Chapter 5: The Periodic Table - Class 9 ICSE Chemistry

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Syllabus Scope & Reference

Covering: Dobereiner's Triads • Newland's Law of Octaves • Mendeleev's Contributions & Periodic Law • Discovery of Atomic Number (Moseley) • Modern Periodic Law • Structure of Modern Periodic Table (Groups 1–18, Periods 1–7) • Special reference to Alkali Metals (Group 1), Alkaline Earth Metals (Group 2), Halogens (Group 17), Noble Gases (Group 18) • Periodic Trends in Properties.

1. Early Attempts at Classification of Elements

1. Reasons for Classification of Elements

Towards the end of the 18th and early 19th centuries, numerous chemical elements were discovered. To study their physical and chemical properties systematically and find logical relationships among them, classification became indispensable.

It helps in studying elements and their compounds in an organized, systematic manner.
It helps in correlating the properties of elements with fundamental atomic parameters.
It aids in predicting the properties of newly discovered or undiscovered elements.

William Prout (1815) suggested that atomic mass of an element never differs and hypothesized that all elements were built from hydrogen atoms. This provided the initial spark for mass-based classification.

2. Dobereiner's Triads (1817)

German chemist Johann Wolfgang Dobereiner observed that certain elements showing close chemical similarities could be arranged in groups of three, called triads.

Dobereiner's Law of Triads

Law of Triads: When chemically similar elements are arranged in increasing order of their atomic masses in groups of three (triads), the atomic mass of the middle element is approximately equal to the arithmetic mean (average) of the atomic masses of the other two elements.

$$\text{Atomic Mass of Middle Element} \approx \frac{\text{Mass of 1st Element} + \text{Mass of 3rd Element}}{2}$$

Classic Examples of Dobereiner's Triads:

Triad Name	Elements & Symbol	Atomic Mass	Arithmetic Mean Calculation	Similar Chemical Properties
1. Alkali Metal Triad	Lithium ($\mathrm{Li}$) Sodium ($\mathrm{Na}$) Potassium ($\mathrm{K}$)	$7.0$ $23.0$ $39.0$	$\frac{7.0 + 39.0}{2} = \mathbf{23.0}$ (Exact match for $\mathrm{Na}$)	All are soft metals, monovalent (+1), react with water to form soluble hydroxides (alkalies) and hydrogen gas.
2. Alkaline Earth Triad	Calcium ($\mathrm{Ca}$) Strontium ($\mathrm{Sr}$) Barium ($\mathrm{Ba}$)	$40.0$ $87.6 \approx 88$ $137.0$	$\frac{40.0 + 137.0}{2} = \mathbf{88.5}$ (Close to $\mathrm{Sr} = 87.6$)	All are reactive metals, divalent (+2), oxides are alkaline in nature.
3. Halogen Triad	Chlorine ($\mathrm{Cl}$) Bromine ($\mathrm{Br}$) Iodine ($\mathrm{I}$)	$35.5$ $80.0$ $127.0$	$\frac{35.5 + 127.0}{2} = \mathbf{81.25}$ (Close to $\mathrm{Br} = 80.0$)	All are non-metals, monovalent (-1), react with water to form acids, form salts with metals.

Reasons for Discarding Law of Triads

1. Limited Applicability: Dobereiner failed to classify all the known elements of his time into triads. Only a few triads could be identified.
2. Failure within Same Family: The law did not hold good even within the same group of elements. For example, in the halogen family starting with Fluorine ($\mathrm{F}=19$), Chlorine ($\mathrm{Cl}=35.5$), and Bromine ($\mathrm{Br}=80$): $$\frac{19 + 80}{2} = 49.5 \neq 35.5 \quad (\text{Mean does not match Chlorine!})$$

3. Newland's Law of Octaves (1864)

English chemist John Newland arranged all known elements (56 elements at that time) in ascending order of their atomic masses and noticed a striking periodic pattern resembling musical notes.

Newland's Law of Octaves

Law of Octaves: When elements are arranged in increasing order of their atomic masses, the physical and chemical properties of every eighth element are a repetition of the properties of the first element, just as the eighth note of a musical octave is a repetition of the first note.
Musical Scale Analogy: Sa (Do), Re (Re), Ga (Mi), Ma (Fa), Pa (So), Dha (La), Ni (Ti), Sa (Do).

