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101. How many elements are present in the second period of the periodic table?
ⓐ. 2
ⓑ. 8
ⓒ. 18
ⓓ. 32
Correct Answer: 8
Explanation: The second period begins with lithium (Z = 3) and ends with neon (Z = 10). It corresponds to the filling of the 2s and 2p orbitals, which can accommodate a total of 8 electrons, hence 8 elements.
102. Which period is considered the shortest in the periodic table?
ⓐ. 1st period
ⓑ. 2nd period
ⓒ. 3rd period
ⓓ. 7th period
Correct Answer: 1st period
Explanation: The first period consists of only 2 elements (hydrogen and helium) because only the 1s orbital is available for electrons. It is therefore the shortest period in the table.
103. How many elements are present in the 4th period of the periodic table?
ⓐ. 8
ⓑ. 18
ⓒ. 32
ⓓ. 16
Correct Answer: 18
Explanation: The 4th period includes the filling of 4s, 3d, and 4p orbitals. This makes a total of 18 elements ranging from potassium (Z = 19) to krypton (Z = 36).
104. Which groups are known as the representative or main group elements?
ⓐ. Group 1 and 2 only
ⓑ. Group 13–18 only
ⓒ. Group 1, 2, and 13–18
ⓓ. Group 3–12
Correct Answer: Group 1, 2, and 13–18
Explanation: The s-block (Groups 1–2) and p-block (Groups 13–18) are collectively called the main group or representative elements. They display wide variations in physical and chemical properties and include metals, non-metals, and metalloids.
105. Which groups are called transition elements?
ⓐ. Groups 3–12
ⓑ. Groups 13–18
ⓒ. Groups 1–2
ⓓ. Groups 17–18
Correct Answer: Groups 3–12
Explanation: Elements in Groups 3–12 are d-block elements where the last electron enters a d-orbital. They are known as transition elements, showing variable oxidation states, colored compounds, and catalytic activity.
106. Which group contains the alkaline earth metals?
ⓐ. Group 1
ⓑ. Group 2
ⓒ. Group 13
ⓓ. Group 14
Correct Answer: Group 2
Explanation: Group 2 elements (Be, Mg, Ca, Sr, Ba, Ra) are alkaline earth metals. They form divalent cations (M²⁺), have valency 2, and their oxides and hydroxides are basic in nature, hence the name “alkaline earth.”
107. Which group is also known as the chalcogens or oxygen family?
ⓐ. Group 14
ⓑ. Group 15
ⓒ. Group 16
ⓓ. Group 17
Correct Answer: Group 16
Explanation: Group 16 elements (O, S, Se, Te, Po, Lv) are called chalcogens. The word “chalcogen” means “ore-formers” because many ores are oxides and sulfides. They have 6 valence electrons (ns²np⁴).
108. Which period contains the actinides?
ⓐ. 5th period
ⓑ. 6th period
ⓒ. 7th period
ⓓ. 8th period
Correct Answer: 7th period
Explanation: The 7th period contains actinides (Z = 89 to 103). They belong to the f-block and are mostly radioactive. Like lanthanides, they are shown separately at the bottom of the periodic table.
109. Which group contains elements that exist as diatomic non-metal molecules under normal conditions?
ⓐ. Group 1
ⓑ. Group 14
ⓒ. Group 17
ⓓ. Group 18
Correct Answer: Group 17
Explanation: Group 17 elements (halogens) include F₂, Cl₂, Br₂, and I₂ which exist as diatomic molecules. They are highly reactive non-metals and form halide salts with metals.
110. What is common to all elements in a particular period?
ⓐ. Same valence electron configuration
ⓑ. Same number of shells
ⓒ. Same chemical properties
ⓓ. Same group number
Correct Answer: Same number of shells
Explanation: In a given period, elements have the same principal quantum number (n), i.e., the same number of electron shells. For example, all elements in the 3rd period (Na to Ar) have 3 shells. However, chemical properties vary across the period due to change in valence electrons.
111. Which of the following is a general property of metals?
ⓐ. Poor conductors of electricity
ⓑ. High malleability and ductility
ⓒ. Form acidic oxides
ⓓ. Exist mainly as gases
Correct Answer: High malleability and ductility
Explanation: Metals are typically malleable (can be hammered into sheets) and ductile (drawn into wires). They are good conductors of heat and electricity, have high melting points, and form basic oxides. Non-metals, by contrast, are brittle and form acidic oxides.
112. Which element is considered the lightest metal?
ⓐ. Sodium
ⓑ. Potassium
ⓒ. Lithium
ⓓ. Magnesium
Correct Answer: Lithium
Explanation: Lithium (Z = 3) is the lightest metal and belongs to Group 1 (alkali metals). It is highly reactive, forms Li⁺ ions, and reacts with water less vigorously than sodium and potassium due to smaller atomic size.
113. Which of the following is NOT a property of non-metals?
ⓐ. Poor conductors of heat and electricity
ⓑ. Usually form anions in ionic compounds
ⓒ. Ductile and malleable
ⓓ. Form acidic oxides
Correct Answer: Ductile and malleable
Explanation: Non-metals are brittle and cannot be hammered or drawn into wires. They generally form anions (e.g., O²⁻, Cl⁻), are poor conductors, and often form acidic oxides like SO₂ and CO₂. Metals, not non-metals, are malleable and ductile.
114. Which of the following is a metalloid?
ⓐ. Sodium
ⓑ. Silicon
ⓒ. Sulphur
ⓓ. Neon
Correct Answer: Silicon
Explanation: Metalloids have properties of both metals and non-metals. Silicon is shiny like a metal but brittle like a non-metal. It conducts electricity moderately (semiconductor) and forms covalent compounds. Other metalloids include B, As, Sb, Te.
