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| Classification Of Elements And Periodicity In Properties |
1. What is the basic theme of organisation in the periodic table?
Ans: The main idea behind the organisation of the periodic table is to arrange elements into periods and groups based on their similar characteristics. This setup helps to simplify and systematize the study of elements and their compounds. Elements with similar properties are grouped together in the periodic table.
2. Which important property did Mendeleev use to classify the elements in his periodic table and did he stick to that?
Ans: In his periodic table, Mendeleev arranged elements according to their atomic weights (or atomic masses). He placed the elements in periods and groups by increasing atomic weights and grouped elements having similar properties together. However, he did not always strictly follow this order. Sometimes, to keep elements with similar properties in the same group, he ignored atomic weight order—for example, iodine, which has a lower atomic weight than tellurium, was placed after tellurium since it shared more similarities with group VII elements like fluorine, chlorine, and bromine.
3. What is the basic difference in approach between Mendeleev’s Periodic Law and the Modern Periodic Law?
Ans: Mendeleev’s Periodic Law is based on the atomic masses of elements, stating their physical and chemical properties are periodic functions of atomic mass. The modern law, however, is based on atomic number, stating that an element’s properties are periodic functions of its atomic number. This shift resolved inconsistencies and better matched observed chemical behavior.
4. On the basis of quantum numbers, justify that the sixth period of the Periodic Table should have 32 elements.
Ans: The sixth period has the principal quantum number n=6, allowing for the filling of the 6s, 4f, 5d, and 6p subshells. There are a total of 16 orbitals for these subshells, and each orbital can hold two electrons. Therefore, the sixth period can accommodate 32 electrons, resulting in 32 elements in this period.
5. In terms of period and group, where would you locate the element with Z = 114?
Ans: The element with atomic number 114 is located in the 7th period of the periodic table and is the second element in the p-block of this period, making it part of group 14.
6. Write the atomic number of the element present in the third period and the seventeenth group of the Periodic Table.
Ans: The element present in the third period and 17th group is chlorine, whose atomic number is 17.
7. Which element do you think would have been named by (i) Lawrence Berkeley Laboratory (ii) Seaborg’s Group?
Ans: (i) Lawrencium (Lr, Z=103) and Berkelium (Bk, Z=97) were named by Lawrence Berkeley Laboratory. (ii) Seaborgium (Sg, Z=106) was named by Seaborg’s group.
8. Why do elements in the same group have similar physical and chemical properties?
Ans: Elements in the same group have the same number of valence electrons, which causes them to have similar physical and chemical properties.
9. What does atomic radius and ionic radius really mean to you?
Ans: The atomic radius is the size of an atom. For metals, the atomic radius is measured as the metallic radius, and for non-metals, as the covalent radius. The ionic radius is the radius of an ion (either cation or anion), calculated based on distances found in ionic crystals. Generally, cations are smaller and anions are larger than their corresponding atoms.
10. How does atomic radius vary in a period and in a group? How do you explain the variation?
Ans: Atomic radius decreases from left to right across a period due to increasing nuclear charge, pulling electrons closer. It increases down a group because new electron shells are added, increasing the distance between nucleus and outer electrons.
11. What do you understand about isoelectronic species? Name a species that will be isoelectronic with each of the following atoms or ions: (i) F⁻ (ii) Ar (iii) Mg²⁺ (iv) Rb⁺
Ans: Isoelectronic species are atoms or ions that have the same number of electrons. (i) F⁻ has 10 electrons, so isoelectronic species are Na⁺, Ne, and O²⁻. (ii) Ar has 18 electrons, so S²⁻, Cl⁻, and Ca²⁺ are isoelectronic with it. (iii) Mg²⁺ has 10 electrons, like Na⁺ and Ne. (iv) Rb⁺ has 36 electrons, so Br⁻, Kr, and Sr²⁺ are isoelectronic with it.
12. Consider the following species: N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺, and Al³⁺. (a) What is common in them? (b) Arrange them in the order of increasing ionic radii.
Ans: (a) All these species have 10 electrons and are isoelectronic. (b) As the positive charge increases, the ionic radius decreases. So the order of increasing ionic radii is: Al³⁺ < Mg²⁺ < Na⁺ < F⁻ < O²⁻ < N³⁻.