Western Musical Note	Do	Re	Mi	Fa	So	La	Ti
Indian Musical Note	Sa	Re	Ga	Ma	Pa	Dha	Ni
Elements Row 1	$\mathrm{H}$ (1)	$\mathrm{Li}$ (2)	$\mathrm{Be}$ (3)	$\mathrm{B}$ (4)	$\mathrm{C}$ (5)	$\mathrm{N}$ (6)	$\mathrm{O}$ (7)
Elements Row 2	$\mathrm{F}$ (8)	$\mathrm{Na}$ (9)	$\mathrm{Mg}$ (10)	$\mathrm{Al}$ (11)	$\mathrm{Si}$ (12)	$\mathrm{P}$ (13)	$\mathrm{S}$ (14)
Elements Row 3	$\mathrm{Cl}$ (15)	$\mathrm{K}$ (16)	$\mathrm{Ca}$ (17)	$\mathrm{Cr}$ (18)	$\mathrm{Ti}$ (19)	$\mathrm{Mn}$ (20)	$\mathrm{Fe}$ (21)

Merits of Newland's Classification:

Worked remarkably well for lighter elements up to Calcium ($\mathrm{Ca}$, Atomic mass 40).
Related the properties of elements directly to their atomic masses.
First scheme to demonstrate the concept of periodicity (regular repetition of properties).
Newland was awarded the prestigious Davy Medal by the Royal Society in 1887 for his work.

Reasons for Discarding Law of Octaves

1. Inapplicable to Heavier Elements: The law was applicable only up to Calcium ($\mathrm{Ca}$). Beyond Calcium, every eighth element did not possess properties similar to the first.
2. Pairing Dissimilar Elements: To fit elements into his table, Newland adjusted two elements into the same slot (e.g., Cobalt $\mathrm{Co}$ and Nickel $\mathrm{Ni}$) and placed them in the same column as Fluorine, Chlorine, and Bromine, which have completely different properties.
3. Separation of Similar Elements: Iron ($\mathrm{Fe}$), which resembles Cobalt and Nickel in properties, was placed far away in a separate column.
4. Discovery of Noble Gases: With the later discovery of noble gases ($\mathrm{He, Ne, Ar}$), the eighth element no longer had similar properties; instead, it became the ninth element!

2. Mendeleev's Periodic Table (1869)

Russian chemist Dmitri Ivanovich Mendeleev published his famous periodic table in 1869. He arranged all the 63 elements known at that time in increasing order of atomic masses in horizontal rows and vertical columns.

Mendeleev's Periodic Law

Mendeleev's Periodic Law: The physical and chemical properties of elements are a periodic function of their atomic masses.
Meaning: When elements are arranged in increasing order of their atomic masses, elements with similar physical and chemical properties recur at regular intervals (periodically).

Essential Features of Mendeleev's Periodic Table:

Groups: 8 vertical columns numbered from Group I to Group VIII. Groups I to VII were divided into sub-groups A and B. Group VIII contained 9 transition metals arranged in three triads ($\text{Fe, Co, Ni}$; $\text{Ru, Rh, Pd}$; $\text{Os, Ir, Pt}$). Noble gases were unknown at that time.
Periods: 6 horizontal rows. Elements change progressively from metallic to non-metallic character across a period.
Valency & Formulas of Compounds: Used general formulas for oxides and hydrides to classify groups:
• Group I: Oxide $\mathrm{R_2O}$, Hydride $\mathrm{RH}$
• Group II: Oxide $\mathrm{RO}$, Hydride $\mathrm{RH_2}$
• Group III: Oxide $\mathrm{R_2O_3}$, Hydride $\mathrm{RH_3}$
• Group IV: Oxide $\mathrm{RO_2}$, Hydride $\mathrm{RH_4}$ (e.g., $\mathrm{CH_4, CO_2}$)

Fig. 5.1: Mendeleev's Periodic Table (Original 8-Group Classification Scheme)

