Class 11 Chemistry Classification of Elements and Periodicity in Properties

Class 11 Chemistry Classification of Elements and Periodicity in Properties

1. What is arguably the most important concept in chemistry? The Periodic Table.
2. What does the Periodic Table provide to the whole of chemistry? A succinct organization.
3. What do chemical elements display in the Periodic Table? Trends and family groupings.
4. Who is Glenn T. Seaborg recognized for? His work leading to the reconfiguration of the periodic table, placing actinoids below lanthanoids, and discovery of transuranium elements.
5. What concept led to the development of the Periodic Table? Grouping elements according to their properties.
6. What law is fundamental to understanding element classification? The Periodic Law.
7. What are the bases for modern periodic classification? Atomic number and electronic configuration.
8. What skill is listed for elements with Z > 100? Naming them according to IUPAC nomenclature.
9. How are elements classified into blocks? Into s, p, d, f blocks.
10. What is a key objective regarding element properties? Recognizing periodic trends in physical and chemical properties.
11. What are the basic units of all types of matter? Elements.
12. How many elements were known in 1800? 31 elements.
13. By 1865, approximately how many elements were known? 63 elements.
14. How many elements are known at present (according to the source)? 114 elements.
15. What type of elements are the recently discovered ones? Man-made.
16. Why was classifying elements necessary for scientists? To systematize knowledge and ease the study of many elements.
17. Who was the first to consider trends among properties of elements? German chemist Johann Dobereiner.
18. What did Dobereiner note in 1829? Similarity among properties of groups of three elements (triads).
19. What was unique about the middle element in Dobereiner’s triads? Its atomic weight was about halfway between the other two.
20. Why was Dobereiner’s law of triads dismissed? It worked only for a few elements and was considered a coincidence.
21. Who was A.E.B. de Chancourtois? A French geologist who attempted to classify elements.
22. How did Chancourtois arrange elements in 1862? In order of increasing atomic weights in a cylindrical table.
23. What did Chancourtois observe about element properties? Their periodic recurrence.
24. Who profounded the Law of Octaves in 1865? English chemist John Alexander Newlands.
25. How did Newlands arrange elements in his Law of Octaves? In increasing order of their atomic weights.
26. What was the central idea of Newlands’ Law of Octaves? Every eighth element had properties similar to the first.
27. For which elements was Newlands’s Law of Octaves primarily true? Elements up to calcium.
28. What award did Newlands receive in 1887? The Davy Medal by the Royal Society, London.
29. Who developed the Periodic Law as we know it today? Russian chemist Dmitri Mendeleev and German chemist Lothar Meyer.
30. In what year did Mendeleev and Meyer independently propose their periodic ideas? 1869.
31. What did Lothar Meyer plot against atomic weight? Physical properties like atomic volume, melting point, and boiling point.
32. Who is generally credited with the development of the Modern Periodic Table? Dmitri Mendeleev.
33. What is Mendeleev’s Periodic Law? “The properties of the elements are a periodic function of their atomic weights“.
34. How did Mendeleev arrange elements in his table? In horizontal rows and vertical columns by increasing atomic weights.
35. What was a key feature of Mendeleev’s classification system? He placed elements with similar properties in the same vertical column or group.
36. What did Mendeleev sometimes ignore when classifying elements? The order of atomic weights.
37. Give an example of Mendeleev ignoring atomic weight order. He placed iodine (lower atomic weight) with halogens, not with tellurium.
38. What did Mendeleev do for undiscovered elements? He left gaps in his table and predicted their existence.
39. What did Mendeleev call the element predicted under aluminium? Eka-aluminium.
40. What was Eka-aluminium later discovered as? Gallium.
41. What did Mendeleev call the element predicted under silicon? Eka-silicon.
42. What was Eka-silicon later discovered as? Germanium.
43. What made Mendeleev and his Periodic Table famous? His bold quantitative predictions and their success.
44. What was unknown to chemists when Mendeleev developed his table? The internal structure of the atom.
45. Who observed regularities in X-ray spectra in 1913? English physicist Henry Moseley.
46. What did Moseley show to be a more fundamental property of an element than atomic mass? Atomic number (Z).
47. What is the Modern Periodic Law? “The physical and chemical properties of the elements are periodic functions of their atomic numbers“.