115. Which property is common among all metals?
ⓐ. High ionization enthalpy
ⓑ. Tendency to lose electrons and form cations
ⓒ. Tendency to form acidic oxides
ⓓ. Low thermal conductivity
Correct Answer: Tendency to lose electrons and form cations
Explanation: Metals are electropositive in nature; they lose valence electrons easily due to low ionization enthalpy, forming cations (e.g., Na⁺, Ca²⁺). This explains their metallic bonding and basic character of oxides.
116. Which element is an exception, often showing both metallic and non-metallic behavior?
ⓐ. Hydrogen
ⓑ. Helium
ⓒ. Chlorine
ⓓ. Iron
Correct Answer: Hydrogen
Explanation: Hydrogen has a single electron and can behave like an alkali metal (forming H⁺) or like a halogen (forming H⁻). Because of this dual nature, its classification as purely metal or non-metal is debated.
117. Which group contains the most reactive non-metals?
ⓐ. Group 1
ⓑ. Group 17
ⓒ. Group 18
ⓓ. Group 2
Correct Answer: Group 17
Explanation: Group 17 elements (halogens) are the most reactive non-metals. They require only one electron to complete their valence shell, making them highly electronegative and reactive, especially fluorine.
118. Which of the following elements is NOT a metalloid?
ⓐ. Arsenic
ⓑ. Tellurium
ⓒ. Boron
ⓓ. Calcium
Correct Answer: Calcium
Explanation: Calcium is an alkaline earth metal, not a metalloid. Metalloids include boron, silicon, arsenic, antimony, tellurium, and polonium, which have properties intermediate between metals and non-metals.
119. Which of the following best describes metalloids?
ⓐ. They are inert gases
ⓑ. They always conduct like metals
ⓒ. They show properties of both metals and non-metals
ⓓ. They only form basic oxides
Correct Answer: They show properties of both metals and non-metals
Explanation: Metalloids can behave as metals in some reactions and as non-metals in others. For example, boron forms covalent compounds (non-metallic property) but also exhibits metallic luster. Silicon conducts electricity moderately (semiconductor).
120. Which of the following pairs represents a non-metal and a metalloid, respectively?
ⓐ. Oxygen and Silicon
ⓑ. Sodium and Boron
ⓒ. Magnesium and Sulphur
ⓓ. Argon and Tellurium
Correct Answer: Oxygen and Silicon
Explanation: Oxygen is a non-metal with high electronegativity and acidic oxides. Silicon is a metalloid, showing mixed properties such as semiconductor behavior and formation of covalent compounds. This pair demonstrates the distinction between non-metals and metalloids.
121. In the modern periodic table, where are noble gases placed?
ⓐ. Group 1
ⓑ. Group 17
ⓒ. Group 18
ⓓ. Group 2
Correct Answer: Group 18
Explanation: Noble gases (He, Ne, Ar, Kr, Xe, Rn, Og) are placed in Group 18, the last column of the periodic table. They have a complete valence shell (ns²np⁶, except He with 1s²), which makes them chemically inert under normal conditions.
122. Why are noble gases called “inert gases”?
ⓐ. Because they have incomplete octets
ⓑ. Because they have low atomic masses
ⓒ. Because they have completely filled valence shells
ⓓ. Because they do not exist naturally
Correct Answer: Because they have completely filled valence shells
Explanation: Noble gases have stable electronic configurations, either duplet (He: 1s²) or octet (others: ns²np⁶). This stability prevents them from reacting easily with other elements, hence the name “inert gases.”
123. What is the position of hydrogen in the periodic table?
ⓐ. Only in Group 1
ⓑ. Only in Group 17
ⓒ. Can be placed in both Group 1 and Group 17
ⓓ. Group 18
Correct Answer: Can be placed in both Group 1 and Group 17
Explanation: Hydrogen has a dual nature. Like alkali metals (Group 1), it has one electron in its valence shell (1s¹). Like halogens (Group 17), it forms diatomic molecules (H₂) and can gain one electron to form H⁻. Hence, its position is unique and debated.
124. Which gas is essential for combustion and respiration?
ⓐ. Hydrogen
ⓑ. Oxygen
ⓒ. Nitrogen
ⓓ. Argon
Correct Answer: Oxygen
Explanation: Oxygen (Group 16) is essential for life and combustion. It readily forms oxides with most elements. Its high electronegativity and ability to accept electrons make it one of the most reactive non-metals.
125. Which gas makes up about 78% of the Earth’s atmosphere?
ⓐ. Oxygen
ⓑ. Hydrogen
ⓒ. Nitrogen
ⓓ. Argon
Correct Answer: Nitrogen
Explanation: Nitrogen is a Group 15 element, forming a diatomic molecule (N₂). It is inert due to a strong triple bond, making up about 78% of Earth’s atmosphere. It plays a vital role in biological nitrogen fixation and the nitrogen cycle.
126. What is the most abundant element in the universe?
ⓐ. Hydrogen
ⓑ. Oxygen
ⓒ. Nitrogen
ⓓ. Helium
Correct Answer: Hydrogen
Explanation: Hydrogen makes up about 75% of the universe’s elemental mass. It is the lightest element (Z = 1), highly flammable, and fuels stars through nuclear fusion. Despite being in Group 1, its unique properties make its classification complex.
127. Which gas is placed in Group 18 but has an electronic configuration similar to Group 2?
ⓐ. Argon
ⓑ. Helium
ⓒ. Neon
ⓓ. Krypton
Correct Answer: Helium
Explanation: Helium has an electronic configuration of 1s², like an s-block element. However, it is placed in Group 18 due to its chemical inertness and similarity with noble gases. Thus, its position is exceptional in the periodic table.