Groups $\to$	Group I	Group II	Group III	Group IV	Group V	Group VI	Group VII	Group VIII
Oxide:	$\mathrm{R_2O}$	$\mathrm{RO}$	$\mathrm{R_2O_3}$	$\mathrm{RO_2}$	$\mathrm{R_2O_5}$	$\mathrm{RO_3}$	$\mathrm{R_2O_7}$	$\mathrm{RO_4}$
Hydride:	$\mathrm{RH}$	$\mathrm{RH_2}$	$\mathrm{RH_3}$	$\mathrm{RH_4}$	$\mathrm{RH_3}$	$\mathrm{RH_2}$	$\mathrm{RH}$	—
Period 1	$\mathrm{H} = 1.008$	—	—	—	—	—	—	—
Period 2	$\mathrm{Li} = 6.94$	$\mathrm{Be} = 9.01$	$\mathrm{B} = 10.81$	$\mathrm{C} = 12.01$	$\mathrm{N} = 14.01$	$\mathrm{O} = 16.00$	$\mathrm{F} = 19.00$	—
Period 3	$\mathrm{Na} = 22.99$	$\mathrm{Mg} = 24.31$	$\mathrm{Al} = 29.98$	$\mathrm{Si} = 28.09$	$\mathrm{P} = 30.97$	$\mathrm{S} = 32.06$	$\mathrm{Cl} = 35.45$	—
Period 4 (Series 1 & 2)	$\mathrm{K} = 39.10$ $\mathrm{Cu} = 63.54$	$\mathrm{Ca} = 40.08$ $\mathrm{Zn} = 65.37$	[Eka-Al] $\mathrm{Ga}=69.72$	[Eka-Si] $\mathrm{Ge}=72.59$	$\mathrm{As} = 74.92$	$\mathrm{Se} = 78.96$	$\mathrm{Br} = 79.91$	$\mathrm{Fe}=55.85$, $\mathrm{Co}=58.93$, $\mathrm{Ni}=58.71$
Period 5 (Series 1 & 2)	$\mathrm{Rb} = 85.47$ $\mathrm{Ag} = 107.87$	$\mathrm{Sr} = 87.62$ $\mathrm{Cd} = 112.41$	$\mathrm{In} = 114.82$	$\mathrm{Sn} = 118.69$	$\mathrm{Sb} = 121.75$	$\mathrm{Te} = 127.60$	$\mathrm{I} = 126.90$	$\mathrm{Ru}=101.1$, $\mathrm{Rh}=102.9$, $\mathrm{Pd}=106.4$

Structure of Mendeleev's 1869 Periodic Table highlighting Groups I-VIII, hydride/oxide formulas, and predicted gaps ([Eka-Al], [Eka-Si]).

1. Merits of Mendeleev's Periodic Table

Systematic Study of Elements: Reduced the study of 63 individual elements to just 8 groups.
Gaps for Undiscovered Elements: Mendeleev boldly left blank spaces in his table for elements yet to be discovered, instead of forcing known elements into wrong places.

Prediction of Properties of Undiscovered Elements: He predicted the atomic masses and chemical properties of unknown elements with extraordinary accuracy:

Property	Predicted: Eka-Aluminium	Discovered: Gallium ($\mathrm{Ga}$)	Predicted: Eka-Silicon	Discovered: Germanium ($\mathrm{Ge}$)
Atomic Mass	$68$	$69.7$	$72$	$72.6$
Density	$5.9\text{ g cm}^{-3}$	$5.91\text{ g cm}^{-3}$	$5.5\text{ g cm}^{-3}$	$5.36\text{ g cm}^{-3}$
Melting Point	Low	$302\text{ K } (29.8^\circ\text{C})$	High	$1231\text{ K}$
Valency	3	3	4	4
Oxide Formula	$\mathrm{E_2O_3}$	$\mathrm{Ga_2O_3}$	$\mathrm{EO_2}$	$\mathrm{GeO_2}$

Correction of Questionable Atomic Masses: Corrected atomic masses of elements based on their position, e.g., Beryllium ($\mathrm{Be}$) mass was corrected from $13.5$ to $9.0$ ($\text{Valency } 2 \times \text{Equivalent mass } 4.5 = 9$). Also corrected Gold ($\mathrm{Au}$) and Platinum ($\mathrm{Pt}$).

2. Defects & Anomalies of Mendeleev's Periodic Table

Limitation & Anomalies in Mendeleev's Table

1. Anomalous Pairs (Inversion of Atomic Weights): Certain elements with higher atomic masses were placed BEFORE elements with lower atomic masses to keep similar elements together:
• Argon ($\text{Mass } 39.9$, $Z=18$) placed BEFORE Potassium ($\text{Mass } 39.1$, $Z=19$).
• Cobalt ($\text{Mass } 58.9$, $Z=27$) placed BEFORE Nickel ($\text{Mass } 58.6$, $Z=28$).
• Tellurium ($\text{Mass } 127.6$, $Z=52$) placed BEFORE Iodine ($\text{Mass } 126.9$, $Z=53$).

