Classification of Elements and Periodicity in Properties - Class Notes

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Student Quick Revision Focus High-yield exam revision guide for CBSE Boards, NEET, JEE Main & Advanced. Master periodic trends (Radii, Ionization Enthalpy anomalies, Electron Gain Enthalpy halogen trends, Electronegativity, and Nature of Oxides).

1. Historical Development of Periodic Table

1.1 Dobereiner's Law of Triads (1829)

Arranged elements of similar properties into groups of three (Triads). Atomic weight of middle element $\approx$ arithmetic mean of 1st and 3rd.

$\text{Li } (7) + \text{K } (39) \rightarrow \text{Na } (\mathbf{23})$
$\text{Ca } (40) + \text{Ba } (137) \rightarrow \text{Sr } (\mathbf{88.5})$
$\text{Cl } (35.5) + \text{I } (127) \rightarrow \text{Br } (\mathbf{81.25})$

1.2 Newlands' Law of Octaves (1865)

Arranged elements by increasing atomic weight. Every 8th element repeats properties of the 1st (like musical octaves). Valid only up to Calcium ($Z=20$).

1.3 Lothar Meyer's Curves (1869)

Plotted Atomic Volume vs. Atomic Mass. Alkali metals ($\text{Li, Na, K, Rb, Cs}$) occupy peaks, alkaline earth metals occupy descending slopes, halogens occupy ascending slopes.

1.4 Mendeleev's Periodic Table (1869)

Mendeleev's Periodic Law: "Properties of elements are periodic functions of their atomic weights."

High-Yield Exam Focus (Mendeleev's Predictions):

Eka-Aluminium $\rightarrow$ Gallium ($\text{Ga}$)
Eka-Silicon $\rightarrow$ Germanium ($\text{Ge}$)
Eka-Boron $\rightarrow$ Scandium ($\text{Sc}$)

Anomalous Pairs (Atomic Mass Inversions): Plotted $\text{Te} (127.6)$ before $\text{I} (126.9)$, $\text{Co} (58.9)$ before $\text{Ni} (58.7)$, and $\text{Ar} (39.9)$ before $\text{K} (39.1)$.

2. Modern Periodic Law & Present Form

2.1 Henry Moseley's Experiment (1913)

Studied X-ray spectra of metals ($\sqrt{\nu} = a(Z - b)$). Proved that Atomic Number ($Z$), not atomic mass, is the fundamental property governing periodic trends.

Modern Periodic Law "The physical and chemical properties of the elements are periodic functions of their atomic numbers ($Z$)."

2.2 Structure of Long Form Periodic Table

7 Periods (Horizontal Rows): Period number = highest principal quantum number ($n$).
• Period 1 ($n=1$): 2 elements ($1s$) • Period 2 ($n=2$): 8 elements ($2s, 2p$)
• Period 3 ($n=3$): 8 elements ($3s, 3p$) • Period 4 ($n=4$): 18 elements ($4s, 3d, 4p$)
• Period 5 ($n=5$): 18 elements ($5s, 4d, 5p$) • Period 6 ($n=6$): 32 elements ($6s, 4f, 5d, 6p$)
• Period 7 ($n=7$): 32 elements ($7s, 5f, 6d, 7p$)
18 Groups (Vertical Columns): Elements in a group share identical outer electronic configuration and similar chemical behavior.

Long Form Modern Periodic Table of Elements

Fig 2.1: Long Form of the Modern Periodic Table of Elements

3. IUPAC Nomenclature for Elements ($Z > 100$)

Derived from digit roots + suffix 'ium'.

**IUPAC Digit Roots**
Digit	0	1	2	3	4	5	6	7	8	9
Root	nil (n)	un (u)	bi (b)	tri (t)	quad (q)	pent (p)	hex (h)	sept (s)	oct (o)	enn (e)

Examples: $Z=104 \rightarrow \text{Unnilquadium (Unq)}$, $Z=117 \rightarrow \text{Ununseptium (Uus)}$, $Z=120 \rightarrow \text{Unbinilium (Ubn)}$.

4. Block Classification ($s, p, d, f$)

**Block Configurations & Key Characteristics**
Block	General Configuration	Groups	Exam Focus Points
s-Block	$ns^{1-2}$	1, 2	Reactive metals, low IE, form $+1, +2$ cations. Highly electropositive.
p-Block	$ns^2 np^{1-6}$	13 to 18	Includes metals, non-metals, metalloids. Representative elements ($s+p$). Halogens (17), Noble gases (18).
d-Block	$(n-1)d^{1-10} ns^{0-2}$	3 to 12	Transition metals. Variable oxidation states, colored ions, paramagnetic, catalysts. $\mathbf{\text{Zn, Cd, Hg}}$ are not transition elements (completely filled $d^{10}$).
f-Block	$(n-2)f^{1-14} (n-1)d^{0-1} ns^2$	Group 3	Inner transition elements. Lanthanoids ($4f, Z=58-71$) and Actinoids ($5f, Z=90-103$). All actinoids are radioactive.