48. What is atomic number equal to? Nuclear charge (number of protons) or number of electrons in a neutral atom.
49. What is the Periodic Law essentially a consequence of? The periodic variation in electronic configurations.
50. What is the most convenient and widely used form of the Periodic Table? The “long form”.
51. What are the horizontal rows in the Periodic Table called? Periods.
52. What are the vertical columns in the Periodic Table called? Groups or families.
53. According to IUPAC, how are groups numbered? From 1 to 18.
54. How many periods are there in total in the Periodic Table? Seven.
55. What does the period number correspond to? The highest principal quantum number (n) of the elements in that period.
56. How many elements are in the first period? 2 elements.
57. How many elements are in the second and third periods? 8 elements each.
58. How many elements are in the fourth and fifth periods? 18 elements each.
59. How many elements are in the sixth period? 32 elements.
60. Where are lanthanoids and actinoids placed in the Periodic Table? In separate panels at the bottom.
61. What element is named Seaborgium (Sg)? Element 106.
62. What organization ratifies names for new elements? IUPAC.
63. Why has naming new elements become controversial? They are unstable, made in minute quantities, and there can be competing claims for discovery.
64. What is the IUPAC recommendation for temporary naming of elements with Z > 100? A systematic nomenclature derived from the atomic number.
65. What numerical root is used for the digit ‘0’ in IUPAC nomenclature? nil.
66. What numerical root is used for the digit ‘1’ in IUPAC nomenclature? un.
67. What suffix is added at the end of the numerical roots for IUPAC temporary names? -ium.
68. What is the IUPAC temporary name for element 101? Unnilunium.
69. What is the IUPAC temporary name for element 118? Ununoctium.
70. What is the symbol for Ununbium? Uub.
71. What is the official name for element 104? Rutherfordium.
72. What is the official name for element 112? Copernicium.
73. What is the official name for element 118? Oganesson.
74. What is the temporary IUPAC name and symbol for the element with atomic number 120? Unbinilium (Ubn).
75. What is the distribution of electrons into orbitals of an atom called? Its electronic configuration.
76. What does an element’s location in the Periodic Table reflect? The quantum numbers of the last orbital filled.
77. What does the period number in the Periodic Table indicate? The value of ‘n’ for the outermost or valence shell.
78. What process is associated with successive periods in the Periodic Table? The filling of the next higher principal energy level.
79. How many electrons can the first period (n=1) accommodate? 2 electrons.
80. What elements complete the K shell in the first period? Hydrogen (1s1) and Helium (1s2).
81. Which period starts with Lithium? The second period (n=2).
82. What orbitals are filled in the second period? 2s and 2p orbitals.
83. How many elements are in the second period? 8 elements.
84. What orbitals are filled in the third period? 3s and 3p orbitals.
85. What transition series is found in the fourth period? The 3d transition series.
86. Where does the 3d transition series start? At scandium (Z=21).
87. What is the electronic configuration of Scandium (Z=21)? 3d14s2.
88. Where does the 3d transition series end? At zinc (Z=30).
89. How many elements are in the fourth period? 18 elements.
90. Which transition series is found in the fifth period? The 4d transition series.
91. Which orbitals are successively filled in the sixth period (n=6)? 6s, 4f, 5d, and 6p orbitals.
92. What is the 4f-inner transition series called? The lanthanoid series.
93. Where does the filling of 4f orbitals begin? With cerium (Z=58).
94. Where does the filling of 4f orbitals end? At lutetium (Z=71).
95. Which period includes most man-made radioactive elements? The seventh period.
96. What is the 5f-inner transition series known as? The actinoid series.
97. Why are inner transition series placed separately in the Periodic Table? To maintain structure and keep elements with similar properties in one column.
98. Why are there 18 elements in the 5th period? Because the 5s, 4d, and 5p orbitals (9 orbitals) are filled, accommodating 18 electrons.
99. What do elements in the same vertical column (group) share? Similar valence shell electronic configurations and properties.
100. What is the valence shell electronic configuration of Group 1 elements (alkali metals)? ns1.
101. What principle provides the theoretical foundation for periodic classification? The aufbau (build up) principle.
102. How are elements classified into four blocks? Based on the type of atomic orbitals being filled.
103. Which block does helium strictly belong to based on its electronic configuration? The s-block.
104. Why is helium placed with Group 18 elements despite being an s-block element? It has a completely filled valence shell (1s2) and noble gas properties.