2. Position of Isotopes: Isotopes are atoms of the same element having different atomic masses (e.g., $\mathrm{^{35}Cl}$ and $\mathrm{^{37}Cl}$). If classification is based on atomic mass, isotopes must occupy separate positions, which would destroy the table layout. Mendeleev gave no separate position to isotopes.

3. Grouping of Chemically Dissimilar Elements: Coinage metals ($\mathrm{Cu, Ag, Au}$) were placed in Group IB together with highly reactive alkali metals ($\mathrm{Li, Na, K}$) in Group IA.

4. Separation of Chemically Similar Elements: Elements with similar properties were placed in different groups, e.g., Copper ($\mathrm{Cu}$) and Mercury ($\mathrm{Hg}$); Platinum ($\mathrm{Pt}$) and Gold ($\mathrm{Au}$).

5. Uncertain Position of Hydrogen: Hydrogen resembles alkali metals (monovalent, forms $\mathrm{H^+}$) as well as halogens (forms diatomic $\mathrm{H_2}$, forms hydrides $\mathrm{NaH}$). Its position was left ambiguous.

3. Modern Periodic Law & Modern Periodic Table

1. Discovery of Atomic Number (Henry Moseley, 1913)

In 1913, English physicist Henry Moseley studied the characteristic X-ray spectra emitted by different metal anodes when bombarded by cathode rays in a discharge tube. He observed that the square root of the frequency ($\sqrt{\nu}$) of characteristic X-rays was directly proportional to the atomic number ($Z$) of the element, NOT its atomic mass.

Moseley's Conclusion

Atomic Number ($Z$): The magnitude of positive charge present in the nucleus of an atom (number of protons) is called its atomic number.

Atomic Number ($Z$) = Number of Protons = Number of Electrons (in a neutral atom)

Moseley's Finding: Atomic number is a more fundamental property of an element than its atomic mass!

2. Modern Periodic Law

Modern Periodic Law

Modern Periodic Law: The physical and chemical properties of elements are a periodic function of their atomic numbers ($Z$).
Significance: When elements are arranged in increasing order of atomic numbers, elements with similar electronic configurations (same number of valence electrons) recur at regular intervals.

3. Resolution of Mendeleev's Anomalies by Modern Periodic Law

Mendeleev's Anomaly	Explanation by Modern Periodic Law ($Z$)
1. Inverted Pair: Argon & Potassium	Argon has $Z=18$ and Potassium has $Z=19$. Arranging by atomic number naturally places Argon ($Z=18$) before Potassium ($Z=19$), resolving the mass discrepancy ($\text{Ar}=39.9, \text{K}=39.1$).
2. Inverted Pair: Cobalt & Nickel	Cobalt has $Z=27$ and Nickel has $Z=28$. Cobalt ($Z=27$) comes before Nickel ($Z=28$) automatically.
3. Position of Isotopes	All isotopes of an element have the same atomic number ($Z$) (e.g., $\mathrm{^{35}_{17}Cl}$ and $\mathrm{^{37}_{17}Cl}$ both have $Z=17$). Thus, all isotopes naturally occupy the same single slot in the periodic table.

4. Structure of the Long Form of Modern Periodic Table

The Long Form of Periodic Table (Bohr's Table) is based on the electronic configuration of elements. It consists of vertical columns called Groups and horizontal rows called Periods.

Fig. 5.2: The Long Form of Modern Periodic Table (Bohr's Classification by $Z$)