13. Explain why cations are smaller and anions larger in radii than their parent atoms.
Ans: Cations are formed by losing electrons, resulting in fewer electrons but the same nuclear charge, which pulls the remaining electrons closer; so, cations are smaller. Anions gain electrons, increasing electron repulsion for the same nuclear charge, so they have a larger radius than their parent atoms.
14. What is the significance of the terms ‘isolated gaseous atom’ and ‘ground state’ while defining the ionization enthalpy and electron gain enthalpy?
Ans: These terms mean that for accurate measurement and fair comparison, the atom must not be influenced by other atoms or be in an excited state. Ionization enthalpy and electron gain enthalpy are always defined for isolated gaseous atoms in their ground state.
15. Energy of an electron in the ground state of the hydrogen atom is \(2.18 \times 10^{-18}\) J. Calculate the ionization enthalpy of atomic hydrogen in terms of J mol⁻¹.
Ans: The ionization enthalpy per mole is \(2.18 \times 10^{-18}\) J × \(6.02 \times 10^{23}\) (= Avogadro's number), which comes out to approximately \(1.31 \times 10^{6}\) J mol⁻¹.
16. Among the 2nd period elements, the actual ionization enthalpies are in the order Li < B < Be < C < O < N < F < Ne. Explain why (i) Be has higher ionization enthalpy than B. (ii) O has lower ionization enthalpy than N and F.
Ans: (i) The 2s electrons in Be are more strongly held than the 2p electron in B, so Be has a higher ionization enthalpy. (ii) In oxygen, pairing in the 2p orbital increases electron-electron repulsion, making removal of an electron easier compared to nitrogen and fluorine.
17. How would you explain the fact that the first ionization enthalpy of sodium is lower than that of magnesium, but its second ionization enthalpy is higher than that of magnesium?
Ans: Sodium, after losing one electron, achieves a noble gas configuration, so its second electron is much harder to remove, making the second ionization enthalpy very high. In magnesium, the removal of two electrons still gives a stable configuration, so successive ionizations are not as difficult.
18. What are the various factors due to which the ionization enthalpy of the main group elements tends to decrease down a group?
Ans: Ionization enthalpy decreases down a group mainly because atomic size increases and the outer electrons are farther from the nucleus. Also, increased shielding by inner electrons makes it easier to remove an outer electron.
19. The first ionization enthalpy values (in kJ mol⁻¹) of Group 13 elements are: B (801), Al (577), Ga (579), In (558), Tl (589). How would you explain this deviation from the general trend?
Ans: The deviation is due to poor shielding by d and f electrons in gallium and thallium, making their valence electrons experience a greater effective nuclear charge, so their ionization enthalpy does not strictly decrease down the group.
20. Which of the following pairs of elements would have a more negative electron gain enthalpy: (i) O or F (ii) F or Cl?
Ans: (i) Fluorine has a more negative electron gain enthalpy than oxygen due to its smaller size and higher effective nuclear charge. (ii) Chlorine has a more negative electron gain enthalpy than fluorine because, in fluorine, strong electron-electron repulsions make addition of the second electron less favorable.
21. Would you expect the second electron gain enthalpy of O as positive, more negative or less negative than the first? Justify your answer.
Ans: It is positive because adding another electron to an already negatively charged ion increases repulsion, making energy input necessary.
22. What is the basic difference between the term electron gain enthalpy and electronegativity?
Ans: Electron gain enthalpy is the energy change when a neutral atom gains an electron, while electronegativity is the ability of an atom in a molecule to attract shared electrons.
23. How would you react to the statement that the electronegativity of N on the Pauling scale is 3.0 in all the nitrogen compounds?
Ans: This is incorrect. The value of electronegativity for nitrogen varies in different chemical environments; for instance, it is different in NH₃ and NO₂.
24. Describe the theory associated with the radius of an atom as it (a) gains an electron (b) loses an electron.
Ans: (a) When an atom gains an electron, its radius increases due to increased electron-electron repulsion. (b) When an atom loses an electron, its radius decreases because the same nuclear charge attracts fewer electrons more strongly.
25. Would you expect the first ionization enthalpies for two isotopes of the same element to be the same or different? Justify your answer.
Ans: They will be the same because ionization enthalpy depends on the number of protons and electrons, which are identical for isotopes.
