The Periodic Table — Class 10 Physical Science
"Mendeleev left gaps in his table. He predicted elements that didn't exist yet — their properties, their atomic masses. And when they were discovered, he was RIGHT. That is the power of a good scientific theory."
1. Early Attempts at Classification
Dobereiner's Triads (1817): Groups of 3 elements with similar properties. Middle element's atomic mass ≈ average of extremes. Limited — very few triads found. Newlands' Law of Octaves (1864): Every 8th element has similar properties — like musical notes. Worked only up to calcium. Ridiculed.
2. Mendeleev's Periodic Table (1869)
Arranged elements by INCREASING ATOMIC MASS. Elements with SIMILAR PROPERTIES in the same VERTICAL COLUMN (Group). GENIUS: Left GAPS for undiscovered elements. PREDICTED their properties (eka-aluminium → gallium. eka-silicon → germanium). When discovered, his predictions were REMARKABLY ACCURATE. Limitations: Position of hydrogen uncertain (resembles both alkali metals AND halogens). Isotopes (same element, different mass) — no place. Anomalous pairs (Argon mass 40 before Potassium mass 39 — based on properties, not mass order).
3. Modern Periodic Law (Mosley, 1913)
Properties of elements are a PERIODIC FUNCTION of their ATOMIC NUMBER. Arranged by increasing ATOMIC NUMBER (Z = number of protons). This SOLVED Mendeleev's anomalies.
4. Structure of the Modern Periodic Table
7 PERIODS (horizontal rows). Period number = number of SHELLS. 18 GROUPS (vertical columns). Group number = number of VALENCE ELECTRONS (for Groups 1-2, 13-18).
Key Groups
| Group | Name | Valence e⁻ | Nature | Examples |
|---|---|---|---|---|
| 1 | Alkali Metals | 1 | Highly reactive metals. Form +1 ions. | Li, Na, K |
| 2 | Alkaline Earth Metals | 2 | Reactive metals. Form +2 ions. | Mg, Ca |
| 17 | Halogens | 7 | Highly reactive non-metals. Form −1 ions. | F, Cl, Br, I |
| 18 | Noble Gases | 8 (full) | EXTREMELY unreactive. | He, Ne, Ar |
5. Periodic Trends
Atomic Size (Radius) : DECREASES across a period (← →). More protons pull electrons CLOSER. INCREASES down a group (↓). More shells added. Metallic Character: DECREASES across period. INCREASES down group. 'The most metallic element: FRANCIUM (bottom left). The most non-metallic: FLUORINE (top right).'
6. Common Mistakes
- 'Mendeleev arranged by atomic number' — He arranged by ATOMIC MASS. Mosley CORRECTED it to atomic number.
- 'Noble gases are in Group 18 because they have 8 valence electrons' — Helium has 2 valence electrons but is STILL in Group 18 (its shell is FULL).
7. AP SSC Exam Focus
| Topic | Marks |
|---|---|
| Mendeleev's table and predictions | 3-4 |
| Modern periodic law | 2-3 |
| Group names and properties | 3-4 |
| Periodic trends | 2-3 |
8. Blocks of the Periodic Table
The modern periodic table is divided into FOUR blocks based on which subshell the LAST electron enters.
s-Block Elements
| Position | Groups 1 and 2 |
|---|---|
| Electron configuration | ns¹ or ns² |
| Properties | Reactive metals. Form basic oxides. Good conductors. |
| Examples | Li, Na, K (Group 1), Be, Mg, Ca (Group 2) |
| Trend | Reactivity INCREASES down the group |
'The most reactive metal in the s-block is FRANCIUM at the very bottom of Group 1 — it reacts EXPLOSIVELY with water.'
p-Block Elements
| Position | Groups 13 to 18 |
|---|---|
| Electron configuration | ns² np¹ to ns² np⁶ |
| Properties | Metals, non-metals, metalloids. Range of properties. |
| Examples | B, C, N, O, F (Group 13-17), Ne, Ar (Group 18) |
| Key group | Halogens (Group 17) — most reactive non-metals |
'Metalloids like Silicon and Germanium are in the p-block. They have properties BETWEEN metals and non-metals — used in SEMICONDUCTORS.'