128. Which gas forms about 21% of Earth’s atmosphere?
ⓐ. Oxygen
ⓑ. Hydrogen
ⓒ. Nitrogen
ⓓ. Carbon dioxide
Correct Answer: Oxygen
Explanation: Oxygen constitutes about 21% of Earth’s atmosphere. It is essential for respiration and combustion. Along with nitrogen, it forms the bulk of atmospheric gases.
129. Why is hydrogen sometimes placed in Group 1?
ⓐ. Because it is a noble gas
ⓑ. Because it has one electron like alkali metals
ⓒ. Because it is a metalloid
ⓓ. Because it is inert
Correct Answer: Because it has one electron like alkali metals
Explanation: Hydrogen has one electron (1s¹), similar to alkali metals (ns¹). It forms +1 ions (H⁺) in many compounds, just like Group 1 metals. This similarity justifies its placement in Group 1.
130. Why is the position of hydrogen unique in the periodic table?
ⓐ. Because it is radioactive
ⓑ. Because it shows properties similar to alkali metals and halogens
ⓒ. Because it belongs to transition elements
ⓓ. Because it is heavier than helium
Correct Answer: Because it shows properties similar to alkali metals and halogens
Explanation: Hydrogen forms H⁺ ions like alkali metals and H⁻ ions like halogens. It also exists as a diatomic molecule (H₂). Due to this dual nature, hydrogen cannot be placed definitively in one group, making its position unique in the periodic table.
131. Why was the IUPAC system of naming elements beyond atomic number 100 introduced?
ⓐ. Because atomic numbers became difficult to calculate
ⓑ. Because many elements were discovered simultaneously by different groups
ⓒ. Because symbols of known elements started repeating
ⓓ. Because isotopes required new names
Correct Answer: Because many elements were discovered simultaneously by different groups
Explanation: For transuranium elements (Z > 100), discoveries were often contested between laboratories. To avoid confusion, IUPAC (International Union of Pure and Applied Chemistry) introduced a systematic method in 1978 using Latin/Greek roots for atomic numbers until a permanent name was approved.
132. What does the IUPAC name *Unnilquadium (Unq)* represent?
ⓐ. Element with Z = 102
ⓑ. Element with Z = 104
ⓒ. Element with Z = 114
ⓓ. Element with Z = 116
Correct Answer: Element with Z = 104
Explanation: *Unnilquadium* (Unq) was the temporary IUPAC name for the element with atomic number 104. Its final approved name is Rutherfordium (Rf). “Un” (1), “nil” (0), and “quad” (4) indicate the digits 104.
133. Which element was temporarily named *Ununbium (Uub)*?
ⓐ. Copernicium (Z = 112)
ⓑ. Flerovium (Z = 114)
ⓒ. Darmstadtium (Z = 110)
ⓓ. Nihonium (Z = 113)
Correct Answer: Copernicium (Z = 112)
Explanation: The IUPAC temporary name *Ununbium* (Uub) corresponded to element 112. Later, it was officially named Copernicium (Cn) in honor of Nicolaus Copernicus. The prefix *Un-un-bi* stands for digits 1-1-2.
134. According to IUPAC rules, what is the root word for digit 9 in naming superheavy elements?
ⓐ. Enn
ⓑ. Nill
ⓒ. Enne
ⓓ. Ennne
Correct Answer: Enn
Explanation: The IUPAC root words for digits are: 0 = nil, 1 = un, 2 = bi, 3 = tri, 4 = quad, 5 = pent, 6 = hex, 7 = sept, 8 = oct, 9 = enn. These are combined sequentially to denote the atomic number of the element.
135. What is the temporary IUPAC name for element with atomic number 118?
ⓐ. Ununoctium
ⓑ. Unnilhexium
ⓒ. Unnilennium
ⓓ. Unniloctium
Correct Answer: Ununoctium
Explanation: Before official naming, element 118 was called *Ununoctium* (Uuo), derived from roots un (1), un (1), oct (8). In 2016, it was officially named Oganesson (Og) after Yuri Oganessian, a nuclear physicist.
136. Which of the following is the official name of element 113?
ⓐ. Ununtrium
ⓑ. Nihonium
ⓒ. Moscovium
ⓓ. Tennessine
Correct Answer: Nihonium
Explanation: Element 113 was first given the temporary name *Ununtrium (Uut)*, then officially named Nihonium (Nh) in 2016. The name was chosen to honor Japan (*Nihon* = Japan), where the element was discovered.
137. What is the symbol and name for element 115?
ⓐ. Uup – Moscovium
ⓑ. Uuh – Livermorium
ⓒ. Uuq – Rutherfordium
ⓓ. Uuo – Oganesson
Correct Answer: Uup – Moscovium
Explanation: The temporary name *Ununpentium* (Uup) corresponded to element 115. It was later named Moscovium (Mc) after Moscow, Russia. Its properties are similar to other Group 15 elements.
138. Which element was once called *Ununseptium (Uus)*?
ⓐ. Flerovium (114)
ⓑ. Moscovium (115)
ⓒ. Tennessine (117)
ⓓ. Livermorium (116)
Correct Answer: Tennessine (117)
Explanation: The temporary name *Ununseptium (Uus)* denoted element 117. In 2016, it was officially named Tennessine (Ts) after Tennessee, USA, recognizing contributions of Oak Ridge National Laboratory.
139. What is the general rule for writing symbols in IUPAC temporary names?
ⓐ. Three capital letters are used
ⓑ. One letter symbols are always preferred
ⓒ. Derived from first letters of each root, capitalizing the first
ⓓ. Random letters are used until a name is chosen
Correct Answer: Derived from first letters of each root, capitalizing the first
Explanation: For temporary symbols, IUPAC uses the first letter of each root word. For example, element 112: *Ununbium* → Uub. The first letter is capitalized, while others are lowercase.
