Classification of Elements and Periodicity in Properties
Welcome to HSLC Guru. This article presents a complete English-medium guide to Class 11 Chemistry Chapter 3 — Classification of Elements and Periodicity in Properties, prepared strictly for ASSEB students. You will find a clear summary, short and long answer questions, MCQs, fill in the blanks, true/false statements and a glossary table to help you revise quickly and score better in your half-yearly and final exams.
Summary
The need to classify elements arose as new elements were continually discovered. Early attempts included Dobereiner’s Triads (1817), in which sets of three chemically similar elements showed the atomic mass of the middle element as nearly the arithmetic mean of the other two (e.g., Li, Na, K). Newlands’ Law of Octaves (1865) arranged elements in increasing order of atomic mass and observed that every eighth element resembled the first, like musical notes; however, it failed beyond calcium. Mendeleev’s Periodic Table (1869) was the first widely accepted classification. He arranged elements in order of increasing atomic mass and stated his Periodic Law: “The properties of elements are a periodic function of their atomic masses.” Mendeleev left gaps for undiscovered elements (eka-aluminium, eka-silicon) and predicted their properties accurately.
The discovery of isotopes and the work of Henry Moseley (1913) led to the Modern Periodic Law: “The physical and chemical properties of elements are a periodic function of their atomic numbers.” The atomic number (Z), being the number of protons, is a more fundamental property than atomic mass. The long form (modern) periodic table contains 7 horizontal periods and 18 vertical groups. Elements are classified into four blocks based on the subshell to which the last electron enters: s-block (Groups 1, 2 and He), p-block (Groups 13–18), d-block or transition elements (Groups 3–12), and f-block or inner transition elements (lanthanoids and actinoids). For elements with Z > 100, the IUPAC nomenclature uses numerical roots (nil = 0, un = 1, bi = 2, tri = 3, quad = 4, pent = 5, hex = 6, sept = 7, oct = 8, enn = 9) followed by “-ium”. Example: Z = 120 = Unbinilium (Ubn).
Periodic Trends: The effective nuclear charge experienced by the valence electrons explains many trends.
$Z_{\text{eff}} = Z – \sigma$
where Z is the atomic number and σ is the screening (shielding) constant. Atomic radius decreases across a period (due to increasing Zeff) and increases down a group (due to addition of new shells). Ionic radius: cations are smaller than the parent atom; anions are larger. For an isoelectronic series, the ion with the largest negative charge is the largest. Ionisation enthalpy (IE) is the minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.
$\text{M}(g) \to \text{M}^+(g) + e^-$
IE generally increases across a period and decreases down a group; exceptions occur due to extra stability of half-filled and fully-filled configurations (e.g., Be > B; N > O). Electron-gain enthalpy is the enthalpy change when an isolated gaseous atom accepts an electron to form a gaseous anion; it becomes more negative across a period and less negative down a group. Halogens have the most negative values; noble gases have positive values. Electronegativity (EN) is the tendency of an atom in a molecule to attract the shared pair of electrons; it increases across a period, decreases down a group, fluorine being the most electronegative (Pauling scale = 4.0). Valency equals the number of valence electrons (for groups 1, 2, 13) or 8 minus the number of valence electrons (for groups 14–17). Metallic character decreases across a period and increases down a group, while non-metallic character shows the opposite trend. Periodic chemical reactivity: reactivity of metals increases down a group and decreases across a period; reactivity of non-metals decreases down a group and increases across a period (up to group 17). The nature of oxides changes from basic (left) through amphoteric (middle) to acidic (right).
Very Short Answer Questions (1 Mark)
Q1. State the Modern Periodic Law.
Answer: The physical and chemical properties of elements are a periodic function of their atomic numbers.
Q2. Who proposed the Law of Octaves?
Answer: John Newlands (1865).
Q3. How many periods and groups are present in the long form of the periodic table?
Answer: 7 periods and 18 groups.
Q4. Name the most electronegative element in the periodic table.
Answer: Fluorine (F), with Pauling electronegativity 4.0.
Q5. Write the IUPAC name and symbol of the element with Z = 120.
