Periodic Table: Complete Chapter in 1 Video || Concepts+PYQs || Class 11 JEE (JEE Wallah)

Summary

This comprehensive video from JEE Wallah provides a complete chapter on the Periodic Table for Class 11 JEE aspirants. It covers fundamental concepts, properties of elements, electronic configurations, and various periodic trends such as atomic radius, ionization energy, electron gain enthalpy, and electronegativity. The video also delves into the modern periodic table's structure, classifications of elements (metals, non-metals, metalloids), lanthanoids and actinoids, and IUPAC nomenclature for elements. Throughout the lecture, concepts are explained with clarity, and Past Year Questions (PYQs) are integrated to aid understanding and exam preparation.

Key Takeaways

* **Modern Periodic Table:** Based on atomic number, arranged in periods (rows) and groups (columns).

* **Periods and Groups:** Elements in the same period have the same number of electron shells; elements in the same group have the same number of valence electrons.

* **Classification of Elements:** Elements are broadly classified into Metals, Non-metals, and Metalloids based on their properties.

* **Lanthanoids and Actinoids:** Inner transition elements located at the bottom of the periodic table.

* **IUPAC Nomenclature:** A systematic naming system for elements, especially those with higher atomic numbers.

* **Electronic Configuration:** Crucial for understanding an element's position and properties in the periodic table.

* **Effective Nuclear Charge ($Z_{eff}$):** The net positive charge experienced by an electron from the nucleus.

* **Penetration Effect:** The ability of an electron to get closer to the nucleus, impacting shielding and $Z_{eff}$.

* **Atomic Radius:** Trends vary across periods and down groups due to changes in electron shells and nuclear attraction.

* **Lanthanide Contraction:** The gradual decrease in atomic and ionic radii across the lanthanide series.

* **Ionic Radius:** Affected by the charge of the ion and the number of electron shells.

* **Ionization Energy (IE):** The energy required to remove an electron from a gaseous atom. Trends are influenced by nuclear charge, shielding, and electron configuration.

* **Electron Gain Enthalpy ($\Delta_{eg}H$) / Electron Affinity (EA):** Energy change when an electron is added to a neutral gaseous atom.

* **Electronegativity (EN):** The tendency of an atom to attract shared electrons in a chemical bond.

* **Oxides:** Properties of oxides vary across the periodic table (acidic, basic, amphoteric).

Detailed Notes

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1. Introduction & Pre-requisites (00:00 - 03:16)

* **Introduction:** Overview of the importance of the Periodic Table in Chemistry for JEE.

* **Topics to be Covered:** Modern Periodic Table, Metals/Non-metals/Metalloids, Atomic Number Distribution, Lanthanoids/Actinoids, IUPAC Names, Electronic Configuration, Effective Nuclear Charge, Penetration Effect, Atomic Radius, Lanthanide Contraction, Ionic Radius, Ionization Energy, Valence Shell Electrons, Electron Gain Enthalpy, Electron Affinity, Electronegativity, Oxides, PYQs.

* **Pre-requisite:** Basic understanding of atomic structure (protons, neutrons, electrons), electronic configuration, and quantum numbers is beneficial.

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2. Modern Periodic Table (06:42 - 22:01)

* **Basis:** Arranged based on **atomic number** (Moseley's Law), not atomic mass.

* **Structure:**

* **Periods (Rows):** 7 periods, correspond to the principal quantum number (n) of the outermost electron shell.

* **Groups (Columns):** 18 groups. Elements in the same group have similar chemical properties due to the same number of valence electrons.

* **Blocks:**

* **s-block:** Groups 1 & 2 (Alkali Metals & Alkaline Earth Metals)

* **p-block:** Groups 13-18 (includes non-metals, metalloids, and some metals)

* **d-block:** Groups 3-12 (Transition Metals)

* **f-block:** Lanthanoids & Actinoids (Inner Transition Metals)

* **General Electronic Configuration:**

* s-block: $ns^1$ or $ns^2$

* p-block: $ns^2np^{1-6}$

* d-block: $(n-1)d^{1-10}ns^{0-2}$

* f-block: $(n-2)f^{1-14}(n-1)d^{0-1}ns^2$

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3. Metals, Non-metals, and Metalloids (22:01 - 31:40)

* **Metals:**

* Tend to lose electrons (electropositive).

* Good conductors of heat and electricity.

* Malleable and ductile.

* Lustrous, usually solid at room temperature (except Hg).

* Located on the left and center of the periodic table.

* **Non-metals:**

* Tend to gain or share electrons (electronegative).

* Poor conductors of heat and electricity (except Graphite).

* Brittle, various states at room temperature.

* Located on the upper right side of the periodic table (e.g., H, C, N, O, P, S, Se, halogens, noble gases).