d-Block Elements (Transition Metals)
| Position | Groups 3 to 12 |
|---|---|
| Electron configuration | (n−1)d¹⁻¹⁰ ns¹⁻² |
| Properties | Hard, high melting points. Form COLOURED compounds. Variable oxidation states. Good catalysts. |
| Examples | Fe, Cu, Zn, Ni, Cr, Mn, Ag, Au, Pt |
| Why 'Transition' | They TRANSITION between s-block and p-block in the table |
'Transition metals form coloured compounds — CuSO₄ is BLUE, KMnO₄ is PURPLE, FeCl₃ is BROWN. This is UNIQUE to d-block elements.'
f-Block Elements (Inner Transition Metals)
| Position | Below the main table (2 rows) |
|---|---|
| Electron configuration | (n−2)f¹⁻¹⁴ (n−1)d⁰⁻¹ ns² |
| Series | Lanthanoids (4f series) and Actinoids (5f series) |
| Properties | Radioactive (actinoids). Lanthanoids are 'rare earth metals'. |
| Examples | Ce, Eu, Gd (Lanthanoids), U, Pu, Th (Actinoids) |
9. Periodic Trends — Detailed
Atomic Size (Atomic Radius)
| Direction | Change | Reason |
|---|---|---|
| Across a period (← →) | DECREASES | Nuclear charge INCREASES. Protons pull electrons CLOSER. Same shell. |
| Down a group (↑ ↓) | INCREASES | NEW SHELLS are added. Each new shell is further from the nucleus. |
'Across a period, although new protons are added, the electrons go into the SAME shell. More protons = stronger pull = smaller radius. Down a group, each new shell is FURTHER out — size ALWAYS increases.'
Example: Arrange Na (Z=11), Mg (Z=12), Al (Z=13), Si (Z=14) in increasing atomic size. All are in Period 3. Across a period, size DECREASES. Increasing order: Si < Al < Mg < Na (correct)
Ionisation Energy (IE)
The energy required to REMOVE the MOST LOOSELY BOUND electron from a gaseous atom.
| Direction | Change | Reason |
|---|---|---|
| Across a period | INCREASES | More protons = tighter hold on electrons |
| Down a group | DECREASES | Outer electron is FURTHER from nucleus, EASIER to remove |
'Noble gases have the HIGHEST ionisation energy — they DON'T want to lose electrons. Alkali metals have the LOWEST — they WANT to lose one electron to achieve noble gas configuration.'
Valency — Across Periods and Groups
| Period 2 Element | Li | Be | B | C | N | O | F | Ne |
|---|---|---|---|---|---|---|---|---|
| Valence Electrons | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 |
| Valency | 1 | 2 | 3 | 4 | 3 | 2 | 1 | 0 |
'Valency across a period: INCREASES from 1 to 4, then DECREASES to 0 at noble gas. This pattern REPEATS in every period.'
Electronegativity
The tendency of an atom to ATTRACT a shared pair of electrons toward itself in a covalent bond.
| Direction | Change | Reason |
|---|---|---|
| Across a period | INCREASES | Smaller size, more protons = stronger pull on bonding electrons |
| Down a group | DECREASES | Larger size, more inner shells SHIELD the nucleus |
'Fluorine (F) is the MOST electronegative element (EN ≈ 4.0). Francium (Fr) is the LEAST (EN ≈ 0.7).'
| Element | F | O | Cl | N | Br | I | C | H | Na | Fr |
|---|---|---|---|---|---|---|---|---|---|---|
| EN | 4.0 | 3.5 | 3.0 | 3.0 | 2.8 | 2.5 | 2.5 | 2.1 | 0.9 | 0.7 |
10. Trends — How to Remember
'Use this mnemonic for the pattern across periods: Ape (Atomic size) Dies (Decreases), IE Goes (Ionisation Energy increases), EN Goes (Electronegativity increases). Down groups: ALL OPPOSITE.'
| Property | Across Period (→) | Down Group (↓) |
|---|---|---|
| Atomic Size | DECREASES | INCREASES |
| Ionisation Energy | INCREASES | DECREASES |
| Electronegativity | INCREASES | DECREASES |
| Metallic Character | DECREASES | INCREASES |
| Electron Affinity | INCREASES (generally) | DECREASES (generally) |
11. Why are Elements in the Same Group Similar?
Elements in the SAME GROUP have the SAME NUMBER of VALENCE ELECTRONS.