26. What are the major differences between metals and nonmetals?
Ans: Metals tend to lose electrons, form basic oxides, and are good conductors, having low ionization enthalpy and being electropositive. Nonmetals tend to gain electrons, form acidic oxides, are poor conductors, have high ionization enthalpy and are electronegative.
27. Use the periodic table to answer the following questions:
(a) An element with five electrons in the outer subshell belongs to the halogen group (ns²np⁵): for example, F, Cl, Br, I.
(b) An element that tends to lose two electrons belongs to group 2: Be, Mg, Ca etc.
(c) An element that tends to gain two electrons belongs to group 16: O, S, Se.
(d) Group 17 contains metals, non-metals, a liquid, and a gas.
28. The increasing order of reactivity among group 1 elements is Li < Na < K < Rb < Cs, whereas that among group 17 elements is F > Cl > Br > I. Explain.
Ans: In group 1, as the size increases down the group, ionization energy decreases and reactivity increases. In group 17, reactivity decreases down the group since larger atoms gain electrons less easily.
29. Write the general outer electronic configuration of s, p, d and f-block elements.
Ans: s-block: ns¹–²; p-block: ns²np¹–⁶; d-block: (n–1)d¹–¹⁰ns⁰–²; f-block: (n–2)f¹–¹⁴(n–1)d¹–¹⁰ns².
30. Assign the position of the element having outer electronic configuration (i) ns²np⁴, n=3 (ii) (n–1)d²ns², n=4 (iii) (n–2)f⁷(n–1)d¹ns², n=6 in the Periodic Table.
Ans: (i) Third period, p-block, group 16 (for example, sulfur).
(ii) Fourth period, d-block, group 4 (Titanium).
(iii) Sixth period, f-block, group 3 (Gadolinium).
31. From given ionization enthalpies and electron gain enthalpy, identify:
(a) The least reactive element is the one with highest ionization enthalpy and positive electron gain enthalpy.
(b) The most reactive metal is the one with lowest first ionization enthalpy and a slightly negative electron gain enthalpy.
(c) The most reactive nonmetal is the one with high first ionization enthalpy and largest negative electron gain enthalpy.
(d) The least reactive nonmetal is the one with very high first ionization enthalpy and positive electron gain enthalpy.
(e) The metal forming MX₂ is the one with lowest second ionization enthalpy.
(f) The metal forming a stable covalent halide MX has low first and high second ionization enthalpy.
32. Predict the formula of the stable binary compounds that would be formed by the combination of the following pairs of elements:
(a) Lithium and Oxygen: Li₂O
(b) Magnesium and Nitrogen: Mg₃N₂
(c) Aluminium and Iodine: AlI₃
(d) Silicon and Oxygen: SiO₂
(e) Phosphorus and Fluorine: PF₃ or PF₅
(f) Element 71 and Fluorine: LuF₃
33. In the modern periodic table, the period indicates the value of— (a) Atomic number (b) Atomic mass (c) Principal quantum number (d) Azimuthal quantum number.
Ans: It indicates the principal quantum number (c) of the valence shell.
34. Which of the following statements related to the modern periodic table is incorrect?
Ans: The statement "The d-block has 8 columns, because a maximum of 8 electrons can occupy all the orbitals in a d-subshell" is incorrect. The d-block actually has 10 columns because a d-subshell can accommodate 10 electrons.
35. Anything that influences the valence electrons will affect the chemistry of the element. Which one of the following factors does not affect the valence shell?
Ans: Nuclear mass does not affect the valence shell of an atom.
36. Which one of the following statements is incorrect in relation to ionization enthalpy?
Ans: "It is easier to remove electrons from orbitals with lower n values than from an orbital having higher n value" is incorrect. Removing electrons from lower n (closer to nucleus) is harder.
37. Considering the elements B, Al, Mg and K, the correct order of their metallic character is:
Ans: K > Mg > Al > B.
38. Considering the elements B, C, N, F and Si, the correct order of their non-metallic character is:
Ans: F > N > C > B > Si.
39. Considering the elements F, Cl, O and N, the correct order of their chemical reactivity in terms of oxidizing property is:
Ans: F > O > Cl > N.
40. (Any additional direct fact recall here as per your PDF can be added; just send the exact question.)
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