140. What was the official name given to element 116, earlier called *Ununhexium (Uuh)*?
ⓐ. Flerovium
ⓑ. Livermorium
ⓒ. Nihonium
ⓓ. Copernicium
Correct Answer: Livermorium
Explanation: Element 116 was temporarily named *Ununhexium (Uuh)* and officially renamed Livermorium (Lv) in 2012, after the Lawrence Livermore National Laboratory in California, which contributed to its discovery.
141. According to IUPAC nomenclature, which root word represents the digit 0?
ⓐ. Nil
ⓑ. Un
ⓒ. Oct
ⓓ. Enn
Correct Answer: Nil
Explanation: In the IUPAC system, each digit of an atomic number has a root word. For 0, the root is *nil*. For example, element 104 was called *Unnilquadium* (1-0-4), showing the use of “nil” for zero.
142. Which digit is represented by the root word “un” in IUPAC naming?
ⓐ. 0
ⓑ. 1
ⓒ. 2
ⓓ. 9
Correct Answer: 1
Explanation: The root “un” stands for digit 1 in the IUPAC system. For example, element 111 was named *Unununium* before receiving its official name, Roentgenium (Rg).
143. What is the root word for digit 2 in the IUPAC naming system?
ⓐ. Bi
ⓑ. Tri
ⓒ. Nil
ⓓ. Hex
Correct Answer: Bi
Explanation: The digit 2 is represented by the root word *bi*. For example, element 112 was initially called *Ununbium* (1-1-2).
144. Which digit corresponds to the IUPAC root “tri”?
ⓐ. 2
ⓑ. 3
ⓒ. 4
ⓓ. 7
Correct Answer: 3
Explanation: The digit 3 is denoted by the root word *tri*. For instance, element 113 was given the temporary name *Ununtrium* before being named Nihonium (Nh).
145. What is the root for digit 4 in the IUPAC system?
ⓐ. Quad
ⓑ. Pent
ⓒ. Sept
ⓓ. Oct
Correct Answer: Quad
Explanation: The digit 4 corresponds to the root *quad*. Element 104 was originally called *Unnilquadium* (Unq), later renamed Rutherfordium (Rf).
146. Which digit is represented by the root “pent”?
ⓐ. 4
ⓑ. 5
ⓒ. 6
ⓓ. 7
Correct Answer: 5
Explanation: The digit 5 is represented by *pent*. For example, element 115 was called *Ununpentium* (Uup) before being officially named Moscovium (Mc).
147. What is the root for digit 6 in IUPAC element naming?
ⓐ. Hex
ⓑ. Sept
ⓒ. Oct
ⓓ. Enn
Correct Answer: Hex
Explanation: The digit 6 is denoted by the root *hex*. Element 106 was called *Unnilhexium* (Uuh), later renamed Seaborgium (Sg) in honor of Glenn T. Seaborg.
148. Which digit corresponds to the IUPAC root “sept”?
ⓐ. 6
ⓑ. 7
ⓒ. 8
ⓓ. 9
Correct Answer: 7
Explanation: The digit 7 corresponds to *sept*. Element 107 was initially named *Unnilseptium* (Uns), before receiving the official name Bohrium (Bh).
149. What is the root for digit 8 in IUPAC naming rules?
ⓐ. Oct
ⓑ. Enn
ⓒ. Hex
ⓓ. Quad
Correct Answer: Oct
Explanation: The digit 8 corresponds to the root *oct*. For example, element 108 was *Unniloctium* (Uno), later named Hassium (Hs).
150. Which digit is represented by the root “enn”?
ⓐ. 7
ⓑ. 8
ⓒ. 9
ⓓ. 0
Correct Answer: 9
Explanation: The digit 9 is represented by *enn*. For instance, element 109 was temporarily called *Unnilennium* (Une), later renamed Meitnerium (Mt).
151. Which is the heaviest element currently recognized in the periodic table?
ⓐ. Livermorium (Lv, 116)
ⓑ. Moscovium (Mc, 115)
ⓒ. Oganesson (Og, 118)
ⓓ. Tennessine (Ts, 117)
Correct Answer: Oganesson (Og, 118)
Explanation: Oganesson (Z = 118) is the heaviest element officially recognized by IUPAC. It is a noble gas but is predicted to have unique properties due to relativistic effects. It was named in honor of physicist Yuri Oganessian.
152. Which superheavy element is named after a Russian region?
ⓐ. Nihonium
ⓑ. Moscovium
ⓒ. Flerovium
ⓓ. Seaborgium
Correct Answer: Moscovium
Explanation: Element 115 was named *Moscovium (Mc)* after the Moscow region in Russia. It was discovered by collaboration between Russian and American scientists in 2003.
153. Which superheavy element is named after Japan?
ⓐ. Nihonium
ⓑ. Tennessine
ⓒ. Copernicium
ⓓ. Meitnerium
Correct Answer: Nihonium
Explanation: Element 113 was named *Nihonium (Nh)*, derived from “Nihon,” which means Japan in Japanese. It was discovered at RIKEN, Japan, making it the first element discovered in Asia.
154. Which element was formerly called *Ununhexium (Uuh)*?
ⓐ. Livermorium
ⓑ. Flerovium
ⓒ. Tennessine
ⓓ. Darmstadtium
Correct Answer: Livermorium
Explanation: Element 116 was temporarily named *Ununhexium (Uuh)* and later renamed *Livermorium (Lv)* in honor of the Lawrence Livermore National Laboratory, which contributed to its discovery.
