Periodic Classification of Element
Short Questions
1. How does Dobereiner’s Triad illustrate the relationship between atomic masses?
Answer : The atomic mass of the middle element is approximately the arithmetic mean of the other two.
2. Why was Newlands’ Law of Octaves limited to elements up to calcium?
Answer : Beyond calcium, the pattern of repeating properties every eighth element broke down due to dissimilar properties.
3. What was the significance of Mendeleev’s prediction of eka-silicon?
Answer : It accurately predicted the properties of germanium, validating his periodic table.
4. Why did Mendeleev leave gaps in his periodic table?
Answer : To accommodate undiscovered elements with predicted properties based on periodic trends.
5. How does the modern periodic law differ from Mendeleev’s periodic law?
Answer : It is based on atomic number rather than atomic mass, providing a more fundamental classification.
6. Why are lanthanides and actinides placed separately in the modern periodic table?
Answer : To maintain a compact table structure, as their properties are similar and fit in the f-block.
7. What determines the period of an element in the modern periodic table?
Answer : The number of electron shells occupied by electrons.
8. Why do isotopes pose no issue in the modern periodic table?
Answer : They have the same atomic number and chemical properties, so they occupy the same position.
9. What is the electron capacity formula for a shell with principal quantum number n?
Answer : 2n² electrons.
10. Why is helium the smallest atom in the periodic table?
Answer : Its high nuclear charge and only two electrons result in the smallest atomic radius.
11. How does the effective nuclear charge affect atomic radius across a period?
Answer : It increases, pulling electrons closer to the nucleus, reducing atomic radius.
12. Why do alkali metals have low electronegativity?
Answer : Their single valence electron is loosely held due to large atomic size and low nuclear charge.
13. What is the trend in electropositivity down Group 1?
Answer : Electropositivity increases as atomic size grows, reducing the nucleus’s hold on valence electrons.
14. Why are nonmetals found in the p-block of the periodic table?
Answer : Their high electronegativity and tendency to gain electrons place them in Groups 13–18.
15. What causes the gradation in reactivity of alkaline earth metals down Group 2?
Answer : Increasing atomic size reduces nuclear attraction, making electron loss easier, thus increasing reactivity.
16. Why does fluorine have the highest electronegativity in the periodic table?
Answer : Its small size and high nuclear charge strongly attract additional electrons.
17. How does the zig-zag line in the periodic table separate element types?
Answer : It divides metals (left) from nonmetals (right), with metalloids along the line.
18. Why does the third period have only eight elements despite a shell capacity of 18?
Answer : The octet rule limits valence electrons to eight, stabilizing the electron configuration.
19. What is the valency of an element with electronic configuration 2, 8, 7?
Answer : 1, as it gains one electron to achieve a stable octet.
20. Why are noble gases chemically inert?
Answer : Their complete valence shells (octet) make them stable, with no tendency to gain or lose electrons.
Long Questions
1. How did Dobereiner’s Triads contribute to the development of the periodic table?
Answer : Dobereiner’s Triads grouped elements with similar properties, showing a pattern in atomic masses. For example, in Cl (35.5), Br (79.9), I (126.9), Br’s mass is nearly the mean of Cl and I. This laid the groundwork for later classifications like Mendeleev’s periodic table.
2. Analyze the limitations of Newlands’ Law of Octaves and their impact on its acceptance.
Answer : The law failed beyond calcium, as heavier elements didn’t fit the octave pattern, and it placed dissimilar elements like Co and Ni together. It also lacked space for new elements, reducing its predictive power. These flaws led to its rejection until more robust systems emerged.
3. Evaluate the significance of Mendeleev’s corrections to atomic masses in his periodic table.
Answer : Mendeleev revised atomic masses, like beryllium’s from 14.09 to 9.4, to align elements with their properties. This ensured correct placement, such as beryllium before boron, enhancing the table’s accuracy. Such corrections bolstered the table’s credibility and predictive success.
4. Why was the discovery of noble gases a challenge for Mendeleev’s periodic table, and how was it resolved?
Answer : Noble gases, discovered later, had no place in Mendeleev’s original table due to their inert nature. Mendeleev created a new ‘zero’ group to accommodate them without disrupting the table. This adaptability demonstrated the table’s flexibility and scientific robustness.
5. How does the modern periodic table overcome the ambiguity in cobalt and nickel’s placement?
Answer : In Mendeleev’s table, cobalt and nickel’s similar atomic masses (59) caused placement issues. The modern table uses atomic numbers (Co: 27, Ni: 28), ensuring cobalt precedes nickel. This reflects their electronic configurations and properties, resolving the ambiguity.
Explain the periodic trend in atomic radius within a period and its underlying cause.
Answer : Atomic radius decreases across a period due to increasing atomic number, adding protons and electrons in the same shell. The higher nuclear charge increases the effective nuclear charge, pulling electrons closer. For example, in period 2, Li (152 pm) is larger than F (72 pm).
Why do halogens exhibit increasing reactivity up Group 17, and how is this related to electronegativity?
Answer : Halogens’ reactivity increases up Group 17 (e.g., I to F) due to decreasing atomic size and increasing electronegativity. Smaller atoms like fluorine have a stronger nuclear pull, attracting electrons more effectively to form anions. This makes fluorine the most reactive halogen.
Discuss the role of electronic configuration in determining an element’s position in the modern periodic table.
Answer : Electronic configuration dictates the number of valence electrons (group) and shells (period). For example, Na (2, 8, 1) is in Group 1, Period 3, due to one valence electron in the third shell. This systematic arrangement predicts chemical properties accurately.
How does the variation in metallic character across a period and down a group influence element properties?
Answer : Metallic character decreases across a period as electronegativity rises, making elements less likely to lose electrons (e.g., Na to Cl). Down a group, metallic character increases as larger atoms lose electrons more easily (e.g., Li to Cs). This affects reactivity and compound formation.
Why is the placement of hydrogen in the periodic table still debated, and what are the arguments for its position?
Answer : Hydrogen resembles both alkali metals (forms H⁺ like Na⁺) and halogens (forms H₂ like Cl₂ and H⁻). Its electronic configuration (1s¹) supports Group 1, but its high electronegativity aligns with Group 17. The modern table leaves its position ambiguous due to these dual properties.
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