWhich is larger: Na or Na⁺?
Na is larger. Na⁺ has lost its 3s electron and its outer occupied shell becomes n=2, so the ion contracts strongly.
Arrange O²⁻, F⁻, Na⁺ and Mg²⁺ in decreasing radius.
Each has 10 electrons. Nuclear charge rises from O to Mg, so radius decreases:
O²⁻ > F⁻ > Na⁺ > Mg²⁺Which generally has higher first ionization energy: Li or Na?
Li, because its valence electron is closer to the nucleus and experiences less shielding.
Why is IE(Be) greater than IE(B)?
Be has a filled 2s² subshell. In B, the electron removed is from the higher-energy 2p¹ subshell, so removal is easier.
Why is IE(N) greater than IE(O)?
N has a relatively stable half-filled 2p³ arrangement. O has 2p⁴, containing one paired set whose electron–electron repulsion makes removal easier.
An element has a large jump between IE₂ and IE₃. What does that suggest?
Two electrons can be removed comparatively more easily before reaching a stable inner shell. This suggests a Group 2-type valence configuration.
Which is more electronegative: C or F?
F. Across Period 2, effective nuclear attraction increases and atomic size decreases, so the attraction for bonding electrons generally rises.
Which is more metallic: Na or Cl?
Na. Sodium loses its valence electron much more readily, while chlorine has a stronger tendency to attract/gain electrons.
Electronic configuration: [Ar] 3d⁵4s².
The differentiating electron enters a d subshell → d-block element.
27. Mendeleev vs Modern Periodic Table
| Feature | Mendeleev system | Modern system |
|---|---|---|
| Fundamental basis | Atomic mass/weight | Atomic number |
| Periodic law | Properties periodic with atomic weight | Properties periodic with atomic number |
| Isotopes | Mass basis creates difficulty | Same Z → same position |
| Ar–K / Co–Ni anomalies | Required exceptions to strict mass order | Correctly ordered by Z |
| Groups | Historical A/B subgroup system | 18 groups |
| Electronic explanation | Not available originally | Periodicity explained by electronic configuration |
| Blocks | Not the organizing principle | s, p, d and f blocks |
28. Source-PDF Points vs Current Main-Syllabus Focus
| Topic in uploaded PDF | How it is used here |
|---|---|
| Mendeleev periodic law, features, defects and merits | Retained as historical foundation and exam-support material |
| Older IA–VIIIA / 0-group terminology | Explained as source terminology; replaced by modern groups 1–18 in main notes |
| Modern periodic law/table | Core current syllabus |
| s/p/d/f block classification | Core current syllabus |
| Atomic radius, ionic radius | Core current syllabus |
| Isoelectronic species | Useful explanatory application of ionic-radius trend |
| Valency | Retained as source-background periodicity concept |
| Ionization energy + factors + successive IE | Core trend plus useful source depth |
| Electron affinity | Core current syllabus |
| Electronegativity | Core current syllabus |
| Metallic character | Core current syllabus; added explicitly as a main section |
29. Common Exam Mistakes
- Writing the modern periodic law in terms of atomic mass instead of atomic number.
- Writing that the modern periodic table has 8 or 9 groups; it has 18 groups.
- Using old A/B subgroup notation as if it were current IUPAC group numbering.
- Calling period a vertical column or group a horizontal row.
- Forgetting that similar group chemistry arises mainly from related valence-shell configurations.
- Calling every d-block element a transition element without considering the transition-element definition.
- Forgetting that f-block elements are shown separately only for compactness; they belong to periods 6 and 7.
- Confusing nuclear charge with effective nuclear charge.
- Thinking shielding increases the attraction of the nucleus for the outer electron; shielding reduces the effective attraction.
- Defining atomic radius as an exact physical boundary of an atom.
- Forgetting the difference among covalent, metallic and van der Waals radii.
- Saying atomic radius increases from left to right across a normal period.
- Saying atomic radius decreases down a group.
- Writing a cation as larger than its parent atom under ordinary comparison.
- Writing an anion as smaller than its parent atom under ordinary comparison.
- For isoelectronic species, comparing radius by electron count alone; they have the same electron count, so nuclear charge decides the order.
- Memorizing periodic arrows without explaining effective nuclear charge, shells and shielding.
- Defining ionization energy for a solid or solution instead of an isolated gaseous species.
- Writing IE₂ < IE₁; successive ionization energies increase.
- Ignoring the Be/B and N/O ionization-energy exceptions.
- Confusing electron affinity with electronegativity.
- Electron affinity concerns adding an electron to an isolated gaseous atom; electronegativity concerns attraction of shared electrons in a bond.
- Assuming electron affinity follows a perfectly smooth trend.
- Using electronegativity for a free isolated atom as though it were an independently measurable atomic energy.
- Forgetting that metallic character generally changes opposite to electronegativity across/down the table.
- Repeating source statements such as “seventh period contains 19 elements and is incomplete” as current fact; the modern seventh period is complete through element 118.
30. Important Exam Questions
Very Short / Short Questions
- Define periodic table and periodicity.
- State Mendeleev’s periodic law.
- Write two merits and four limitations of Mendeleev’s periodic table.