Fig 4.1: Division of the Periodic Table into s, p, d, and f Blocks

Positional Exceptions Helium ($\text{He}$): $1s^2$ (s-block config), placed in Group 18 p-block due to noble gas inertness.
Hydrogen ($\text{H}$): $1s^1$, placed separately at top due to similarity with both alkali metals and halogens.

5. Periodic Trends in Physical Properties (Core Exam Focus)

5.1 Screening Effect & $Z_{eff}$

$$\mathbf{Z_{eff} = Z - \sigma}$$

Screening power order: $\mathbf{s > p > d > f}$. Across a period, $Z_{eff}$ increases $\Rightarrow$ size decreases.

5.2 Atomic Radii ($r_{vdw} > r_{met} > r_{cov}$)

Across Period: Atomic radius decreases (increasing $Z_{eff}$).
Down Group: Atomic radius increases (addition of principal shells $n$).
Noble Gases: Have non-bonded Van der Waals radii, which are larger than covalent radii of halogens!

5.3 Ionic Radii & Isoelectronic Series (100% Exam Favorite)

$\text{Cation} < \text{Parent Atom}$ ($\text{Na}^+ 95\text{ pm} < \text{Na } 186\text{ pm}$).
$\text{Anion} > \text{Parent Atom}$ ($\text{F}^- 136\text{ pm} > \text{F } 64\text{ pm}$).

Isoelectronic Species Radius Rule For species with same number of electrons ($\text{N}^{3-}, \text{O}^{2-}, \text{F}^-, \text{Na}^+, \text{Mg}^{2+}, \text{Al}^{3+}$):
Higher positive nuclear charge ($Z$) $\Rightarrow$ Smaller radius. $$\text{N}^{3-} > \text{O}^{2-} > \text{F}^- > \text{Na}^+ > \text{Mg}^{2+} > \text{Al}^{3+}$$

Lanthanoid Contraction Exception Poor shielding by $14$ inner $4f$ electrons causes size of 5d series elements to be almost identical to 4d series elements: $$\mathbf{\text{Zr } (4d) \approx \text{Hf } (5d), \quad \text{Nb } \approx \text{Ta}, \quad \text{Mo } \approx \text{W}}$$

5.4 Ionization Enthalpy ($\Delta_i H$) & Exceptions

Energy required to remove most loosely bound electron from isolated gaseous ground state atom. Always endothermic ($\Delta_i H > 0$). $\Delta_i H_1 < \Delta_i H_2 < \Delta_i H_3$.

TOP EXAM EXCEPTIONS — Ionization Enthalpy 1. $\text{Be} > \text{B}$ ($\Delta_i H_1$): $\text{Be} (2s^2)$ has higher 1st IE ($899\text{ kJ/mol}$) than $\text{B} (2p^1, 801\text{ kJ/mol})$ due to stable filled $2s$ subshell. (Same: $\mathbf{\text{Mg} > \text{Al}}$).
2. $\text{N} > \text{O}$ ($\Delta_i H_1$): $\text{N} (2p^3)$ has stable half-filled $2p^3$ configuration ($1402\text{ kJ/mol}$). $\text{O} (2p^4)$ has paired electron repulsion ($1314\text{ kJ/mol}$). (Same: $\mathbf{\text{P} > \text{S}}$).
3. Period 2 IE Order (JEE/NEET): $$\text{Li} < \text{B} < \text{Be} < \text{C} < \text{O} < \text{N} < \text{F} < \text{Ne}$$ 4. Group 13 IE Anomaly: $\text{B} > \text{Tl} > \text{Ga} > \text{Al} > \text{In}$.
5. Deductive IE Jumps: Sudden huge jump between $IE_n$ and $IE_{n+1}$ indicates $n$ valence electrons.

5.5 Electron Gain Enthalpy ($\Delta_{eg} H$)

Enthalpy change when neutral gaseous atom gains an electron: $X_{(g)} + e^- \rightarrow X^-_{(g)}$.

TOP EXAM EXCEPTIONS — Electron Gain Enthalpy 1. Halogen Trend ($\mathbf{\text{Cl} > \text{F}}$): Chlorine has the most negative $\Delta_{eg} H$ in periodic table! $$\text{Cl } (-349\text{ kJ/mol}) > \text{F } (-328) > \text{Br } (-325) > \text{I } (-295)$$ Reason: Fluorine is extremely small; incoming electron suffers high inter-electronic repulsion in compact $2p$ subshell.

