Periodic Classification of Elements MCQs
Multiple Choice Questions for NEET and IIT-JEE
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The neutron/proton (N/Z) ratio is a nuclear property that determines nuclear stability and radioactivity. It does not change predictably across a period or down a group in the periodic table, and it even varies among different isotopes of the exact same element.
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The s-block elements consist of Group 1 (alkali metals, historically designated as IA) and Group 2 (alkaline earth metals, historically designated as IIA), as their valence electrons enter the s-subshell.
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The tenth element is Neon (Ne), which is a noble gas located at the end of the second period. Elements belonging to the same period or group share periodic relationships, and neon structurally completes the second period shell ($2s^2 2p^6$).
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While Henry Moseley discovered the atomic number basis, the structural framework of the long form of the periodic table based on electronic configuration was developed primarily by Niels Bohr (along with Rang and Werner).
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The electronic configuration of the element with atomic number 16 (Sulfur) is $1s^2 2s^2 2p^6 3s^2 3p^4$. Having 6 valence electrons, it is placed in the sixth group (Group VIA or Group 16) of the periodic table.
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Representative elements comprise s-block and p-block elements (excluding noble gases under some definitions). Aluminium is a p-block representative element, whereas Chromium is a transition metal, Argon is a noble gas, and Lanthanum is an inner-transition/d-block element.
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Californium (Cf) has an atomic number of 98. It belongs to the 5f-block series, which constitutes the Actinide series in the inner-transition elements.
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As we move from left to right across a period, the nuclear charge increases and the atomic radius decreases. This makes it harder for atoms to lose electrons, resulting in a decrease in metallic character (electropositive nature).
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The electronic configuration of atomic number 20 (Calcium) is $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2$. The highest principal quantum number is $n = 4$, which indicates it belongs to the 4th period.
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The general valence shell electronic configuration of transition elements (d-block elements) is $(n-1)d^{1-10}ns^{0-2}$, representing the progressive filling of the inner d-orbitals.
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Cerium (Ce, atomic number 58) is the first element of the lanthanide series (4f-block), spanning from atomic numbers 58 to 71.
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Both Magnesium (Mg) and Barium (Ba) belong to Group 2 (alkaline earth metals) of the periodic table, possessing two electrons in their outermost s-subshell.
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Sodium (Na, $Z=11$) and Chlorine (Cl, $Z=17$) both have their valence electrons in the third shell ($n=3$), putting them in the 3rd period of the periodic table.
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The last electron enters the 5d orbital ($5d^2$) since the 4f subshell is completely filled ($4f^{14}$). Therefore, this element belongs to the d-block (specifically Hafnium, $Z=72$).
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Lithium (Period 2, Group 1) and Magnesium (Period 3, Group 2) exhibit a diagonal relationship due to their similar ionic sizes and comparable ionic potential (charge/size ratio).
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According to the Modern Periodic Law, the physical and chemical properties of elements are a periodic function of their atomic numbers.
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Group IIIA (Group 13) contains Boron, which is a metalloid, while the rest of the members (Al, Ga, In, Tl) are metals. Group IA includes Hydrogen, but as a pure metallic group designation, Group 13 explicitly mixes metalloids and post-transition metals.
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Inert gases (or noble gases) have a stable, completely filled valence shell configuration of $ns^2 np^6$ (except Helium, which is $1s^2$).
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The element with atomic number 106 (Seaborgium) is located in the 7th period and Group 6. Its outer electronic configuration ends in $6d^4 7s^2$, placing it firmly in the d-block.
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The electronic configuration of atomic number 33 (Arsenic) is $[Ar] 3d^{10} 4s^2 4p^3$. With 5 valence electrons ($4s^2 4p^3$), it belongs to the fifth main group (Group VA or Group 15).
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Atomic number 38 belongs to Strontium (Sr), which is an alkaline earth metal in Group 2. In contrast, 32 (Ge) is a metalloid, 34 (Se) is a non-metal, and 36 (Kr) is a noble gas.
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Moving left to right across a transition series, the atomic volume generally decreases initially due to an increasing effective nuclear charge that draws the electrons closer, before flattening or slightly increasing at the end due to electron-electron repulsion.
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Chalcogens are elements of Group 16 (the Oxygen family). They have 6 electrons in their outermost shell, giving them a valence configuration of $ns^2 np^4$.
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Atomic number 3 is Lithium ($1s^2 2s^1$) and atomic number 12 is Magnesium ($[Ne] 3s^2$). Both have their differentiating valence electrons in an s-subshell, making them s-block elements.
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Going down the halogen group (Group 17), the number of electron shells increases, which leads to a clear and steady increase in both atomic and ionic radii.
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Beryllium (Be) shares a diagonal relationship with Aluminium (Al) because of their similar ionic sizes and high charge densities, yielding similar chemical attributes.