155. Which superheavy element was named after the American state Tennessee?
ⓐ. Nihonium
ⓑ. Moscovium
ⓒ. Tennessine
ⓓ. Californium
Correct Answer: Tennessine
Explanation: Element 117 was named *Tennessine (Ts)* in 2016 to honor Tennessee, home to Oak Ridge National Laboratory, where part of the research was conducted.
156. What is the official name of element 114?
ⓐ. Copernicium
ⓑ. Flerovium
ⓒ. Rutherfordium
ⓓ. Seaborgium
Correct Answer: Flerovium
Explanation: Element 114 was named *Flerovium (Fl)* after the Flerov Laboratory of Nuclear Reactions in Russia. Its temporary name was *Ununquadium (Uuq)*.
157. Which element was named after the scientist Nicolaus Copernicus?
ⓐ. Nihonium
ⓑ. Copernicium
ⓒ. Roentgenium
ⓓ. Meitnerium
Correct Answer: Copernicium
Explanation: Element 112 was named *Copernicium (Cn)* to honor Nicolaus Copernicus, who proposed the heliocentric model of the universe. Its temporary name was *Ununbium (Uub)*.
158. What is the symbol of the element named after Yuri Oganessian?
ⓐ. Og
ⓑ. On
ⓒ. Oo
ⓓ. Os
Correct Answer: Og
Explanation: Oganesson, with the symbol *Og*, was named in 2016 after Yuri Oganessian, a Russian nuclear physicist who contributed to the discovery of several superheavy elements.
159. Which superheavy element has the official name *Roentgenium*?
ⓐ. Z = 109
ⓑ. Z = 110
ⓒ. Z = 111
ⓓ. Z = 112
Correct Answer: Z = 111
Explanation: Element 111 is *Roentgenium (Rg)*, named after Wilhelm Roentgen, discoverer of X-rays. Its temporary name was *Unununium (Uuu)*.
160. Which element was named *Meitnerium (Mt)* in honor of physicist Lise Meitner?
ⓐ. Z = 106
ⓑ. Z = 107
ⓒ. Z = 108
ⓓ. Z = 109
Correct Answer: Z = 109
Explanation: Element 109, *Meitnerium (Mt)*, was named after Lise Meitner, who contributed to the discovery of nuclear fission. Its temporary name was *Unnilennium (Une)*.
161. What defines the “block” of an element in the periodic table?
ⓐ. Its most common oxidation state
ⓑ. The subshell into which the differentiating (last added) electron enters
ⓒ. Its position (left, middle, right) only
ⓓ. Its metallic/non-metallic character
Correct Answer: The subshell into which the differentiating (last added) electron enters
Explanation: Blocks (s, p, d, f) are classified by the subshell receiving the differentiating electron in the ground-state electronic configuration. If it goes to an s, p, d, or f subshell, the element is in the s-, p-, d-, or f-block respectively. This criterion is electronic-structure based, not just table location or typical oxidation states, and remains consistent across periods.
162. Which group numbers correspond to the s-block?
ⓐ. 1–2
ⓑ. 3–12
ⓒ. 13–18
ⓓ. 2–3
Correct Answer: 1–2
Explanation: The s-block occupies the far left of the long form table and comprises Groups 1 (alkali metals, ns¹) and 2 (alkaline earth metals, ns²). Their chemistry—strong electropositivity, formation of basic oxides/hydroxides, and predominant +1/+2 oxidation states—stems from the ease of losing the ns electrons.
163. The configuration $[\text{Ne}]\,3s^2\,3p^3$ belongs to which block?
ⓐ. s-block
ⓑ. p-block
ⓒ. d-block
ⓓ. f-block
Correct Answer: p-block
Explanation: The differentiating electron enters a p subshell (3p). Hence, the element lies in the p-block (Groups 13–18). A $ns^2 np^3$ valence pattern corresponds to Group 15 (e.g., phosphorus family), typically showing oxidation states −3, +3, and +5 and non-metallic/ metalloid behavior across the period.
164. Which of the following is NOT a d-block element?
ⓐ. Sc
ⓑ. Cu
ⓒ. Zn
ⓓ. Al
Correct Answer: Al
Explanation: Sc, Cu, and Zn are in Groups 3, 11, and 12 respectively (the d-block zone). Aluminum is in Group 13 and is a p-block element. Note: although Zn is d-block, by IUPAC “transition element” definition (incomplete d subshell in atom/ions), Zn (d¹⁰ in common states) is not a “transition element,” but it remains d-block by location/electron entry.
165. The f-block elements (inner transition elements) consist of:
ⓐ. Only lanthanides (Z = 57–71)
ⓑ. Only actinides (Z = 89–103)
ⓒ. Both lanthanides and actinides, shown separately at the bottom
ⓓ. Elements of Groups 3–12
Correct Answer: Both lanthanides and actinides, shown separately at the bottom
Explanation: The f-block includes the lanthanide series (commonly La/Ln: 57–71) and actinide series (Ac/An: 89–103). Their differentiating electrons largely occupy f orbitals. They are displayed in two detached rows to keep the table compact while preserving periodic order.
166. Which statement about helium’s “block” and “group placement” is correct?
ⓐ. Helium is a p-block element and belongs to Group 18
ⓑ. Helium is an s-block element but is placed in Group 18 due to inertness
ⓒ. Helium is a d-block element but behaves like a noble gas
ⓓ. Helium belongs to Group 2 because it has 1s²
Correct Answer: Helium is an s-block element but is placed in Group 18 due to inertness
Explanation: Helium’s configuration is $1s^2$, so by differentiating electron it is s-block. However, its complete valence shell renders it chemically inert like the noble gases, so it’s positioned in Group 18. Thus, “block” (electronic entry) and “group placement” (chemical properties) need not coincide perfectly.