105. Why is hydrogen placed separately in the Periodic Table? It has only one s-electron, can be Group 1 or act like Group 17, making it a special case.
106. Which groups belong to the s-Block Elements? Group 1 (alkali metals) and Group 2 (alkaline earth metals).
107. What are the outermost electronic configurations of s-block elements? ns1 and ns2.
108. What is a key characteristic of s-block elements concerning ionization enthalpy? They are reactive metals with low ionization enthalpies.
109. What ions do alkali metals readily form? 1+ ions.
110. What ions do alkaline earth metals readily form? 2+ ions.
111. How do metallic character and reactivity change as you go down an s-block group? They increase.
112. Why are s-block elements not found pure in nature? Due to their high reactivity.
113. What is the predominant character of compounds formed by most s-block elements? Ionic.
114. Which groups comprise the p-Block Elements? Group 13 to 18.
115. What are s-block and p-block elements collectively called? Representative Elements or Main Group Elements.
116. How does the outermost electronic configuration vary in p-block elements across a period? From ns2np1 to ns2np6.
117. What elements are found at the end of each p-block period? Noble gas elements.
118. Why do noble gases exhibit very low chemical reactivity? They have completely filled valence shells (ns2np6), a very stable arrangement.
119. What are the halogens? Group 17 elements.
120. What are the chalcogens? Group 16 elements.
121. What is characteristic of halogens and chalcogens regarding electron gain enthalpy? They have highly negative electron gain enthalpies.
122. How do halogens attain stable noble gas configurations? By readily adding one electron.
123. How does non-metallic character change across a period (left to right)? It increases.
124. How does metallic character change down a p-block group? It increases.
125. Which groups comprise the d-Block Elements? Group 3 to 12.
126. What are d-block elements also known as? Transition Elements.
127. What characterizes d-block elements electronically? The filling of inner d orbitals.
128. What is the general outer electronic configuration of d-block elements? (n-1)d1-10ns0-2.
129. What is an exception to the general outer electronic configuration for Pd? 4d105s0.
130. What is the physical state of all d-block elements? They are all metals.
131. Name three characteristic properties of d-block elements. They form coloured ions, exhibit variable valence, and are often used as catalysts.
132. Which d-block elements do not show most properties of transition elements? Zn, Cd, and Hg.
133. Why are Zn, Cd, and Hg considered exceptions to typical transition elements? They have completely filled (n-1)d10ns2 configurations.
134. What is the role of transition metals between s-block and p-block elements? They form a bridge between the chemically active s-block metals and less active p-block elements.
135. What are the two rows of elements at the bottom of the Periodic Table called? The Lanthanoids and Actinoids.
136. What is the atomic number range for Lanthanoids? Ce(Z=58) – Lu(Z=71).
137. What is the atomic number range for Actinoids? Th(Z=90) – Lr(Z=103).
138. What is the general outer electronic configuration of f-block elements? (n-2)f1-14(n-1)d0–1ns2.
139. What is another name for f-block elements? Inner-Transition Elements.
140. How do the properties of elements within an f-block series compare? They are quite similar.
141. Why is the chemistry of early actinoids more complicated than lanthanoids? Due to the large number of oxidation states they can exhibit.