Block	Group 1 (IA)	Group 2 (IIA)	Groups 3 to 12 (d-block)	Group 13 (IIIA)	Group 14 (IVA)	Group 15 (VA)	Group 16 (VIA)	Group 17 (VIIA)	Group 18 (Zero)
Period 1	₁H 2,1	—	d-block Transition Metals (3 outer shells filling, variable valencies)	—	—	—	—	—	₂He 2
Period 2	₃Li 2,1	₄Be 2,2		₅B 2,3	₆C 2,4	₇N 2,5	₈O 2,6	₉F 2,7	₁₀Ne 2,8
Period 3	₁₁Na 2,8,1	₁₂Mg 2,8,2		₁₃Al 2,8,3	₁₄Si 2,8,4	₁₅P 2,8,5	₁₆S 2,8,6	₁₇Cl 2,8,7	₁₈Ar 2,8,8
Period 4	₁₉K	₂₀Ca	₂₁Sc ... ₃₀Zn	₃₁Ga	₃₂Ge	₃₃As	₃₄Se	₃₅Br	₃₆Kr
Period 5	₃₇Rb	₃₈Sr	₃₉Y ... ₄₈Cd	₄₉In	₅₀Sn	₅₁Sb	₅₂Te	₅₃I	₅₄Xe
Period 6	₅₅Cs	₅₆Ba	₅₇La – ₈₀Hg	₈₁Tl	₈₂Pb	₈₃Bi	₈₄Po	₈₅At	₈₆Rn
Period 7	₈₇Fr	₈₈Ra	₈₉Ac – ₁₁₂Cn	₁₁₃Nh	₁₁₄Fl	₁₁₅Mc	₁₁₆Lv	₁₁₇Ts	₁₁₈Og
f-block Series:		Lanthanides (Period 6): ₅₈Ce ... ₇₁Lu (14 Rare Earth elements) Actinides (Period 7): ₉₀Th ... ₁₀₃Lr (14 Radioactive elements)

Color-coded Long Form Modern Periodic Table showing 18 Groups, 7 Periods, s/p/d/f blocks, and electronic configurations for representative elements.

A. GROUPS (18 Vertical Columns):

Numbered from 1 to 18 according to IUPAC recommendations (replacing old notations IA..VIIA, VIII, IB..VIIB, 0).
Elements in a group have the same number of valence electrons and hence show similar chemical properties.
Group 1: Alkali metals ($\mathrm{Li, Na, K, Rb, Cs, Fr}$) — 1 valence electron.
Group 2: Alkaline earth metals ($\mathrm{Be, Mg, Ca, Sr, Ba, Ra}$) — 2 valence electrons.
Group 13 to 16: Boron, Carbon, Nitrogen, Oxygen families.
Group 17: Halogens ($\mathrm{F, Cl, Br, I, At}$) — 7 valence electrons.
Group 18 (Zero Group): Inert / Noble gases ($\mathrm{He, Ne, Ar, Kr, Xe, Rn}$) — stable octet (duplet for $\mathrm{He}$).

B. PERIODS (7 Horizontal Rows):

Period No.	Type of Period	Number of Elements	Atomic Number Range	Electron Shells Involved	Elements Present
1	Shortest Period	2	Z = 1 to 2	K shell (n = 1)	Hydrogen (H₁) to Helium (He₂)
2	Short Period	8	Z = 3 to 10	L shell (n = 2)	Lithium (Li₃) to Neon (Ne₁₀)
3	Short Period	8	Z = 11 to 18	M shell (n = 3)	Sodium (Na₁₁) to Argon (Ar₁₈)
4	Long Period	18	Z = 19 to 36	N shell (n = 4)	Potassium (K₁₉) to Krypton (Kr₃₆)
5	Long Period	18	Z = 37 to 54	O shell (n = 5)	Rubidium (Rb₃₇) to Xenon (Xe₅₄)
6	Longest Period	32	Z = 55 to 86	P shell (n = 6)	Caesium (Cs₅₅) to Radon (Rn₈₆) Includes Lanthanides (Ce₅₈ → Lu₇₁)
7	Incomplete / Longest	32	Z = 87 to 118	Q shell (n = 7)	Francium (Fr₈₇) to Oganesson (Og₁₁₈) Includes Actinides (Th₉₀ → Lr₁₀₃)

4. Classification & Types of Elements

Based on electronic configurations and shell completion, elements are classified into four main categories:

1. Representative / Normal Elements (s and p-block)

Elements of Group 1, 2, 13, 14, 15, 16, and 17. They have their outermost shell incomplete. They include both metals, metalloids, and non-metals.

2. Transition Elements (d-block)

Elements of Groups 3, 4, 5, 6, 7, 8, 9, 10, 11, and 12. They have their two outermost shells incomplete. All are metals having high melting/boiling points, variable valencies, and forming colored ions/compounds.

3. Inner Transition Elements (f-block)

Elements having their three outermost shells incomplete. Placed separately at the bottom of the periodic table in two horizontal series of 14 elements each:

Lanthanides (Rare Earths): 14 elements from Cerium ($\mathrm{Ce}_{58}$) to Lutetium ($\mathrm{Lu}_{71}$) in Period 6 following Lanthanum ($\mathrm{La}_{57}$).
Actinides (Radioactive Series): 14 elements from Thorium ($\mathrm{Th}_{90}$) to Lawrencium ($\mathrm{Lr}_{103}$) in Period 7 following Actinium ($\mathrm{Ac}_{89}$). All are radioactive.