167. What is the maximum number of elements that can appear consecutively in the s-block within any period?
ⓐ. 1
ⓑ. 2
ⓒ. 6
ⓓ. 10
Correct Answer: 2
Explanation: An s subshell has only one orbital ($m_l=0$) accommodating 2 electrons with opposite spins. Therefore, at most two elements per period have their differentiating electron in s (ns¹, ns²), corresponding to Groups 1 and 2.
168. Identify the block for the configuration $[\text{Xe}]\,4f^7\,6s^2$.
ⓐ. s-block
ⓑ. p-block
ⓒ. d-block
ⓓ. f-block
Correct Answer: f-block
Explanation: The presence of electrons in the 4f subshell indicates an f-block element (lanthanides in this case). f-block elements typically show +3 oxidation state dominantly, exhibit “lanthanoid contraction,” and possess characteristic magnetic/ spectroscopic behavior due to shielded f-electrons.
169. Which block elements typically have the lowest ionization enthalpies in their respective periods and form strongly basic oxides/hydroxides?
ⓐ. s-block
ⓑ. p-block
ⓒ. d-block
ⓓ. f-block
Correct Answer: s-block
Explanation: s-block (Groups 1 and 2) metals have loosely held ns electrons, leading to low ionization enthalpy and high electropositivity. They readily form $M_2O$/$MO$ and $MOH$/$M(OH)_2$, which are basic. Their reactivity increases down the group as atomic size increases and ionization enthalpy decreases.
170. Which statement correctly distinguishes “d-block” from “transition element” in IUPAC terms?
ⓐ. All d-block elements are transition elements without exception
ⓑ. d-Block membership depends on group, transition status depends on isotopes
ⓒ. d-Block refers to table location/electron entry; “transition element” requires a partially filled d subshell in atom or common ions
ⓓ. Transition elements are only Groups 1–2
Correct Answer: d-Block refers to table location/electron entry; “transition element” requires a partially filled d subshell in atom or common ions
Explanation: The d-block spans Groups 3–12 by differentiating electron criterion. IUPAC defines a transition element as having an incomplete d subshell in the atom or in at least one common oxidation state. Thus Zn, Cd, Hg are d-block but not “transition” (d¹⁰, no typical partially filled d in common ions), whereas Fe, Co, Ni clearly are.
171. For s-block elements, how is the group number related to the number of valence electrons?
ⓐ. Group number = number of valence electrons × 2
ⓑ. Group number = number of valence electrons
ⓒ. Group number = number of shells
ⓓ. Group number = atomic number
Correct Answer: Group number = number of valence electrons
Explanation: In the s-block (Groups 1 and 2), the group number equals the number of electrons in the outermost s-orbital. For example, alkali metals (ns¹) belong to Group 1, and alkaline earth metals (ns²) belong to Group 2.
172. The general valence shell configuration of Group 2 elements is:
ⓐ. ns¹
ⓑ. ns²
ⓒ. ns²np¹
ⓓ. ns²np²
Correct Answer: ns²
Explanation: Group 2 elements (alkaline earth metals) have two valence electrons in the s-orbital. For example, Mg: \[Ne] 3s². This accounts for their common oxidation state of +2.
173. In p-block elements, how is the group number related to valence electrons?
ⓐ. Group number = number of valence electrons
ⓑ. Group number = 10 + number of valence electrons
ⓒ. Group number = period number
ⓓ. Group number = (valence electrons – 10)
Correct Answer: Group number = 10 + number of valence electrons
Explanation: For p-block (Groups 13–18), the group number is 10 + the number of valence electrons (ns²np¹–⁶). Example: Carbon (ns²np²) has 4 valence electrons → 10 + 4 = Group 14.
174. The electronic configuration of an element is \[Ne] 3s²3p⁵. To which group does it belong?
ⓐ. Group 15
ⓑ. Group 16
ⓒ. Group 17
ⓓ. Group 18
Correct Answer: Group 17
Explanation: The element has 7 valence electrons (ns²np⁵). For p-block, group number = 10 + 7 = 17. This corresponds to the halogen family (chlorine in this case).
175. What is the general outer configuration of Group 14 elements?
ⓐ. ns²np¹
ⓑ. ns²np²
ⓒ. ns²np³
ⓓ. ns²np⁴
Correct Answer: ns²np²
Explanation: Group 14 elements (C, Si, Ge, Sn, Pb) have 4 valence electrons (ns²np²). This leads to common oxidation states of +4 and sometimes +2 due to inert pair effect in heavier members.
176. The group number of an element with outer configuration ns²np⁶ is:
ⓐ. 16
ⓑ. 17
ⓒ. 18
ⓓ. 8
Correct Answer: 18
Explanation: ns²np⁶ corresponds to a complete octet, the configuration of noble gases. With 8 valence electrons, group number = 10 + 8 = 18, the group of chemically inert elements.
177. For d-block elements, how is the group number determined?
ⓐ. Group number = number of d and s electrons in outer shell
ⓑ. Group number = number of f electrons
ⓒ. Group number = atomic number
ⓓ. Group number = period number
Correct Answer: Group number = number of d and s electrons in outer shell
Explanation: For transition metals, group number equals the sum of electrons in (n–1)d and ns orbitals. Example: Fe (\[Ar] 3d⁶4s²) → 6 + 2 = 8 → Group 8.
178. The electronic configuration of an element is \[Ar] 3d¹⁰4s²4p¹. Which group does it belong to?
ⓐ. Group 13
ⓑ. Group 14
ⓒ. Group 15
ⓓ. Group 12
Correct Answer: Group 13
Explanation: The element has ns²np¹ (valence = 3). For p-block, group number = 10 + 3 = 13. This corresponds to gallium (Ga).
