Classification of Elements and Periodicity in Properties Class 11 Notes

• Mendeleev predicted the properties of unknown elements (Ga, Ge, Sc) from gaps in his table.
• Moseley's X-ray work fixed the ordering by nuclear charge, removing the atomic-mass anomalies.
• The modern table has 118 elements in 18 groups and 7 periods.

• Group = number of valence electrons (or, for d/f, related to the outer configuration).
• Period = the principal quantum number of the outermost shell.
• Elements with the same valence-shell configuration share similar chemical behaviour within a group.

• Cations are smaller than their parent atoms; anions are larger.
• Isoelectronic species (same electron count, e.g. Na⁺, Mg²⁺, Al³⁺, O²⁻, F⁻) shrink as charge increases.
• d- and f-orbital contraction: the lanthanide contraction makes post-lanthanide elements smaller than otherwise expected.

• Across a period IE rises: nuclear charge grows and radius shrinks.
• Down a group IE falls: bigger atom, better shielding, outer electron held weakly.
• IE₂ > IE₁ > IE₃ always — removing a further electron from a more positive ion needs more energy.
• Exceptions: Be (1s²2s²) shows higher IE₁ than B (2s²2p¹); N (2p³, half-filled) beats O (2p⁴).

• Valence electrons define the combining ratio and the oxides formed: metals give basic oxides, non-metals acidic oxides.
• Metallic character decreases across a period and increases down a group.
• Reducing power follows ionisation enthalpy; oxidising power follows electron gain enthalpy.
• Noble gases are the least reactive — high IE plus a stable closed shell (except Kr and Xe which form compounds with F and O).

Example: Arrange N, O and F in order of increasing ionisation enthalpy.

Solution: F has the highest IE (small size, high Z), then N, then O. Nitrogen's IE₁ is higher than oxygen's despite being to its left, because the half-filled 2p³ configuration is extra stable.

Example: Which is smaller — Na⁺ or F⁻ — and why?

Solution: Na⁺ and F⁻ are isoelectronic ions (10 electrons each). For the same electron count, the ion with the higher nuclear charge pulls the electrons in tighter, so Na⁺ is smaller than F⁻.

Exercises (NCERT Q3.1–Q3.40)