4. Inert / Noble Gases (Group 18 / Zero Group)

Elements $\mathrm{He_2, Ne_{10}, Ar_{18}, Kr_{36}, Xe_{54}, Rn_{86}}$. They have their outermost shell completely filled ($8$ valence electrons, except Helium which has $2$). Owing to stable octet/duplet, they are unreactive monoatomic gases.

5. Bridge Elements & Diagonal Relationship

Elements of Period 2 show resemblance in properties with elements of the next group in Period 3 diagonally adjacent to them. Such Period 2 elements are called bridge elements.

Fig 5.3: Bridge Elements & Diagonal Relationship (Period 2 & Period 3)

Fig. 5.3: Diagonal Relationship: Period 2 elements (Li, Be, B) exhibit striking similarities in chemical properties with diagonally lower Period 3 elements (Mg, Al, Si).

Lithium ($\mathrm{Li}$) & Magnesium ($\mathrm{Mg}$): Both form nitrides ($\mathrm{Li_3N, Mg_3N_2}$), carbonates decompose on heating.
Beryllium ($\mathrm{Be}$) & Aluminium ($\mathrm{Al}$): Both form amphoteric oxides ($\mathrm{BeO, Al_2O_3}$) and hydroxides.
Boron ($\mathrm{B}$) & Silicon ($\mathrm{Si}$): Both are metalloids forming acidic oxides ($\mathrm{B_2O_3, SiO_2}$).

6. Typical Elements

Elements of Period 3 ($\mathrm{Na_{11}, Mg_{12}, Al_{13}, Si_{14}, P_{15}, S_{16}, Cl_{17}}$) are called typical elements because each summarizes the general characteristics of its respective group.

5. Special Study of Specific Groups

1. Group 1: Alkali Metals ($\mathrm{Li, Na, K, Rb, Cs, Fr}$)

Valence Configuration: 1 valence electron (ns¹). Valency = 1.
Physical Nature: Soft metals (can be cut with a knife), silvery white, low melting and boiling points, light density ($\mathrm{Li, Na, K}$ float on water).
Flame Colors: Impart characteristic colors to flame ($\mathrm{Li}$: Crimson red, $\mathrm{Na}$: Golden yellow, $\mathrm{K}$: Pale violet).
Chemical Properties: Extremely reactive, strong reducing agents, form unipositive cations ($\mathrm{M^+}$).
• Action of Air: $4\mathrm{Na} + \mathrm{O_2} \to 2\mathrm{Na_2O}$ (tarnish quickly).
• Action of Water: React violently with water producing alkali and hydrogen gas: $$2\mathrm{M} + 2\mathrm{H_2O} \to 2\mathrm{MOH} + \mathrm{H_2 \uparrow}$$ • Action of Acids: $2\mathrm{M} + 2\mathrm{HCl} \to 2\mathrm{MCl} + \mathrm{H_2 \uparrow}$

2. Group 2: Alkaline Earth Metals ($\mathrm{Be, Mg, Ca, Sr, Ba, Ra}$)

Valence Configuration: 2 valence electrons (ns²). Valency = 2.
Physical Nature: Greyish white, harder than alkali metals. Oxides occur in earth's crust and form alkaline solutions.
Flame Colors: $\mathrm{Ca}$: Brick red, $\mathrm{Sr}$: Crimson red, $\mathrm{Ba}$: Apple green ($\mathrm{Be}$ and $\mathrm{Mg}$ do not impart color).
Chemical Properties: Divalent electropositive metals ($\mathrm{M^{2+}}$), good reducing agents.
• Action of Water: $\mathrm{Ca} + 2\mathrm{H_2O} \to \mathrm{Ca(OH)_2} + \mathrm{H_2 \uparrow}$ (Magnesium reacts with hot water/steam).
• Action of Acids: $\mathrm{Mg} + 2\mathrm{HCl} \to \mathrm{MgCl_2} + \mathrm{H_2 \uparrow}$