Answer: Unbinilium (Ubn).
Q6. Which block do lanthanoids and actinoids belong to?
Answer: The f-block (inner transition elements).
Q7. What is the SI unit of ionisation enthalpy?
Answer: kJ mol−1.
Q8. Give one example of Dobereiner’s triad.
Answer: Lithium (Li, 7), Sodium (Na, 23), Potassium (K, 39); the atomic mass of Na is approximately the average of Li and K.
Q9. Define isoelectronic species with one example.
Answer: Species having the same number of electrons. Example: N3−, O2−, F−, Ne, Na+, Mg2+ (each has 10 electrons).
Q10. Which element has the highest first ionisation enthalpy?
Answer: Helium (He).
Short Answer Questions (2–3 Marks)
Q1. Distinguish between Mendeleev’s Periodic Law and the Modern Periodic Law.
Answer: Mendeleev’s Periodic Law states that the properties of elements are a periodic function of their atomic masses. The Modern Periodic Law (Moseley) states that the properties of elements are a periodic function of their atomic numbers. The change was necessary because atomic number (number of protons) is a more fundamental property than atomic mass; it also resolved the anomalies of pairs like Ar–K, Co–Ni and Te–I.
Q2. Why does atomic radius decrease across a period but increase down a group?
Answer: Across a period, electrons are added to the same shell while the nuclear charge increases. The effective nuclear charge (Zeff) experienced by valence electrons rises, pulling them closer; hence atomic radius decreases. Down a group, new principal shells are added; the increased shielding outweighs the rise in nuclear charge, so the outer electrons are farther from the nucleus and the radius increases.
Q3. Arrange the following in increasing order of ionic radii: N3−, O2−, F−, Na+, Mg2+.
Answer: All five are isoelectronic (10 electrons). The smaller the positive charge or the greater the negative charge, the larger the ionic radius. Order: Mg2+ < Na+ < F− < O2− < N3−.
Q4. First ionisation enthalpy of nitrogen is greater than that of oxygen. Explain.
Answer: Nitrogen has the configuration [He] 2s2 2p3, in which the 2p subshell is exactly half-filled — an extra-stable arrangement. Oxygen has [He] 2s2 2p4; removing the paired 2p electron is easier because doing so relieves electron–electron repulsion and gives a stable half-filled p3 configuration. Hence IE1(N) > IE1(O).
Q5. Why are noble gases placed in a separate group?
Answer: Noble gases (group 18) have completely filled valence shells (ns2np6, except He = 1s2). They are chemically inert under ordinary conditions, possess very high ionisation enthalpies and nearly zero electron-gain enthalpies. Their unique chemistry justifies a separate group.
Q6. Define electronegativity. State Pauling’s scale.
Answer: Electronegativity is the tendency of an atom in a covalent bond to attract the shared pair of electrons towards itself. On Pauling’s scale, fluorine is assigned the highest value of 4.0. Other commonly cited values are O = 3.5, N = 3.0, Cl = 3.0, C = 2.5.
Long Answer Questions (5–7 Marks)
Q1. Discuss the merits and demerits of Mendeleev’s periodic table.
Answer: Merits: (i) It systematically arranged 63 known elements in order of increasing atomic mass. (ii) Vacant places were left for undiscovered elements such as eka-boron (Sc), eka-aluminium (Ga) and eka-silicon (Ge); their properties were predicted with remarkable accuracy. (iii) Elements with similar properties were placed in the same group, simplifying the study of their chemistry. (iv) Atomic masses of some elements (e.g., Be, In, U) were corrected on the basis of the table.
Demerits: (i) The position of hydrogen was not definite — it resembles both alkali metals and halogens. (ii) Isotopes of an element could not be accommodated, as they have the same chemical properties but different atomic masses. (iii) Anomalous pairs such as Ar (40)–K (39), Co (58.9)–Ni (58.7) and Te (127.6)–I (126.9) had to be placed against the trend of atomic mass. (iv) Lanthanoids and actinoids could not be placed properly. (v) The cause of periodicity was not explained.
Q2. Describe the classification of elements into s-, p-, d- and f-blocks. Give general electronic configurations and characteristic properties of each block.