• 3.1) What is the basic theme of organisation in the periodic table?
• 3.2) Which important property did Mendeleev use to classify elements, and did he always follow it?
• 3.3) What is the basic difference in approach between Mendeleev's periodic law and the modern periodic law?
• 3.4) On the basis of quantum numbers, justify that the sixth period of the periodic table should contain 32 elements.
• 3.5) In terms of period and group, where is the element with atomic number Z = 114 located?
• 3.6) Write the atomic number of the element in the third period and group 17 of the periodic table.
• 3.7) Which elements do you think would have been named by (i) Lawrence Berkeley Laboratory and (ii) Seaborg's group?
• 3.8) Why do elements in the same group have similar physical and chemical properties?
• 3.9) What do atomic radius and ionic radius mean?
• 3.10) How does atomic radius vary across a period and down a group, and why?
• 3.11) What are isoelectronic species? Name one species isoelectronic with each of (i) F⁻, (ii) Ar, (iii) Mg²⁺ and (iv) Rb⁺.
• 3.12) Consider N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺ and Al³⁺. (a) What is common to all of them? (b) Arrange them in order of increasing ionic radius.
• 3.13) Explain why a cation is smaller and an anion larger than its parent atom.
• 3.14) Why are the terms 'isolated gaseous atom' and 'ground state' used when defining ionization enthalpy and electron gain enthalpy? Hint: consider the requirements for comparison.
• 3.15) The energy of an electron in the ground state of a hydrogen atom is −2.18 × 10⁻¹⁸ J. Calculate the ionization enthalpy of atomic hydrogen in J mol⁻¹. Hint: use the mole concept.
• 3.16) The actual first ionization enthalpies of the second-period elements follow Li < B < Be < C < O < N < F < Ne. Why is (i) the value for Be higher than that for B, and (ii) the value for O lower than those for both N and F?
• 3.17) Explain why the first ionization enthalpy of sodium is lower than that of magnesium, but its second ionization enthalpy is higher.
• 3.18) What factors cause the ionization enthalpy of main-group elements to decrease down a group?
• 3.19) The exercise gives first ionization enthalpies for group 13 as B = 801, Al = 577, Ga = 579, In = 558 and Tl = 589 kJ mol⁻¹. How is this deviation from the general down-group decrease explained?
• 3.20) Which pair in each case has the more negative electron gain enthalpy: (i) O or F and (ii) F or Cl?
• 3.21) Would oxygen's second electron gain enthalpy be positive, more negative or less negative than its first? Justify the answer.
• 3.22) What is the basic difference between electron gain enthalpy and electronegativity?
• 3.23) How should the statement be evaluated that nitrogen has Pauling electronegativity 3.0 in all nitrogen compounds?
• 3.24) Describe the theory associated with atomic radius when (a) an atom gains an electron and (b) an atom loses an electron.
• 3.25) Would the first ionization enthalpies of two isotopes of the same element be the same or different? Justify the answer.
• 3.26) What are the major differences between metals and non-metals? Consider: (i) electron loss or gain, (ii) ionization and electron-gain tendencies, (iii) electronegativity, (iv) type of compounds, (v) oxide acidity or basicity, (vi) reducing or oxidizing action, and (vii) physical properties.
• 3.27) Use the periodic table to answer: (a) identify an element with five electrons in its outer subshell, (b) identify an element that tends to lose two electrons, (c) identify an element that tends to gain two electrons, and (d) identify the group containing metallic as well as non-metallic character and liquid as well as gaseous members at room temperature.
• 3.28) Explain why reactivity increases down group 1 in the order Li < Na < K < Rb < Cs, whereas halogen oxidizing reactivity decreases in the order F > Cl > Br > I.
• 3.29) Write the general outer electronic configurations of s-, p-, d- and f-block elements.
• 3.30) Assign the position in the periodic table for an element with outer configuration: (i) ns²np⁴ with n = 3, (ii) (n − 1)d²ns² with n = 4, and (iii) (n − 2)f⁷(n − 1)d¹ns² with n = 6.
• 3.31) The exercise gives the following first ionization enthalpy, second ionization enthalpy and electron gain enthalpy values, all in kJ mol⁻¹: I = 520, 7300, −60; II = 419, 3051, −48; III = 1681, 3374, −328; IV = 1008, 1846, −295; V = 2372, 5251, +48; VI = 738, 1451, −40. Which element is (a) least reactive, (b) most reactive metal, (c) most reactive non-metal, (d) least reactive non-metal, (e) the metal forming a stable binary halide MX₂, and (f) the metal forming a predominantly stable covalent halide MX, where X is a halogen?
• 3.32) Predict the formulas of stable binary compounds formed by: (a) lithium and oxygen, (b) magnesium and nitrogen, (c) aluminium and iodine, (d) silicon and oxygen, (e) phosphorus and fluorine, and (f) element 71 and fluorine.
• 3.33) In the modern periodic table, a period indicates the value of: (a) atomic number, (b) atomic mass, (c) principal quantum number or (d) azimuthal quantum number.
• 3.34) Which statement about the modern periodic table is incorrect? (a) The p-block has 6 columns because a p subshell can hold 6 electrons. (b) The d-block has 8 columns because a d subshell can hold 8 electrons. (c) Each block has as many columns as the maximum number of electrons in the corresponding subshell. (d) The block identifies the azimuthal quantum number l of the last subshell filled.
• 3.35) Which factor does not affect the valence shell: (a) valence principal quantum number n, (b) nuclear charge Z, (c) nuclear mass or (d) number of core electrons?
• 3.36) The size of the isoelectronic species F⁻, Ne and Na⁺ is affected by: (a) nuclear charge Z, (b) valence principal quantum number n, (c) electron–electron interaction in outer orbitals, or (d) none of these because their sizes are the same.
• 3.37) Which statement about ionization enthalpy is incorrect? (a) It increases for every successive electron. (b) The greatest increase occurs when an electron is removed from a core noble-gas configuration. (c) The end of valence electrons is marked by a large jump. (d) An electron from an orbital with lower n is easier to remove than one with higher n.
• 3.38) Considering B, Al, Mg and K, choose the correct order of metallic character: (a) B > Al > Mg > K, (b) Al > Mg > B > K, (c) Mg > Al > K > B, or (d) K > Mg > Al > B.
• 3.39) Considering B, C, N, F and Si, choose the correct order of non-metallic character: (a) B > C > Si > N > F, (b) Si > C > B > N > F, (c) F > N > C > B > Si, or (d) F > N > C > Si > B.
• 3.40) Considering F, Cl, O and N, choose the correct order of oxidizing property: (a) F > Cl > O > N, (b) F > O > Cl > N, (c) Cl > F > O > N, or (d) O > F > N > Cl.