3. Group 17: Halogens ($\mathrm{F, Cl, Br, I, At}$)

Meaning: 'Halo' = sea salt, 'gens' = producer (sea salt producers).
Valence Configuration: 7 valence electrons (ns²np⁵). Valency = 1. Form diatomic non-metallic molecules ($\mathrm{F_2, Cl_2, Br_2, I_2}$).
Physical States & Colors:
• Fluorine ($\mathrm{F_2}$): Pale yellow gas.
• Chlorine ($\mathrm{Cl_2}$): Greenish yellow gas (pungent, poisonous).
• Bromine ($\mathrm{Br_2}$): Reddish brown liquid.
• Iodine ($\mathrm{I_2}$): Dark violet lustrous solid.
Reactivity Trend: Most reactive non-metals! Reactivity decreases down the group: $$\mathbf{\mathrm{F_2 > Cl_2 > Br_2 > I_2}}$$
Action with Metals: Form ionic metal halides, e.g., $2\mathrm{Na} + \mathrm{Cl_2} \to 2\mathrm{NaCl}$.

4. Group 18: Noble / Inert Gases ($\mathrm{He, Ne, Ar, Kr, Xe, Rn}$)

Valence Configuration: Completely filled 8 electrons (ns²np⁶), except He (1s²). Valency = 0.
Properties: Colorless, odorless, tasteless, monoatomic gases. Liquefied with great difficulty.
Uses:
• Helium ($\mathrm{He}$): Filling weather balloons and airships (light and non-inflammable).
• Neon ($\mathrm{Ne}$): Neon sign boards for advertising (glows reddish orange).
• Argon ($\mathrm{Ar}$): Filling electric light bulbs to prevent tungsten filament oxidation.

6. Periodicity & Periodic Trends in Properties

Periodicity: The regular recurrence of elements with similar properties at definite intervals when arranged in increasing order of atomic numbers is called periodicity.

Cause of Periodicity: Recurrence of similar valence shell electronic configurations (same number of outermost electrons) at regular intervals.

1. Trends Down a Group (Top to Bottom)

Fig 5.4: Atomic Radii Variation Down Group 1 & Across Period 3

Fig. 5.4: Atomic Size Variations: (Left) Down Group 1, new electron shells increase atomic radius. (Right) Across Period 3, increasing nuclear charge (+11 to +17) pulls electrons inward, decreasing atomic radius.

Property	Trend Down a Group (↓)	Reason / Mechanism
1. Number of Shells	Increases by 1 at each successive step.	New shell added at each period (e.g., Li: 2,1 → Na: 2,8,1 → K: 2,8,8,1).
2. Valence Electrons	Remains SAME.	Group number equals valence electrons (e.g., all Group 1 elements have 1 valence electron).
3. Valency	Remains SAME.	Depends on valence electrons, which stay constant down the group.
4. Atomic Size (Radius)	Increases down the group.	Addition of a new electron shell at each step increases distance between nucleus and outermost shell, dominating the nuclear charge pull.
5. Metallic Character	Increases down the group.	As atomic size increases, nuclear attraction on valence electrons decreases, making it easier for atom to lose electrons (electropositive).
6. Non-Metallic Character	Decreases down the group.	Tendency to gain electrons decreases due to increased atomic radius and weaker nuclear pull.
7. Chemical Reactivity	• Metals: Increases (Li < Na < K < Rb < Cs) • Non-Metals: Decreases (F > Cl > Br > I)	• Metals: Lower ionization energy down the group increases reactivity. • Non-Metals: Smaller atomic size attracts electrons better, so Fluorine (F) is most reactive.

2. Trends Across a Period (Left to Right $\to$)

Property	Trend Across a Period (→)	Reason / Mechanism
1. Number of Shells	Remains SAME.	All elements in a period have the same number of electron shells (equal to period number).
2. Valence Electrons	Increases from 1 to 8.	Protons and valence electrons increase by 1 at each element across Period.
3. Valency w.r.t. Hydrogen	Increases from 1 to 4, then decreases to 0.	NaH(1), MgH₂(2), AlH₃(3), SiH₄(4), PH₃(3), H₂S(2), HCl(1), Ar(0).
4. Atomic Size (Radius)	Decreases across the period.	Nuclear charge (protons) increases while number of shells remains constant, pulling valence shell closer to nucleus.
5. Metallic Character	Decreases from left to right.	Elements transition from Metals → Metalloids → Non-metals → Inert gases across period.
6. Nature of Oxides	Strongly Basic → Basic → Amphoteric → Acidic → Strongly Acidic.	Period 3 Oxides: • Na₂O (Strong basic), MgO (Basic) • Al₂O₃ (Amphoteric), SiO₂ (Weak acidic) • P₂O₅ (Acidic), SO₃ (Acidic), Cl₂O₇ (Strong acidic)

7. High-Yield ICSE Reasoning & Exam Practice Drills

✍ ICSE EXAM HIGH-YIELD DRILL 1 — Dobereiner's Triad Numerical

Q1. Elements A, B, and C form a Dobereiner's triad. If the atomic mass of A is 7 and that of C is 39, calculate the atomic mass of element B. Name elements A, B, and C.

