Grade 11 : Chemistry - Periodic Table (Set 1)

Explanation: Moseley demonstrated that atomic number ($Z$) is a more fundamental property of an element than its atomic mass. Hence, modern periodic law states that properties depend on atomic number.

Explanation: Elements in which the last electron enters the s-orbital of the outermost shell are s-block elements. Their general outer configuration is $ns^1$ (Group 1) or $ns^2$ (Group 2).

Explanation: Fluorine is the most electronegative element in the periodic table with an assigned Pauling value of 4.0 due to its small size and high nuclear attraction.

Explanation: Transition or d-block elements fill the penultimate $(n-1)d$ subshell while the outermost shell has configuration $ns^{1-2}$.

Explanation: Moving across a period, effective nuclear charge ($Z_{eff}$) increases while electrons enter the same valence shell, pulling electrons closer to the nucleus and decreasing atomic radius.

Explanation: Metallic character depends on electron-donating ability. Down a group, atomic size increases and ionization enthalpy decreases, making it easier to lose electrons.

Explanation: Argon ($ ext{Ar}$, $Z=18$) belongs to Group 18, possessing a completely filled valence octet ($3s^2 3p^6$), making it a noble/inert gas.

Explanation: Atomic number 17 is Chlorine with configuration $1s^2 2s^2 2p^6 3s^2 3p^5$. Highest principal quantum number $n=3$ (Period 3) and group is $10 + 7 = 17$.

Explanation: Down a group, atomic radius increases and inner-shell shielding increases, lowering electrostatic attraction on outer electrons, thus reducing first ionization enthalpy.

Explanation: For isoelectronic species (10 electrons each), ionic radius decreases as nuclear charge ($Z$) increases. $\text{Al}^{3+}$ has $Z=13$, which exerts the strongest attraction on electrons.

Explanation: Elements in the same group have identical valence shell electronic configurations (same number of valence electrons), leading to similar chemical properties.

Explanation: f-block elements (lanthanoids and actinoids) are called inner-transition elements because they fill the $(n-2)f$ subshell within the transition series.

Explanation: Nitrogen, Oxygen, and Fluorine have their last valence electrons entering $2p$ subshells, placing them in the p-block (Groups 15, 16, and 17).

Explanation: Halogens require only one electron to attain stable noble gas configuration ($ns^2 np^6$), releasing high energy when acquiring an electron ($\Delta_{eg}H$ is highly negative).

Explanation: Alkaline earth metals ($ ext{Be}, ext{Mg}, ext{Ca}, ext{Sr}, ext{Ba}, ext{Ra}$) belong to Group 2 and have outer configuration $ns^2$.

Explanation: Chlorine has a more negative electron gain enthalpy than Fluorine because Fluorine's tiny $2p$ subshell suffers significant inter-electronic repulsion upon adding an extra electron.

Explanation: Element 104 is named Rutherfordium ($ ext{Rf}$) in honor of Ernest Rutherford. Systematic IUPAC name was Unnilquadium.

Explanation: $\text{Al}_2\text{O}_3$ reacts with both acids and bases to form salts and water, making it an amphoteric oxide.

Explanation: Nuclear charge $Z$ increases by 1 unit for each step left-to-right while shielding constant $\sigma$ increases minimally, resulting in net increase of $Z_{eff} = Z - \sigma$.

Explanation: Group 16 elements ($ ext{O}, ext{S}, ext{Se}, ext{Te}, ext{Po}$) are known as chalcogens, meaning 'ore-forming' elements.