Answers at a glance

1. 1
   Q3.1 → The basic theme is the classification of elements into periods and groups according to recurring physical and chemical properties.
2. 2
   Q3.2 → Mendeleev used increasing atomic weight, but he departed from strict atomic-weight order when a more chemically coherent family order was required.
3. 3
   Q3.3 → Mendeleev used atomic weight, whereas the modern periodic law uses the more fundamental and unambiguous atomic number.
4. 4
   Q3.4 → The sixth period has the theoretical capacity for 32 elements because its 16 orbitals can accommodate 32 electrons.
5. 5
   Q3.5 → The element with Z = 114 is flerovium in period 7 and group 14.
6. 6
   Q3.6 → The element is chlorine with atomic number 17.
7. 7
   Q3.7 → (i) Lawrencium (Lr), Z = 103, and berkelium (Bk), Z = 97; (ii) seaborgium (Sg), Z = 106.
8. 8
   Q3.8 → Elements in a group have similar valence-shell electron configurations, so they form bonds and exhibit related physical and chemical properties.
9. 9
   Q3.9 → Atomic radius measures an atom through metallic or covalent bond distances, while ionic radius measures a cation or anion in an ionic lattice; the exercise's examples give Cu = 128 pm, Cl = 99 pm, Na⁺ = 95 pm, Na = 186 pm, F⁻ = 136 pm and F = 64 pm.
10. 10
    Q3.10 → Atomic radius generally decreases across a period because effective nuclear charge increases, and it increases down a group because shell number and shielding increase.
11. 11
    Q3.11 → Isoelectronic species have equal electron counts: examples are Ne for F⁻, Cl⁻ for Ar, F⁻ for Mg²⁺ and Br⁻ for Rb⁺.
12. 12
    Q3.12 → All six ions are isoelectronic with 10 electrons, and their radii increase in the order Al³⁺ < Mg²⁺ < Na⁺ < F⁻ < O²⁻ < N³⁻.
13. 13
    Q3.13 → A cation contracts because electron loss raises effective nuclear attraction, whereas an anion expands because electron gain increases repulsion and reduces attraction per electron.
14. 14
    Q3.14 → The terms specify a minimum-interaction gaseous atom in its lowest-energy state, allowing enthalpy changes to represent atomic energetics and be compared consistently.
15. 15
    Q3.15 → The ionization enthalpy of atomic hydrogen is 1.31 × 10⁶ J mol⁻¹.
16. 16
    Q3.16 → Be's 2s electron is more tightly bound than B's 2p electron, while O's paired 2p electron is easier to remove than N's unpaired electron and F's increased effective nuclear charge binds its 2p electron more strongly.
17. 17
    Q3.17 → Na has the lower first ionization enthalpy because Mg binds 3s electrons more strongly, but Na⁺ has the higher second ionization enthalpy because its second electron is removed from a stable noble-gas core.
18. 18
    Q3.18 → Increasing shell number, increasing atomic size and stronger shielding by inner-shell electrons cause the down-group decrease in ionization enthalpy.
19. 19
    Q3.19 → The small Al-to-Ga and In-to-Tl increases arise from poor shielding by inserted d and f electrons, whereas the B-to-Al and Ga-to-In decreases follow the normal size-and-shielding trend.
20. 20
    Q3.20 → F has the more negative electron gain enthalpy than O, while Cl has the more negative electron gain enthalpy than F.
21. 21
    Q3.21 → The second electron gain enthalpy of oxygen is positive because strong repulsion in the small O⁻ ion makes addition of a second electron endothermic.
22. 22
    Q3.22 → Electron gain enthalpy is the enthalpy change for electron addition to an isolated gaseous atom, while electronegativity is the ability of an atom in a compound to attract a shared electron pair.
23. 23
    Q3.23 → The statement is incorrect because the effective electronegativity of nitrogen depends on its bonded environment and can differ between compounds.
24. 24
    Q3.24 → An anion is larger than its parent atom because added electrons increase repulsion, whereas a cation is smaller because electron loss increases effective nuclear attraction.
25. 25
    Q3.25 → The first ionization enthalpies are expected to be the same for isotopes because they have identical nuclear charge and electronic structure.
26. 26
    Q3.26 → Metals generally lose electrons, form cations and ionic compounds, have basic oxides and reducing character, while non-metals generally gain electrons, form covalent compounds, have acidic oxides and oxidizing character.
27. 27
    Q3.27 → (a) Br, for example; (b) Mg, for example; (c) O or S; (d) group 17.
28. 28
    Q3.28 → Metallic reactivity increases down group 1 as ionization enthalpy falls, whereas halogen reactivity decreases down group 17 as electron acceptance becomes less favourable, with F₂ remaining most reactive because of its low bond dissociation enthalpy.
29. 29
    Q3.29 → The general patterns are ns¹⁻² for s-block, ns²np¹⁻⁶ for p-block, (n − 1)d¹⁻¹⁰ns⁰⁻² for d-block and (n − 2)f¹⁻¹⁴(n − 1)d⁰⁻¹ns² for f-block elements.
30. 30
    Q3.30 → (i) S: period 3, group 16; (ii) Ti: period 4, group 4; (iii) Gd: period 6, f-block (group 3), Z = 64.
31. 31
    Q3.31 → (a) V; (b) II; (c) III; (d) V; (e) VI; (f) I.
32. 32
    Q3.32 → The formulas are (a) Li₂O, (b) Mg₃N₂, (c) AlI₃, (d) SiO₂, (e) PF₃ or PF₅, and (f) LuF₃.
33. 33
    Q3.33 → A period indicates the principal quantum number, so option (c) is correct.
34. 34
    Q3.34 → Statement (b) is incorrect: the d-block has 10 columns because a d subshell can accommodate 10 electrons.
35. 35
    Q3.35 → Nuclear mass, option (c), does not affect the valence shell in this periodic-trend treatment.
36. 36
    Q3.36 → The size is affected by nuclear charge, option (a), with the radius order Na⁺ < Ne < F⁻.
37. 37
    Q3.37 → Statement (d) is incorrect because an electron with higher n is easier to remove than one with lower n.
38. 38
    Q3.38 → The correct order is K > Mg > Al > B, option (d).
39. 39
    Q3.39 → The correct order is F > N > C > B > Si, option (c).
40. 40
    Q3.40 → The correct oxidizing-character order is F > O > Cl > N, option (b).