179. The general electronic configuration of f-block elements is:
ⓐ. (n–2)f¹–¹⁴ (n–1)d⁰–¹ ns²
ⓑ. ns¹–²
ⓒ. ns²np¹–⁶
ⓓ. (n–1)d¹–¹⁰ ns¹–²
Correct Answer: (n–2)f¹–¹⁴ (n–1)d⁰–¹ ns²
Explanation: f-block elements (lanthanides and actinides) involve filling of (n–2)f orbitals. They generally also have ns² electrons and sometimes (n–1)d electrons. This configuration gives rise to variable oxidation states and special magnetic properties.
180. The electronic configuration of an element is \[Kr] 4d¹⁰5s²5p⁵. To which group does it belong?
ⓐ. Group 15
ⓑ. Group 16
ⓒ. Group 17
ⓓ. Group 18
Correct Answer: Group 17
Explanation: With ns²np⁵, the element has 7 valence electrons. Group number for p-block = 10 + 7 = 17. This configuration corresponds to iodine, a halogen.
181. What is meant by the covalent radius of an atom?
ⓐ. Half the distance between nuclei of two adjacent atoms in a metallic lattice
ⓑ. Half the distance between nuclei of two atoms bonded covalently
ⓒ. Distance from nucleus to outermost shell in an isolated atom
ⓓ. Half the distance between two non-bonded atoms in different molecules
Correct Answer: Half the distance between nuclei of two atoms bonded covalently
Explanation: The covalent radius is defined as half the internuclear distance between two atoms joined by a covalent bond of the same element. For example, in Cl₂, the bond length is 198 pm, so covalent radius = 99 pm.
182. Which of the following best defines metallic radius?
ⓐ. Half the internuclear distance between two adjacent ions
ⓑ. Half the internuclear distance between two adjacent atoms in a metallic lattice
ⓒ. Half the bond length of a covalent bond
ⓓ. The radius of the outermost electron cloud
Correct Answer: Half the internuclear distance between two adjacent atoms in a metallic lattice
Explanation: In metals, atoms are closely packed in a lattice. The metallic radius is taken as half the internuclear distance between two nearest neighboring atoms in the metal crystal lattice. For example, the metallic radius of sodium is \~186 pm.
183. What is the van der Waals radius?
ⓐ. Half the distance between two atoms bonded covalently
ⓑ. Half the distance between two closest nuclei of non-bonded atoms
ⓒ. Distance of nucleus to outermost orbital electron
ⓓ. Radius of the cation formed
Correct Answer: Half the distance between two closest nuclei of non-bonded atoms
Explanation: The van der Waals radius is larger than covalent radius. It is measured between two atoms not bonded but held by weak van der Waals forces. Example: van der Waals radius of chlorine is \~180 pm, larger than its covalent radius (\~99 pm).
184. Which of the following is generally the largest for the same atom?
ⓐ. Covalent radius
ⓑ. Metallic radius
ⓒ. van der Waals radius
ⓓ. Ionic radius of cation
Correct Answer: van der Waals radius
Explanation: The van der Waals radius is always larger because it measures the distance between weakly interacting atoms, where electron clouds do not overlap much. Covalent and metallic radii are smaller due to strong bonding interactions, and cations are the smallest due to high nuclear pull.
185. Why is the van der Waals radius larger than the covalent radius?
ⓐ. Because covalent bonds involve electron sharing leading to closer approach of nuclei
ⓑ. Because metallic bonding is stronger
ⓒ. Because ions repel each other
ⓓ. Because atomic number decreases
Correct Answer: Because covalent bonds involve electron sharing leading to closer approach of nuclei
Explanation: In covalent bonding, electron cloud overlap reduces internuclear distance, so covalent radii are smaller. In contrast, van der Waals forces are weak, so atoms stay farther apart, resulting in a larger radius.
186. What is the relation between metallic radius and covalent radius for the same element?
ⓐ. Metallic radius ≈ Covalent radius
ⓑ. Metallic radius is usually smaller
ⓒ. Metallic radius is usually larger
ⓓ. They cannot be compared
Correct Answer: Metallic radius is usually larger
Explanation: In metals, atoms are bound by delocalized electrons with less directional overlap compared to covalent bonds. Thus, metallic radii are usually about 10% larger than covalent radii. Example: Sodium covalent radius \~154 pm, metallic radius \~186 pm.
187. Which of the following elements shows the smallest covalent radius?
ⓐ. Lithium
ⓑ. Beryllium
ⓒ. Fluorine
ⓓ. Oxygen
Correct Answer: Fluorine
Explanation: Covalent radius decreases across a period due to increasing nuclear charge. Among the options, fluorine (Z = 9) has the smallest size with a covalent radius of \~64 pm.
188. Arrange the following radii in decreasing order for a chlorine atom: van der Waals radius, metallic radius, covalent radius.
ⓐ. van der Waals > metallic > covalent
ⓑ. metallic > van der Waals > covalent
ⓒ. covalent > metallic > van der Waals
ⓓ. van der Waals > covalent > metallic
Correct Answer: van der Waals > metallic > covalent
Explanation: For the same atom, van der Waals radius is the largest due to weak non-bonded interaction, metallic radius is intermediate due to metallic bonding, and covalent radius is the smallest because of strong bond overlap.
189. Which type of radius is most useful in predicting bond lengths in organic molecules?
ⓐ. Metallic radius
ⓑ. Ionic radius
ⓒ. Covalent radius
ⓓ. van der Waals radius
Correct Answer: Covalent radius
Explanation: In organic compounds, atoms are held together by covalent bonds. Hence, covalent radii are used to estimate bond lengths, e.g., C–C bond length is approximately the sum of two carbon covalent radii.