142. Are actinoid elements radioactive? Yes.
143. What are elements after uranium called? Transuranium Elements.
144. In which family would element Z=117 belong? The halogen family (Group 17).
145. What block does element Z=117 belong to? p-block.
146. In which group would element Z=120 be placed? Group 2 (alkaline earth metals).
147. What block does element Z=120 belong to? s-block.
148. What percentage of all known elements do metals comprise? More than 78%.
149. Where are metals generally found in the Periodic Table? On the left side.
150. What is an exception to metals being solid at room temperature? Mercury.
151. What are two properties of metals regarding shaping? They are malleable and ductile.
152. Where are non-metals located in the Periodic Table? At the top right-hand side.
153. What is the trend of properties across a horizontal row from left to right? From metallic to non-metallic.
154. What are characteristic properties of most non-metallic solids? They are brittle, neither malleable nor ductile.
155. How does metallic character change down a group? It increases.
156. How does non-metallic character change across a period from left to right? It increases.
157. What are elements bordering the zig-zag line in Fig. 3.3 called? Semi-metals or metalloids.
158. Name two elements classified as metalloids. Silicon (Si) and Germanium (Ge).
159. Arrange Si, Be, Mg, Na, P in increasing order of metallic character. P < Si < Be < Mg < Na.
160. Where is chemical reactivity highest in a period for metals? In Group 1 metals (extreme left).
161. Where is chemical reactivity highest in a period for non-metals? In Group 17 non-metals (extreme right).
162. How does chemical reactivity change down a group for alkali metals? It increases.
163. How does chemical reactivity change down a group for halogens? It decreases.
164. What properties are discussed as physical periodic trends? Atomic/ionic radii, ionization enthalpy, electron gain enthalpy, electronegativity.
165. Why is atomic size difficult to measure precisely? Atoms are very small, and electron clouds lack sharp boundaries.
166. How is the “Covalent radius” estimated for non-metallic elements? As half the bond distance between two atoms in a covalent molecule.
167. How is “metallic radius” defined for metals? As half the internuclear distance in a metallic crystal.
168. How does atomic size generally change across a period (left to right)? It decreases.
169. What causes the decrease in atomic size across a period? Increasing effective nuclear charge.
170. How does atomic radius change down a group? It increases regularly.
171. What causes the increase in atomic radius down a group? Increasing principal quantum number (n) and shielding effect of inner electrons.
172. Why is a cation smaller than its parent atom? It has fewer electrons but the same nuclear charge.
173. Why is an anion larger than its parent atom? It has more electrons, leading to increased electron-electron repulsion and decreased effective nuclear charge.
174. What are isoelectronic species? Atoms and ions that contain the same number of electrons.
175. For isoelectronic species, which one will have a smaller radius? The cation with the greater positive charge.
176. What is Ionization Enthalpy (∆iH)? The energy required to remove an electron from an isolated gaseous atom in its ground state.
177. Are ionization enthalpies always positive or negative? Always positive.
178. Why is the second ionization enthalpy higher than the first? It’s harder to remove an electron from a positively charged ion.
179. Where do maxima occur in the plot of first ionization enthalpies versus atomic number? At the noble gases.
180. Where do minima occur in the plot of first ionization enthalpies versus atomic number? At the alkali metals.
181. How does ionization enthalpy generally change across a period? It increases.
182. How does ionization enthalpy generally change down a group? It decreases.
183. What is the “shielding” or “screening” effect? Inner core electrons reduce the effective nuclear charge experienced by valence electrons.
184. Why does ionization enthalpy increase across a period? Increasing nuclear charge outweighs shielding.
185. Why does ionization enthalpy decrease down a group? Increased shielding outweighs increasing nuclear charge and valence electrons are farther away.
186. Why is the first ionization enthalpy of Boron slightly less than Beryllium? A 2p-electron (Boron) is easier to remove than a 2s-electron (Beryllium) due to less penetration and greater shielding.
187. Why is the first ionization enthalpy of Oxygen lower than Nitrogen? Oxygen has increased electron-electron repulsion due to two 2p-electrons occupying the same orbital.
188. What is Electron Gain Enthalpy (∆egH)? The enthalpy change when an electron is added to a neutral gaseous atom to form a negative ion.
189. For many elements, is electron gain enthalpy positive or negative? Negative (energy is released).
190. Why do halogens have very high negative electron gain enthalpies? They can attain stable noble gas electronic configurations by gaining an electron.
191. Do noble gases have positive or negative electron gain enthalpies? Large positive values.
192. How does electron gain enthalpy generally change across a period? Becomes more negative.
193. How does electron gain enthalpy generally change down a group? Becomes less negative.
194. Why is electron gain enthalpy of O or F less negative than S or Cl? The added electron goes to a smaller n=2 quantum level, experiencing significant repulsion.
195. What is Electronegativity? A qualitative measure of an atom’s ability to attract shared electrons in a chemical compound.
196. Is electronegativity a measurable quantity like ionization enthalpy? No, it is not a measurable quantity.
197. Who developed the most widely used electronegativity scale? Linus Pauling.
198. What arbitrary value did Pauling assign to fluorine for electronegativity? 4.0.
199. How does electronegativity generally change across a period (left to right)? It increases.
200. How does electronegativity generally change down a group? It decreases.