MCQ Questions

1 Mark Questions

1. 1
   Who proposed the modern periodic law? (Ans: Henry Moseley)
2. 2
   Write the group number of chlorine. (Ans: 17 (halogen))
3. 3
   Define electron gain enthalpy. (Ans: the energy released when an isolated gaseous atom gains an electron to form an anion — the more negative, the more readily the electron is gained)
4. 4
   What is the period number of sodium? (Ans: period 3)

2 Mark Questions

1. 1
   Explain why ionisation enthalpy decreases down a group. (Ans: increasing atomic size and greater shielding; the outermost electron is farther from the nucleus)
2. 2
   Give two anomalies of the Mendeleev periodic table that the modern table removed. (Ans: position of isotopes, misplaced elements by atomic mass (e.g. Ar before K), no place for noble gases until added)
3. 3
   Arrange F, Cl, Br and I in increasing order of electronegativity. (Ans: I < Br < Cl < F)

3 Mark Questions

1. 1
   Explain the periodic trends in atomic radius across a period and down a group with reasons. (Ans: decreases across due to increasing effective nuclear charge; increases down a group due to added shells)
2. 2
   Compare ionisation enthalpy with electron gain enthalpy and state how each varies along a period and a group.
3. 3
   Classify elements into s, p, d and f blocks and locate the position (group and period) of an element with configuration [Ar] 3d⁵ 4s¹. (Ans: d-block, group 6, period 4 — chromium)

5 Mark Questions

1. 1
   Describe Mendeleev's periodic table, its merits and limitations, and how it was replaced by the modern (Moseley) periodic law. (Ans: based on atomic mass, gaps predicted new elements, isotopes/order anomalies; modern law based on atomic number)
2. 2
   Discuss the trends of the following properties in the modern periodic table with reasons: atomic radius, ionisation enthalpy, electron gain enthalpy, electronegativity and metallic character — both across a period and down a group. (Ans: radius decreases/hypothetical; IE increases across decreases down; EGE becomes more negative across; EN increases across; metallic character decreases across and increases down)