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The configuration 2, 8, 2 belongs to Magnesium ($Z=12$). It has 2 valence electrons which it can easily lose to form a stable cation, a hallmark characteristic of metals.
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Atom A has 5 valence electrons in its outermost shell ($4s^2 4p^3$). It belongs to Group 15 (Pnictogens), making its chemical behaviors highly similar to Nitrogen, which shares the same $ns^2 np^3$ valence configuration.
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Elements within the same group possess the same number of valence (outermost) electrons, which is the underlying cause for their similar chemical properties and bonding behaviors.
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The modern periodic table organizes chemical elements in order of increasing atomic number (number of protons in the nucleus).
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Alkali metals have the largest atomic radii in their respective periods. Due to this large size and low effective nuclear charge on the single valence electron, they possess the lowest ionization potential.
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The shielding (screening) effect order for orbitals in the same shell is $s > p > d > f$. Therefore, d-electrons exhibit a weaker screening effect compared to p-electrons due to their diffused shape.
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Lithium (Li) has an atomic mass of approximately 6.94 u and possesses the lowest density of all solid elements/metals at room temperature.
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Except for Hydrogen (which sits structurally above the s-block), all stable non-metals and metalloids are exclusively clustered in the upper right quadrant within the p-block.
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Lithium (Group 1, Period 2) displays a diagonal relationship with Magnesium (Group 2, Period 3) due to their comparable atomic/ionic size and charge density.
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Atomic number 17 belongs to Chlorine. Its valence shell configuration is $3s^2 3p^5$ (7 valence electrons), meaning it falls into Group VIIA (Group 17) of the periodic table.
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Dmitri Mendeleev is widely credited with the most monumental step in developing the periodic table by organizing elements by atomic mass trends and accurately predicting undiscovered elements.
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While the broader definitions can include Scandium and Yttrium, the explicit inner-transition block series of rare-earth lanthanoids historically counts 14 elements (filling the 4f subshell from Cerium to Lutetium).
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Antimony (Sb) displays physical and chemical traits intermediate between metals and non-metals, classifying it cleanly as a metalloid.
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Atomic number 55 is Cesium (Cs). Following Xenon ($Z=54$), its valence configuration is $[Xe] 6s^1$, which puts it in Group 1 of the s-block.
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Transition metals frequently form coloured ions and salts because they contain partially filled d-orbitals which allow for visible light absorption during d-d electron transitions.
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Arsenic (As) is located along the dividing line between metals and nonmetals in the p-block and functions chemically and physically as a metalloid.
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Cerium ($Z=58$) marks the beginning of the Lanthanide series, where the filling of the 4f subshell commences, identifying it as an f-block element.
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Across a typical period (like period 2), increasing effective nuclear charge pulls electron clouds closer, steadily decreasing the atomic radii and overall atomic/gram volume from Group 1 towards the halogens.
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Lanthanides consist of the 14 elements spanning from atomic numbers 58 to 71, during which the inner 4f subshell is progressively filled.
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Diagonal relationships occur due to competing trends across a period and down a group, which yield a strikingly similar ionic potential or charge/size ratio ($e/r$) between diagonal neighbors.
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Ernest Rutherford is renowned for discovering the atomic nucleus and planetary atom model, but he did not directly work on establishing or designing classification arrangements for the periodic table.
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Tungsten (W) holds the highest melting point among all pure elements in metallic form ($3422^\circ\text{C}$). Note that while diamond is higher, it is an allotrope of carbon, a nonmetal.
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These atomic numbers correspond to Fluorine (9), Chlorine (17), Bromine (35), Iodine (53), and Astatine (85), which form Group 17: the Halogen family.
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When descending a group, the total number of shell levels changes, but the count of valence electrons in the outermost orbit remains completely unchanged.
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Among the isoelectronic series ($\text{Na}^+$, $\text{Mg}^{2+}$, and $\text{Al}^{3+}$), $\text{Al}^{3+}$ has the highest nuclear charge (13 protons acting on 10 electrons), causing the electron cloud to contract the most.
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A neutral atom is always larger than its respective cations. As electrons are progressively removed ($\text{Al} \rightarrow \text{Al}^+ \rightarrow \text{Al}^{2+} \rightarrow \text{Al}^{3+}$), the electron-electron repulsion decreases and the nuclear pull per remaining electron intensifies, shrinking the size.
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Non-metallic character increases across a period from left to right. Oxides of highly non-metallic elements are strongly acidic. Since Sulfur (S) is located further to the right in the third period compared to Mg, Ca, and Si, its oxide ($\text{SO}_2$) acts as a strong acidic oxide when interacting with water to form sulfurous acid.
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Generally, ionization energy increases across a period. However, Nitrogen ($1s^2 2s^2 2p^3$) has a stable, half-filled p-subshell configuration which requires more energy to disrupt than the valence configuration of Oxygen ($1s^2 2s^2 2p^4$). Fluorine remains the highest due to its much smaller atomic radius and higher effective nuclear charge.
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