* **Metalloids (Semimetals):**

* Exhibit properties of both metals and non-metals.

* Located along the zig-zag line separating metals and non-metals.

* Examples: Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), Tellurium (Te), Polonium (Po).

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4. Atomic Number Distribution (31:40 - 43:42)

* **General Trends:**

* Atomic number increases from left to right across a period.

* Atomic number increases from top to bottom down a group.

* **Elements and their Atomic Numbers:** Discussion of specific elements and their positions based on atomic number.

* **Periods and Electron Shells:**

* Period 1: n=1 (H, He)

* Period 2: n=2 (Li to Ne)

* Period 3: n=3 (Na to Ar)

* Period 4: n=4 (K to Kr) - Starts filling 4s, then 3d.

* Period 5: n=5 (Rb to Xe) - Starts filling 5s, then 4d.

* Period 6: n=6 (Cs to Rn) - Starts filling 6s, then 4f (Lanthanoids), then 5d.

* Period 7: n=7 (Fr to Og) - Starts filling 7s, then 5f (Actinoids), then 6d.

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5. Lanthanoid and Actinoid Series (43:42 - 58:37)

* **f-block Elements:** Located at the bottom of the periodic table.

* **Lanthanoids:**

* Start from Cerium (Ce, Z=58) to Lutetium (Lu, Z=71).

* Filling of 4f orbitals.

* First series of inner transition elements.

* Generally exhibit a +3 oxidation state.

* **Actinoids:**

* Start from Thorium (Th, Z=90) to Lawrencium (Lr, Z=103).

* Filling of 5f orbitals.

* Second series of inner transition elements.

* Exhibit a wider range of oxidation states due to the involvement of 5f, 6d, and 7s electrons.

* Most actinoids are radioactive.

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6. IUPAC Name of Elements (58:37 - 1:09:20)

* **Purpose:** To systematically name elements with atomic numbers 100 and above, until official names are approved.

* **Nomenclature Rules:** Based on digits:

* 0: nil (n)

* 1: un (u)

* 2: bi (b)

* 3: tri (t)

* 4: quad (q)

* 5: pent (p)

* 6: hex (h)

* 7: sept (s)

* 8: oct (o)

* 9: enn (e)

* **Suffix:** "-ium"

* **Examples:**

* 101: Un+nil+ium = Unnilium (Unh)

* 102: Un+nil+bium = Unnilbium (Unb)

* 104: Un+nil+quadium = Unnilquadium (Unq)

* 118: Un+un+octium = Ununoctium (Uuo) - Official name: Oganesson (Og)

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7. Electronic Configuration (1:09:20 - 1:43:24)

* **Aufbau Principle:** Electrons fill orbitals starting from the lowest energy levels.

* **Pauli Exclusion Principle:** No two electrons in an atom can have the same set of four quantum numbers.

* **Hund's Rule of Maximum Multiplicity:** Electrons will occupy degenerate orbitals singly before pairing up.

* **Key Concepts:**

* Orbitals and their energy levels (e.g., 1s, 2s, 2p, 3s, 3p, 4s, 3d...).

* Order of filling: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p.

* **Anomalies:**

* Half-filled and fully-filled orbitals are more stable (e.g., Cr, Cu).

* Cr (Z=24): [Ar] 4s¹ 3d⁵ (instead of 4s² 3d⁴)

* Cu (Z=29): [Ar] 4s¹ 3d¹⁰ (instead of 4s² 3d⁹)

* **Noble Gas Configuration:** A shorthand for writing electronic configurations.

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8. Effective Nuclear Charge ($Z_{eff}$) (1:43:24 - 1:52:51)

* **Definition:** The net positive charge experienced by an electron in an atom due to the nucleus.

* **Factors:**

* Actual nuclear charge (Z, atomic number).

* Shielding or screening effect (S).

* **Formula:** $Z_{eff} = Z - S$

* **Shielding Effect (S):** Electrons in inner shells (and to some extent, electrons in the same shell) reduce the attraction of the nucleus for the valence electrons.

* **Trends:**

* **Across a Period:** $Z_{eff}$ increases because Z increases while S increases only slightly (electrons are added to the same shell). This leads to a stronger pull on valence electrons.

* **Down a Group:** $Z_{eff}$ increases slightly because although Z increases significantly, the shielding effect (S) also increases significantly due to the addition of inner electron shells.

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9. Penetration Effect (1:52:51 - 2:04:03)

* **Definition:** The ability of an electron in a particular orbital to get closer to the nucleus.

* **Order of Penetration:** s > p > d > f

* **s-orbitals** have the highest probability of being found close to the nucleus.

* **p-orbitals** have less penetration than s.

* **d-orbitals** have even less penetration.

* **f-orbitals** have the least penetration.