Calculation: According to Dobereiner's Law of Triads: $$\begin{aligned} \text{Atomic mass of B} &= \frac{\text{Atomic mass of A} + \text{Atomic mass of C}}{2} \\ &= \frac{7 + 39}{2} = \frac{46}{2} = \mathbf{23} \end{aligned}$$

Elements: Element A = Lithium ($\mathrm{Li}$, Mass 7), Element B = Sodium ($\mathrm{Na}$, Mass 23), Element C = Potassium ($\mathrm{K}$, Mass 39).

✍ ICSE EXAM HIGH-YIELD DRILL 2 — Mendeleev's Predictions

Q2. What names were given by Mendeleev to the undiscovered elements corresponding to Gallium and Germanium? What was the basis of his predictions?

Ans: Mendeleev named Gallium as Eka-Aluminium (meaning one place below aluminium) and Germanium as Eka-Silicon (meaning one place below silicon).

The basis of his prediction was the periodic repetition of chemical properties and vacant spaces left in vertical groups of his table.

✍ ICSE EXAM HIGH-YIELD DRILL 3 — Atomic Size Trend Reasoning

Q3. Explain why the atomic radius of Sodium ($\mathrm{Na}$, $Z=11$) is larger than that of Chlorine ($\mathrm{Cl}$, $Z=17$), even though Chlorine has more protons.

Ans: Both Sodium ($\mathrm{Na}: 2,8,1$) and Chlorine ($\mathrm{Cl}: 2,8,7$) belong to Period 3 and have the same number of electron shells (3 shells).

However, Chlorine has 17 protons in its nucleus compared to Sodium's 11 protons. The higher effective nuclear charge in Chlorine exerts a much stronger inward pull on the valence electrons, pulling the outer shell closer and making Chlorine's atomic radius smaller than Sodium's.

✍ ICSE EXAM HIGH-YIELD DRILL 4 — Group Identification

Q4. An element X has atomic number 17. State: (a) Its electronic configuration, (b) Its period number, (c) Its group number (IUPAC), (d) Valency, (e) Whether it is a metal or non-metal.

(a) Electronic Configuration: $2, 8, 7$ (Total 17 electrons).

(b) Period Number: 3 (since electrons occupy 3 shells: K, L, M).

(c) Group Number: Group 17 (Halogens, $7 + 10 = 17$).

(d) Valency: $8 - 7 = \mathbf{1}$.

(e) Nature: Non-metal (Chlorine, forms $\mathrm{Cl^-}$ ion).

8. Master Formula & Rapid Revision Summary

Concept / Law	Key Formula / Definition	ICSE Exam Tip
Dobereiner's Triad Formula	Mass of middle element = (Mass₁ + Mass₃) / 2	Applies to Li-Na-K, Ca-Sr-Ba, Cl-Br-I. Fails for F-Cl-Br.
Mendeleev's Law	Properties = Periodic function of Atomic Mass	Inverted pairs: Ar-K, Co-Ni, Te-I. Left gaps for Eka-Al, Eka-Si.
Modern Periodic Law	Properties = Periodic function of Atomic Number ($Z$)	Discovered by Henry Moseley (1913) via X-ray spectra of metal anodes.
Period Number	Period Number = Total number of electron shells	7 Periods (Period 1 shortest: 2 elements; Period 6 longest: 32 elements).
Group Number (IUPAC)	Group No. = Valence electrons (for valence ≥ 3, add 10)	18 Groups (Group 1 Alkali, Group 2 Alkaline earth, Group 17 Halogens, Group 18 Noble gases).
Atomic Radius Trend	• Down Group: Increases • Across Period: Decreases	Down group: Shell addition dominates. Across period: Nuclear charge pull dominates.
Bridge Elements	Period 2 elements (Li, Be, B) resemble Period 3 (Mg, Al, Si)	Diagonal relationship due to similar ionic radii and polarising power.
Typical Elements	Period 3 elements (Na, Mg, Al, Si, P, S, Cl)	Summarize properties of their respective main groups.