2. Group 16 Trend ($\mathbf{\text{S} > \text{O}}$): Sulfur ($-200\text{ kJ/mol}$) is more negative than Oxygen ($-141\text{ kJ/mol}$).

3. 2nd Electron Gain Enthalpy ($\Delta_{eg} H_2$): ALWAYS POSITIVE / ENDOTHERMIC for all elements ($\text{O}^- + e^- \rightarrow \text{O}^{2-}, \Delta_{eg}H_2 = +780\text{ kJ/mol}$) due to anion-electron repulsion.

5.6 Electronegativity (EN)

Tendency of a bonded atom to attract shared pair of electrons.

Key Electronegativity Formulas & Scale Values:
• Pauling Scale: $\text{F} (4.0) > \text{O} (3.5) > \text{N} (3.0) \approx \text{Cl} (3.0) > \text{Br} (2.8) > \text{I} (2.5) \approx \text{S} (2.5) \approx \text{C} (2.5) > \text{H} (2.1)$
• Mulliken Scale: $\text{EN}_{\text{Mulliken}} = \frac{\text{IE} + \text{EA}}{2} \quad (\text{in eV}) \quad \Rightarrow \text{EN}_{\text{Pauling}} \approx \frac{\text{EN}_{\text{Mulliken}}}{2.8}$
• Hybridization Effect: $\text{EN} \propto \% s\text{-character} \Rightarrow sp (50\%) > sp^2 (33.3\%) > sp^3 (25\%)$

6. Periodic Trends in Chemical Properties

6.1 Valence & Oxidation States

Valence = Number of outer electrons (Group 1, 2, 13, 14) or $(8 - \text{outer electrons})$ (Group 15 to 18).

Example ($\text{OF}_2$ vs $\text{Na}_2\text{O}$): In $\text{OF}_2$, $\text{F} = -1$ and $\text{O} = +2$. In $\text{Na}_2\text{O}$, $\text{O} = -2$ and $\text{Na} = +1$.

6.2 Anomalous Properties of 2nd Period & Diagonal Relationship

2nd period elements ($\text{Li, Be, B, C, N, O, F}$) differ from group members due to small size, high IE/EN, absence of vacant $d$-orbitals (max covalency = 4), and $p\pi-p\pi$ multiple bonding ability.

Diagonal Relationship: Similarities between $\text{Li}-\text{Mg}, \text{Be}-\text{Al}, \text{B}-\text{Si}$ due to nearly equal polarizing power ($\text{Ionic Charge} / \text{Radius}^2$).

6.3 Oxides Nature Classification (Direct Exam MCQs)

Basic Oxides: $\text{Na}_2\text{O}, \text{K}_2\text{O}, \text{CaO}, \text{BaO}$ ($\text{Na}_2\text{O} + \text{H}_2\text{O} \rightarrow 2\text{NaOH}$).
Acidic Oxides: $\text{Cl}_2\text{O}_7, \text{SO}_3, \text{N}_2\text{O}_5, \text{CO}_2, \text{P}_4\text{O}_{10}$ ($\text{Cl}_2\text{O}_7 + \text{H}_2\text{O} \rightarrow 2\text{HClO}_4$).
Amphoteric Oxides (High-Yield): $\mathbf{\text{Al}_2\text{O}_3, \text{Ga}_2\text{O}_3, \text{ZnO}, \text{BeO}, \text{PbO}, \text{SnO}, \text{As}_2\text{O}_3, \text{V}_2\text{O}_5, \text{Cr}_2\text{O}_3}$.
Neutral Oxides (High-Yield): $\mathbf{\text{CO}, \text{NO}, \text{N}_2\text{O}, \text{H}_2\text{O}}$.

6.4 Chemical Reactivity & Oxidizing Power

Group 1 Reactivity: Increases down group ($\text{Li} < \text{Na} < \text{K} < \text{Rb} < \text{Cs}$).
Group 17 Reactivity: Decreases down group ($\text{F}_2 > \text{Cl}_2 > \text{Br}_2 > \text{I}_2$).
Oxidizing Power Order: $\mathbf{\text{F} > \text{O} > \text{Cl} > \text{N}}$.

Student High-Yield Quick Self-Test Q1 (CBSE Board / NEET): Which element has the most negative electron gain enthalpy in the periodic table?

Solution: Chlorine ($\text{Cl}$) ($-349\text{ kJ/mol}$) because Fluorine is extremely small and experiences heavy $2p$ inter-electronic repulsion.