- Why was the position of hydrogen difficult in Mendeleev’s table?
- How did isotopes create a problem for a mass-based periodic law?
- State the modern periodic law.
- Why is atomic number more fundamental than atomic mass for classification?
- What is the fundamental cause of periodicity?
- Define group and period.
- How many groups and periods are in the modern periodic table?
- Define an s-block element and give its general configuration.
- Define a p-block element and give its general configuration.
- Give the general configuration of d-block elements.
- What are inner-transition elements?
- Define nuclear charge.
- Define effective nuclear charge.
- What is shielding/screening effect?
- Define atomic radius.
- Differentiate covalent radius and van der Waals radius.
- Explain the trend of atomic radius across a period.
- Explain the trend of atomic radius down a group.
- Why is Na⁺ smaller than Na?
- Why is Cl⁻ larger than Cl?
- Define isoelectronic species.
- Arrange O²⁻, F⁻, Na⁺ and Mg²⁺ in decreasing ionic radius with reason.
- Define first ionization energy.
- Explain the trend of ionization energy across a period and down a group.
- State factors affecting ionization energy.
- Why is IE(Be) greater than IE(B)?
- Why is IE(N) greater than IE(O)?
- Define successive ionization energies.
- Define electron affinity.
- State the general trend of electron affinity.
- Define electronegativity.
- State the general trend of electronegativity.
- Define metallic character and explain its periodic trend.
Long / Descriptive Questions
- Discuss Mendeleev’s periodic law, main features, merits and defects.
- Explain how the modern periodic law overcomes major limitations of Mendeleev’s system.
- Describe the modern periodic table and classification into groups, periods and blocks.
- Compare the general characteristics and electronic configurations of s-, p-, d- and f-block elements.
- Explain nuclear charge, shielding and effective nuclear charge and connect them with periodic trends.
- Define atomic radius and explain its trends across periods and down groups.
- Explain ionic radius, parent-atom/ion size relations and isoelectronic-series ordering.
- Define ionization energy and discuss the factors affecting it.
- Explain periodic variation of ionization energy with important configuration-based exceptions.
- Explain successive ionization energies and their significance.
- Define electron affinity and discuss its general periodic variation and irregularities.
- Define electronegativity and explain its general periodic trend.
- Explain metallic and non-metallic character on the basis of periodic trends.
- Prepare a comparative table showing the direction and causes of all major periodic properties.
Reasoning / “Give Reason” Questions
- Atomic radius decreases across a period.
- Atomic radius increases down a group.
- Cations are smaller than their parent atoms.
- Anions are larger than their parent atoms.
- In an isoelectronic series, radius decreases with increasing nuclear charge.
- Ionization energy generally rises across a period.
- Ionization energy generally falls down a group.
- Be has higher first ionization energy than B.
- N has higher first ionization energy than O.
- Electronegativity generally rises across a period.
- Metallic character increases down a group.
- Noble gases show very low chemical reactivity.
Study Unit 4 in three blocks: (1) Mendeleev → modern periodic law, (2) groups/periods/s-p-d-f blocks, (3) periodic properties and their causes. In trend questions, always write both the direction and the reason using shells, shielding and effective nuclear charge.
31. One-Minute Revision
- Modern periodic law: properties are periodic functions of atomic number.
- Mendeleev’s historical law used atomic mass/weight.
- Modern table: 7 periods and 18 groups.
- Period = horizontal row; group = vertical column.
- Periodicity comes from recurring valence-shell electronic configurations.
- s block: ns¹–².
- p block: ns²np¹–⁶.
- d block: generally (n−1)d¹–¹⁰ns⁰–².
- f block: differentiating electron enters an f subshell.
- Groups 1–2 mainly form the s block; Groups 13–18 mainly form the p block.
- Groups 3–12 occupy the d block.
- Lanthanides and actinides form the f block.
- Nuclear charge depends on proton number Z.
- Effective nuclear charge is the net attraction after shielding.
- Across a period, Zeff generally increases.
- Atomic radius generally decreases across a period.
- Atomic radius generally increases down a group.
- Cation < parent atom; anion > parent atom.
- Isoelectronic radius decreases as nuclear charge increases.
- Ionization energy generally increases across a period.
- Ionization energy generally decreases down a group.
- Successive ionization energies increase: IE₁ < IE₂ < IE₃.
- Large successive-IE jumps help reveal valence-electron count.
- Be > B and N > O are important first-IE exceptions.
- Electron affinity describes electron addition to a gaseous atom.
- Electron-affinity trends are irregular compared with simple radius trends.
- Electronegativity = attraction for shared bonding electrons.
- Electronegativity generally increases across and decreases down.
- Fluorine is the most electronegative element on the Pauling scale.
- Metallic character generally decreases across and increases down.
- Non-metallic character generally changes in the opposite direction.
32. Diagram Practice
- Electronic origin of periodicity.
- Mendeleev mass-order problems and modern atomic-number solution.
- s-, p-, d- and f-block layout.
- Shielding and effective nuclear charge.
- Atomic-radius trend.
- Cation vs atom vs anion size.
- Ionization-energy trend.
- Electronegativity trend.
- Metallic-character trend.
- Master periodic-trend compass.
Discussion
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