* **Impact on Shielding and $Z_{eff}$:** Electrons in orbitals with higher penetration effect experience greater attraction from the nucleus and are shielded less effectively by other electrons. This contributes to the observed trends in atomic radius and ionization energy.

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10. Atomic Radius (2:04:03 - 2:28:32)

* **Definition:** The distance from the center of the nucleus to the outermost electron shell. Various definitions exist (van der Waals, covalent, metallic radii).

* **Trends:**

* **Across a Period (Left to Right):** Decreases.

* Reason: Nuclear charge (Z) increases, pulling the electron shells closer, while the number of shells (n) remains constant. $Z_{eff}$ increases.

* **Down a Group (Top to Bottom):** Increases.

* Reason: New electron shells are added with each period, increasing the distance of the outermost electrons from the nucleus. Shielding also increases.

* **Special Cases:**

* Noble gases are larger than halogens in the same period due to repulsive forces between their lone pairs and the van der Waals radius definition.

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11. Lanthanide Contraction (2:28:32 - 2:41:31)

* **Definition:** The gradual and steady decrease in atomic and ionic radii across the lanthanide series (elements 57-71).

* **Reason:**

* The f-orbitals are filled across the series.

* Electrons in f-orbitals are poor at shielding each other and the outer electrons.

* As Z increases by 1 for each subsequent element, the nuclear attraction increases, but the shielding effect from the poorly shielding 4f electrons does not fully compensate.

* **Consequences:**

* Elements in the 3rd transition series (after lanthanides) have atomic radii similar to their counterparts in the 2nd transition series (e.g., Zr and Hf, Nb and Ta). This is because the lanthanide contraction effectively "shrinks" the atoms of the 4th period elements, making them comparable in size to the 5th period elements.

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12. Ionic Radius (2:41:31 - 2:50:35)

* **Definition:** The radius of an ion.

* **Trends:**

* **Cations (Positive Ions):** Smaller than their parent atoms.

* Reason: Loss of electrons reduces electron-electron repulsion, and the remaining electrons are pulled more strongly by the nucleus.

* For a given element, $|Z_{eff}|$ increases as positive charge increases (e.g., $Fe^{3+}$ is smaller than $Fe^{2+}$).

* **Anions (Negative Ions):** Larger than their parent atoms.

* Reason: Gain of electrons increases electron-electron repulsion, and the same nuclear charge has to attract more electrons, leading to expansion.

* For a given element, $|Z_{eff}|$ decreases as negative charge increases (e.g., $O^{2-}$ is larger than $O^{-}$).

* **Isoelectronic Species:** Ions with the same number of electrons. Their ionic radii decrease with increasing nuclear charge (e.g., $N^{3-}$, $O^{2-}$, $F^{-}$, $Na^{+}$, $Mg^{2+}$, $Al^{3+}$ all have 10 electrons; radii decrease from $N^{3-}$ to $Al^{3+}$).

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13. Ionisation Energy (IE) / Ionisation Potential (IP) (2:50:35 - 3:30:47)

* **Definition:** The minimum energy required to remove the most loosely bound electron from a neutral gaseous atom in its ground state.

* 1st IE: $X(g) \rightarrow X^+(g) + e^-$

* 2nd IE: $X^+(g) \rightarrow X^{2+}(g) + e^-$ (Always higher than 1st IE)

* **Factors Affecting IE:**

* **Nuclear Charge (Z):** Higher Z $\implies$ Higher IE.

* **Atomic Size:** Larger size $\implies$ Lower IE.

* **Shielding Effect:** Greater shielding $\implies$ Lower IE.

* **Penetration Effect:** Greater penetration $\implies$ Higher IE.

* **Stability of Half-filled/Full-filled Configurations:** Higher stability $\implies$ Higher IE.

* **Trends:**

* **Across a Period (Left to Right):** Generally increases.

* Reason: Z increases, atomic size decreases, $Z_{eff}$ increases.

* **Down a Group (Top to Bottom):** Generally decreases.

* Reason: Atomic size increases, shielding effect increases.

* **Irregularities:**

* **Group 2 (Be) vs Group 1 (Li):** IE of Be > IE of Li (because 2s² is more stable than 2s¹ and Be has higher Z).

* **Group 13 (B) vs Group 2 (Be):** IE of B < IE of Be (because removal of 2p¹ electron from B is easier than removing 2s² electron from Be).

* **Group 15 (N) vs Group 16 (O):** IE of N > IE of O (because N has a stable half-filled 2p³ configuration, and removing an electron from O breaks the pairing in 2p⁴, reducing repulsion).

* **Group 18 (Noble Gases):** Highest IE in their respective periods due to stable full outer shells.