Answer: Based on the subshell that receives the differentiating electron, the periodic table is divided into four blocks.
(a) s-block: Groups 1 (alkali metals) and 2 (alkaline earth metals). General configuration: ns1–2. They are soft, highly reactive metals, strong reducing agents, form ionic compounds, low IE, low EN, basic oxides.
(b) p-block: Groups 13 to 18. General configuration: ns2 np1–6. They include metals, metalloids and non-metals; show variable oxidation states; halogens and chalcogens are the most reactive non-metals; group 18 contains noble gases.
(c) d-block (transition elements): Groups 3 to 12. General configuration: (n−1)d1–10 ns0–2. They are hard, high-melting metals, exhibit variable oxidation states, form coloured ions and complexes, and act as catalysts.
(d) f-block (inner transition elements): Lanthanoids (Ce–Lu) and actinoids (Th–Lr). General configuration: (n−2)f1–14 (n−1)d0–1 ns2. They show lanthanoid contraction, multiple oxidation states; actinoids are radioactive.
Q3. Explain the variation of ionisation enthalpy along a period and down a group, with reasons and exceptions.
Answer: Ionisation enthalpy (ΔiH) is the energy required to remove the most loosely bound electron from one mole of isolated gaseous atoms.
Across a period: Nuclear charge increases, atomic radius decreases and Zeff rises. Therefore, IE generally increases from left to right. Maximum value occurs at the noble gas. Exceptions: (i) IE1(B) < IE1(Be) because removal of the 2p electron of B is easier than the paired 2s electron of Be. (ii) IE1(O) < IE1(N) because N has an extra-stable half-filled 2p3 configuration.
Down a group: Atomic size increases and shielding by inner shells increases, so the outermost electron is loosely held. IE decreases from top to bottom. Hence, Cs has one of the lowest IE values among stable elements (Fr being radioactive).
Q4. Define electron-gain enthalpy. Discuss its variation in the periodic table and explain why the value of ΔegH of fluorine is less negative than that of chlorine.
Answer: Electron-gain enthalpy (ΔegH) is the enthalpy change when an isolated gaseous atom in its ground state accepts an electron to form a gaseous negative ion: X(g) + e− → X−(g).
Variation: Across a period, ΔegH becomes more negative (energy released increases) as atomic size decreases and Zeff rises. Halogens have the most negative values. Down a group, ΔegH becomes less negative because added shells increase the size and reduce the attraction for the incoming electron.
Why ΔegH(F) is less negative than ΔegH(Cl): Fluorine has a very small 2p subshell. The incoming electron experiences strong inter-electronic repulsion in the compact 2p orbitals, so less energy is released. Chlorine, with a larger 3p subshell, accommodates the incoming electron more comfortably, giving a more negative ΔegH. Hence chlorine has the most negative electron-gain enthalpy in the periodic table.
Q5. Write the IUPAC nomenclature for elements with Z > 100. Derive the names and symbols for Z = 101, 109, 116 and 118.
Answer: The IUPAC system uses numerical roots: 0 = nil (n), 1 = un (u), 2 = bi (b), 3 = tri (t), 4 = quad (q), 5 = pent (p), 6 = hex (h), 7 = sept (s), 8 = oct (o), 9 = enn (e). The roots for the digits of the atomic number are written in order, followed by the suffix “-ium”. The symbol is formed from the first letter of each root.
Z = 101 → un + nil + un + ium = Unnilunium, symbol Unu (now Mendelevium, Md).
Z = 109 → un + nil + enn + ium = Unnilennium, symbol Une (now Meitnerium, Mt).
Z = 116 → un + un + hex + ium = Ununhexium, symbol Uuh (now Livermorium, Lv).
Z = 118 → un + un + oct + ium = Ununoctium, symbol Uuo (now Oganesson, Og).
Multiple Choice Questions (MCQs)
Q1. The Modern Periodic Law was proposed by:
(a) Mendeleev (b) Newlands (c) Moseley (d) Dobereiner
Answer: (c) Moseley.