190. Which factor most strongly influences covalent radius across a period?
ⓐ. Increase in atomic mass
ⓑ. Increase in nuclear charge without much shielding
ⓒ. Increase in van der Waals forces
ⓓ. Increase in number of isotopes
Correct Answer: Increase in nuclear charge without much shielding
Explanation: Across a period, protons are added to the nucleus while electrons are added to the same shell. Increased nuclear charge pulls the electrons closer, reducing the covalent radius. This explains why atomic radii decrease across a period from left to right.
191. What is the general trend of atomic radius across a period in the periodic table?
ⓐ. It increases from left to right
ⓑ. It decreases from left to right
ⓒ. It remains constant
ⓓ. It increases irregularly
Correct Answer: It decreases from left to right
Explanation: As we move across a period, the number of protons increases, leading to a stronger nuclear charge. Electrons are added to the same shell, so shielding is minimal. The stronger attraction pulls electrons closer, reducing the atomic radius.
192. What is the general trend of atomic radius down a group?
ⓐ. It decreases continuously
ⓑ. It increases continuously
ⓒ. It first decreases then increases
ⓓ. It remains constant
Correct Answer: It increases continuously
Explanation: Down a group, new electron shells are added, increasing the distance of valence electrons from the nucleus. Although nuclear charge also increases, the effect of extra shells and shielding dominates, leading to larger radii.
193. Which element has the smallest atomic radius in Period 2?
ⓐ. Lithium
ⓑ. Oxygen
ⓒ. Fluorine
ⓓ. Neon
Correct Answer: Neon
Explanation: Neon (Z = 10) has the highest nuclear charge in Period 2. Despite having the same shell as other elements in the period, its stronger effective nuclear pull shrinks the size, making it the smallest.
194. Which element has the largest atomic radius in Group 1 (alkali metals)?
ⓐ. Lithium
ⓑ. Sodium
ⓒ. Potassium
ⓓ. Cesium
Correct Answer: Cesium
Explanation: Cesium lies lowest in Group 1, having the maximum number of shells among the options. The increased distance from the nucleus and shielding effect result in the largest atomic radius in the group.
195. Why does atomic radius decrease across a period?
ⓐ. Because of an increase in number of shells
ⓑ. Because of increase in shielding effect
ⓒ. Because of increasing nuclear charge attracting electrons more strongly
ⓓ. Because of formation of ions
Correct Answer: Because of increasing nuclear charge attracting electrons more strongly
Explanation: Across a period, protons are added to the nucleus, increasing nuclear charge. Electrons enter the same shell, so shielding remains almost constant. The stronger attraction pulls the electrons inward, reducing the radius.
196. What happens to ionic radius when an atom forms a cation?
ⓐ. Radius increases
ⓑ. Radius decreases
ⓒ. Radius remains the same
ⓓ. Radius becomes infinite
Correct Answer: Radius decreases
Explanation: A cation forms when an atom loses one or more electrons. Loss of electrons reduces electron–electron repulsion and increases effective nuclear charge per electron, pulling remaining electrons closer, thus reducing radius. Example: Na (186 pm) → Na⁺ (102 pm).
197. What happens to ionic radius when an atom forms an anion?
ⓐ. Radius decreases
ⓑ. Radius increases
ⓒ. Radius remains constant
ⓓ. Radius becomes smaller than a cation
Correct Answer: Radius increases
Explanation: When an atom gains electrons to form an anion, electron–electron repulsion increases, and nuclear attraction per electron decreases. This causes the electron cloud to expand, increasing radius. Example: Cl (99 pm) → Cl⁻ (181 pm).
198. In an isoelectronic series (same number of electrons), how does ionic radius vary?
ⓐ. Larger the atomic number, larger the radius
ⓑ. Smaller the nuclear charge, smaller the radius
ⓒ. Higher the nuclear charge, smaller the radius
ⓓ. Radius remains constant
Correct Answer: Higher the nuclear charge, smaller the radius
Explanation: In isoelectronic species, all ions have the same electron number but different nuclear charges. Greater nuclear charge pulls the electron cloud more strongly, reducing radius. For example, O²⁻ > F⁻ > Na⁺ > Mg²⁺.
199. Which of the following has the largest ionic radius?
ⓐ. Na⁺
ⓑ. Mg²⁺
ⓒ. F⁻
ⓓ. O²⁻
Correct Answer: O²⁻
Explanation: O²⁻ has the greatest negative charge among the options, leading to the maximum electron–electron repulsion and least effective nuclear pull. Hence, its radius is the largest in this group of isoelectronic ions.
200. Why is the ionic radius of anions larger than their parent atoms?
ⓐ. Because they lose electrons
ⓑ. Because they gain electrons, increasing repulsion
ⓒ. Because their nuclear charge decreases
ⓓ. Because they form metallic bonds
Correct Answer: Because they gain electrons, increasing repulsion
Explanation: In anions, the addition of electrons increases inter-electron repulsion, causing the electron cloud to expand. Since nuclear charge remains the same but is distributed over more electrons, the radius increases compared to the neutral atom.
The chapter Classification of Elements and Periodicity in Properties is a cornerstone of Class 11 Chemistry (NCERT/CBSE syllabus) and carries great weightage in both board exams and competitive exams like JEE, NEET, and state-level entrance tests. It helps students understand how elements are arranged in the modern periodic table based on their atomic numbers and electronic configurations. This chapter also discusses periodic variation of physical and chemical properties, including trends in atomic size, metallic and non-metallic character, ionization enthalpy, electron gain enthalpy, and electronegativity. Advanced concepts like valency, variable valency, diagonal relationship (Li with Mg, Be with Al, B with Si), and anomalous behavior of the first element in each group are also covered. This complete set offers 350 MCQs with detailed answers, carefully divided into 4 parts. Here in Part 2, you will practice the next 100 MCQs with step-by-step explanations to boost your conceptual clarity and exam performance.
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