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14. Number of Valence Shell Electrons (3:30:47 - 3:38:04)

* **Definition:** Electrons in the outermost shell, involved in chemical bonding.

* **Group Number and Valence Electrons:**

* For s-block and p-block elements (except He): Group number = Number of valence electrons.

* Groups 1 & 2: 1 or 2 valence electrons.

* Groups 13-18: Group number - 10 = number of valence electrons (e.g., Group 13 has 3, Group 17 has 7).

* For d-block elements: Valence electrons are in the outermost ns and (n-1)d orbitals.

* **Importance:** Determines the chemical properties and bonding behavior of an element.

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15. Electron Gain Enthalpy ($\Delta_{eg}H$) & Electron Affinity (EA) (3:38:04 - 4:28:00)

* **Electron Gain Enthalpy ($\Delta_{eg}H$):** The energy change that occurs when an electron is added to a neutral gaseous atom to form a gaseous anion.

* $\Delta_{eg}H = E.A.$ (if E.A. is defined as energy released)

* If $\Delta_{eg}H$ is negative, energy is released (exothermic, favorable).

* If $\Delta_{eg}H$ is positive, energy is absorbed (endothermic, unfavorable).

* **Electron Affinity (EA):** Often used interchangeably with $\Delta_{eg}H$. Defined as the energy released when an electron is added. A more negative EA means greater affinity.

* **Trends:**

* **Across a Period (Left to Right):** Generally becomes more negative (increases).

* Reason: Z increases, atomic size decreases, $Z_{eff}$ increases, making it easier to attract an electron.

* **Down a Group (Top to Bottom):** Generally becomes less negative (decreases).

* Reason: Atomic size increases, shielding effect increases, making it harder to attract an electron.

* **Irregularities:**

* **Noble Gases (Group 18):** Positive $\Delta_{eg}H$ (energy absorbed) because they already have stable, filled shells.

* **Group 2 (Alkaline Earth Metals):** Positive or slightly negative $\Delta_{eg}H$ because electrons are added to a filled s-orbital (ns²), which is energetically unfavorable.

* **Group 15 (N):** Less negative $\Delta_{eg}H$ than Group 14 (C) due to relatively poor shielding and the addition of an electron to a half-filled p-subshell (p³), which is quite stable.

* **Halogens (Group 17):** Most negative $\Delta_{eg}H$ because they are one electron away from a stable noble gas configuration. Fluorine has a slightly less negative EA than Chlorine due to its smaller size (inter-electron repulsion in the small 2p subshell).

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16. Electronegativity (EN) (4:28:00 - 4:38:24)

* **Definition:** A measure of the tendency of an atom to attract a bonding pair of electrons. It's a relative concept, not an experimentally measured quantity like IE or EA.

* **Pauling Scale:** The most common scale.

* **Factors Affecting EN:** Similar to IE and EA.

* Nuclear Charge

* Atomic Size

* Shielding Effect

* **Trends:**

* **Across a Period (Left to Right):** Increases.

* Reason: Z increases, size decreases, $Z_{eff}$ increases.

* **Down a Group (Top to Bottom):** Decreases.

* Reason: Size increases, shielding increases.

* **General Values:**

* Fluorine (F): Most electronegative element (EN ≈ 3.98 on Pauling scale).

* Cesium (Cs) and Francium (Fr): Least electronegative elements.

* Non-metals have high EN; metals have low EN.

* **Application:** Used to determine bond polarity (ionic, polar covalent, nonpolar covalent). The difference in electronegativity ($\Delta EN$) between two bonded atoms is key.

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17. Oxides (4:38:24 - 4:52:29)

* **Definition:** Compounds formed by an element with oxygen.

* **Acidic Oxides:**

* Formed by non-metals.

* React with water to form acids.

* React with bases to form salt and water.

* Examples: $SO_2$, $CO_2$, $P_4O_{10}$.

* **Basic Oxides:**

* Formed by metals (especially alkali and alkaline earth metals).

* React with water to form bases.

* React with acids to form salt and water.

* Examples: $Na_2O$, $CaO$, $MgO$.

* **Amphoteric Oxides:**

* Exhibit both acidic and basic properties.

* React with both acids and bases.

* Examples: $Al_2O_3$, $ZnO$, $PbO$, $SnO_2$.

* **Neutral Oxides:**

* Do not react with acids or bases.

* Examples: $CO$, $N_2O$, $NO$.

* **Trend:** As we move from left to right across a period, the acidic nature of oxides increases, while the basic nature decreases. Oxides of elements at the beginning of a period are basic, becoming amphoteric in the middle, and acidic towards the end.

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18. Thank You & Closing Remarks (4:52:29 - End)

* Recap of topics covered.

* Encouragement for practice, especially PYQs.

* Information on further resources (DPPs, other courses, PW app).