Q2. The element with the highest first ionisation enthalpy is:
(a) H (b) He (c) Ne (d) F
Answer: (b) He.
Q3. The most electronegative element on the Pauling scale is:
(a) Cl (b) O (c) F (d) N
Answer: (c) F.
Q4. Which one of the following is an isoelectronic pair?
(a) Na+, K+ (b) O2−, F− (c) Cl−, Br− (d) Mg2+, Ca2+
Answer: (b) O2−, F− (both have 10 electrons).
Q5. The total number of groups in the modern periodic table is:
(a) 7 (b) 8 (c) 16 (d) 18
Answer: (d) 18.
Q6. The element with electronic configuration [Ar] 3d10 4s2 4p3 belongs to:
(a) s-block (b) p-block (c) d-block (d) f-block
Answer: (b) p-block (As, Z = 33).
Q7. Among the following, the element with the largest atomic radius is:
(a) Li (b) Na (c) K (d) Cs
Answer: (d) Cs.
Q8. The IUPAC symbol for element Z = 120 is:
(a) Ubn (b) Uun (c) Uno (d) Ueb
Answer: (a) Ubn (Unbinilium).
Q9. The element with the most negative electron-gain enthalpy is:
(a) F (b) Cl (c) O (d) S
Answer: (b) Cl.
Q10. Lanthanoids and actinoids belong to:
(a) s-block (b) p-block (c) d-block (d) f-block
Answer: (d) f-block.
Fill in the Blanks
Q1. The modern periodic law was given by __________.
Answer: Henry Moseley.
Q2. The general electronic configuration of p-block elements is __________.
Answer: ns2 np1–6.
Q3. The effective nuclear charge is given by Zeff = __________.
Answer: Z − σ.
Q4. The atomic radius __________ across a period from left to right.
Answer: decreases.
Q5. The element with the highest electronegativity is __________.
Answer: Fluorine.
True or False
Q1. Mendeleev arranged elements in the order of increasing atomic number.
Answer: False. He used atomic mass; atomic number was used by Moseley.
Q2. Anion is always larger than its parent neutral atom.
Answer: True.
Q3. Ionisation enthalpy increases down a group.
Answer: False. It decreases down a group.
Q4. Noble gases have positive electron-gain enthalpy values.
Answer: True.
Q5. The element with Z = 118 is named Oganesson.
Answer: True.
Glossary
| Term | Meaning |
|---|---|
| Periodic Law | Properties of elements are a periodic function of their atomic numbers (Moseley). |
| Period | A horizontal row in the periodic table; there are 7 periods. |
| Group | A vertical column in the periodic table; there are 18 groups. |
| s-block | Elements of groups 1 and 2 (and He); valence electron enters the s subshell. |
| p-block | Elements of groups 13–18; valence electron enters the p subshell. |
| d-block | Transition elements of groups 3–12; differentiating electron enters (n−1)d. |
| f-block | Lanthanoids and actinoids; differentiating electron enters (n−2)f. |
| Atomic radius | Distance from the nucleus to the outermost shell of an isolated atom. |
| Ionic radius | Effective radius of an ion in an ionic crystal. |
| Isoelectronic species | Species having the same number of electrons. |
| Ionisation enthalpy | Energy needed to remove the most loosely bound electron from a gaseous atom. |
| Electron-gain enthalpy | Enthalpy change when a gaseous atom accepts an electron. |
| Electronegativity | Tendency of an atom in a bond to attract the shared pair of electrons. |
| Effective nuclear charge | Net positive charge experienced by valence electrons (Zeff = Z − σ). |
| Shielding effect | Reduction in the effective nuclear charge due to inner-shell electrons. |
| Valency | Combining capacity of an element; for s/p-block, related to the number of valence electrons. |
| Metallic character | Tendency of an element to lose electrons; increases down a group, decreases across a period. |
| Non-metallic character | Tendency to gain electrons; decreases down a group, increases across a period. |
| IUPAC nomenclature (Z > 100) | Systematic naming using numerical roots and the suffix “-ium”. |
| Lanthanoid contraction | Steady decrease in atomic and ionic size across the lanthanoid series. |