CBSE Class 11 Chemistry Chapter 3 Notes: Classification of Elements

The Modern Periodic Law and Division of Blocks

The modern organization of elements is based entirely on the Modern Periodic Law, which asserts that the physical and chemical properties of elements are periodic functions of their atomic numbers ($Z$). This fundamental law corrected the discrepancies of Mendeleev's mass-based system by highlighting that electronic configuration controls chemical behaviors.

The modern long-form periodic table organizes the 118 elements into 18 vertical columns (groups) and 7 horizontal rows (periods). These elements are categorized into four distinct blocks ($s$, $p$, $d$, and $f$) depending on the orbital into which the last electron enters:

• s-Block Elements (Groups 1 and 2): Highly reactive metals with outer configuration $ns^1$ or $ns^2$. They exhibit low ionization enthalpies.
• p-Block Elements (Groups 13 to 18): Contains metals, metalloids, and non-metals with configuration $ns^2 np^{1-6}$. Together with s-block elements, they are called Representative Elements.
• d-Block Elements (Groups 3 to 12): Known as Transition Elements, characterized by partially filled $(n-1)d$ subshells and general configuration $(n-1)d^{1-10} ns^{1-2}$.
• f-Block Elements (Lanthanoids and Actinoids): Known as Inner Transition Elements, with valence shell configuration $(n-2)f^{1-14} (n-1)d^{0-1} ns^2$, typically placed separately at the bottom.

Core Revision Glossary

Atomic Radius

One-half the distance between the nuclei of two identical atoms bonded together in a single covalent bond (or metallic bond).

Effective Nuclear Charge ($Z_{eff}$)

The net positive charge experienced by an outer-shell electron from the nucleus, which is less than the actual nuclear charge due to screening by inner electrons.

Ionization Enthalpy ($\Delta_i H$)

The minimum quantitative energy required to remove the most loosely bound electron from an isolated neutral gaseous atom in its ground state.

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

The enthalpy change occurring when an electron is added to a neutral isolated gaseous atom to form a uninegative ion.

Electronegativity

The qualitative measure of the power or ability of an atom in a chemical compound to attract the shared pair of electrons toward itself.

Isoelectronic Species

Atoms or ions that contain the same total number of electrons but differ in nuclear charge (e.g., $N^{3-}$, $O^{2-}$, $F^-$, $Na^+$, $Mg^{2+}$).

Lanthanoid Contraction

The steady, cumulative decrease in the atomic and ionic radii of the transition elements from Lanthanum to Lutetium due to the poor shielding effect of $4f$ electrons.

Key Points & Exception Alert

• Atomic radii decrease across a period due to increasing $Z_{eff}$, and increase down a group because new principal shells are added.
• Cations are always smaller than their parent atoms because of increased effective nuclear charge, whereas anions are always larger due to increased inter-electronic repulsion.
• First ionization enthalpy of Nitrogen ($1s^2 2s^2 2p^3$) is higher than that of Oxygen ($1s^2 2s^2 2p^4$) because of the extra thermodynamic stability of the half-filled $2p$ subshell in Nitrogen.
• First ionization enthalpy of Beryllium ($1s^2 2s^2$) is higher than Boron ($1s^2 2s^2 2p^1$) because a $2s$ electron is closer to the nucleus and requires more energy to remove than a $2p$ electron.
• The electron gain enthalpy of Fluorine is unexpectedly less negative than that of Chlorine because Fluorine's small size leads to severe inter-electronic repulsion in its compact $2p$ subshell.
• Similarly, Oxygen has a less negative electron gain enthalpy than Sulfur due to high electron density in its small $2p$ orbital.
• Noble gases have highly positive electron gain enthalpies because their outermost shells are completely filled, and any added electron must enter a higher-energy shell.
• Metallic character (electropositive nature) decreases across a period and increases down a group, while non-metallic character follows the inverse trend.

Summary of Periodic Trends

Finding the Period, Block, and Group of an Element

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Solved Mini-Examples for Revision

• {"title":"Example 1: Position Determination","description":"An element has atomic number $Z = 23$. Predict its position (Period, Block, and Group) in the periodic table.","solution":"1. Write electronic configuration: $[Ar] 3d^3 4s^2$.\n2. Period = Highest quantum number $n = 4$.\n3. Block = Last electron enters $d$-orbital, so it belongs to the $d$-block.\n4. Group = Number of $(n-1)d$ electrons + $ns$ electrons = $3 + 2 = 5$.\nConclusion: Period 4, Block $d$, Group 5."}
• {"title":"Example 2: Isoelectronic Size Ordering","description":"Arrange the following species in increasing order of their ionic radius: $Mg^{2+}$, $O^{2-}$, $Na^+$, $F^-$.","solution":"All four species are isoelectronic (they all have 10 electrons). For isoelectronic species, the ionic radius decreases as the nuclear charge (number of protons) increases because the nucleus pulls the same number of electrons more tightly.\nProton counts: $O^{2-} (8)$, $F^- (9)$, $Na^+ (11)$, $Mg^{2+} (12)$.\nTherefore, the increasing order of size is: $Mg^{2+} < Na^+ < F^- < O^{2-}$."}

Board Exam Trap Alerts & Cues

1. Be vs B and N vs O Explanations: When answering why Nitrogen has a higher first ionization enthalpy than Oxygen, do not just write 'Nitrogen is stable.' You must write the configuration ($1s^2 2s^2 2p^3$) and explicitly state that the half-filled $p$-orbital has extra exchange energy and symmetric charge distribution.

1. The Halogen Trap: If asked which halogen has the highest electron gain enthalpy, the correct answer is Chlorine, NOT Fluorine. Fluorine's small atomic size leads to high valence electron density and repulsion, limiting its electron-accepting thermodynamic favorability.

1. Ionic Radii of Isoelectronic Species: CBSE frequently asks for ordering. Always count protons ($Z$) first. High protons = small size. No exception here.

Quick Revision Check

• Why does the atomic size decrease across a period? Across a period, electrons are added to the same shell while nuclear charge increases. This increases the effective nuclear charge ($Z_{eff}$), pulling the valence shell closer to the nucleus.
• Which has a lower first ionization enthalpy: Sulfur or Phosphorus? Sulfur has a lower first ionization enthalpy. Phosphorus has a stable, half-filled outer p-subshell ($3p^3$), making it harder to remove an electron compared to Sulfur ($3p^4$).
• What is the general electronic configuration of d-block elements? The general electronic configuration of d-block elements is $(n-1)d^{1-10} ns^{1-2}$.
• Why do noble gases have positive values of electron gain enthalpy? Noble gases have fully filled stable configurations ($ns^2 np^6$). The addition of an extra electron requires placing it in a higher energy shell, which is an endothermic process requiring energy input.

Frequently Asked Questions