- Group 1 elements all have 1 valence electron → ALL form +1 ions, react similarly with water.
- Group 17 elements all have 7 valence electrons → ALL form −1 ions, react similarly with metals.
'The chemistry of an element is determined by its VALENCE ELECTRONS. Same group = same valence electrons = similar chemical properties.'
12. Common Mistakes — Expanded
| Mistake | Correct Fact |
|---|---|
| 'Mendeleev arranged by atomic number' | He arranged by ATOMIC MASS. Mosley corrected to atomic number. |
| 'All noble gases have 8 electrons' | Helium has 2 valence electrons. Its shell is FULL with 2. |
| 'Atomic size increases across a period' | FALSE. It DECREASES — more protons pull electrons closer. |
| 'Electronegativity decreases across a period' | FALSE. It INCREASES — atoms get smaller, pull electrons more strongly. |
| 'Group number = valence electrons for all groups' | True only for Groups 1, 2, and 13-18. For transition metals (Groups 3-12), it's NOT this simple. |
| 'Metals are only on the left' | True for most metals, but p-block has some metals too (Al, Sn, Pb, Bi). |
13. Worked Examples
Example 1: Which element has the larger atomic radius: Na or Cl? Why? Na has LARGER atomic radius. Both are in Period 3, but Na is on the LEFT. Across a period, atomic size DECREASES due to increasing nuclear charge. Na (Z=11) has less pull on its outermost electron than Cl (Z=17). (correct)
Example 2: Arrange F, Cl, Br, I in increasing order of electronegativity. Down Group 17, electronegativity DECREASES. Increasing order: I < Br < Cl < F (correct)
Example 3: An element has electronic configuration 2, 8, 5. Find its: a) Period number — 3 (3 shells) b) Group number — 15 (5 valence electrons, Group 15) c) Block — p-block (last electron enters p subshell) d) Nature — Non-metal (on the right side of periodic table) e) Valency — 3 (needs 3 electrons to complete octet) (correct)
14. AP SSC Board Exam Focus — Expanded
| Topic | Marks | Frequency |
|---|---|---|
| Mendeleev's predictions and limitations | 3-4 | Very High |
| Modern periodic law (Mosley) | 2-3 | Very High |
| Group names and properties | 3-4 | Very High |
| Periodic trends (size, IE, EN) | 4-5 | Very High |
| Blocks (s, p, d, f) | 2-3 | High |
| Identifying period/group/block from config | 3 | High |
| Comparison between Mendeleev and Modern | 4 | Moderate |
15. Self-Test Questions
- State the Modern Periodic Law. How did it IMPROVE upon Mendeleev's periodic table?
- What is the relationship between valence electrons and group number for Groups 1, 2, and 13-18?
- Explain why atomic size DECREASES across a period but INCREASES down a group.
- Which element has higher electronegativity: oxygen or sulphur? Why?
- An element X has electronic configuration 2, 8, 8, 2. Find its period, group, and block. Name the element.
- Why did Mendeleev leave GAPS in his periodic table? Give an example.
- Arrange the following elements in increasing order of metallic character: Na, Al, Si, Mg.
- What are transition metals? Why are they called 'transition' elements?
- Calcium (Z=20) and Bromine (Z=35) — which has a LARGER atomic radius? Justify.
- Match the following: | Element | Block | |---------|-------| | Fe | s-block | | Na | p-block | | Cl | d-block | | U | f-block |
Answers: 1) Properties are a periodic function of ATOMIC NUMBER. Solved the anomalous pair issue. 2) For Groups 1-2: group number = valence electrons. For Groups 13-18: group number − 10 = valence electrons. 3) Across: same shell, more protons = stronger pull. Down: new shells added = larger distance. 4) OXYGEN — it is ABOVE sulphur in Group 16. Electronegativity DECREASES down a group. 5) Period 4, Group 2, s-block. CALCIUM (Ca). 6) He predicted undiscovered elements — e.g., eka-aluminium (gallium) and eka-silicon (germanium). 7) Si < Al < Mg < Na (increasing metallic character). 8) d-block elements with variable oxidation states, coloured compounds, good catalysts. Called 'transition' because they are between s-block and p-block. 9) Ca — Calcium is on the LEFT of Period 4. Atomic size DECREASES across a period. 10) Fe → d-block, Na → s-block, Cl → p-block, U → f-block.
