Chapter 3: Classification of Elements and Periodicity in Properties

Link to Std 9–10 : You studied Mendeleev's table and the modern table briefly; now we see how the idea of "periodic" properties grew and why atomic number replaced atomic mass.

In the figure : Timeline of attempts at classification: Döbereiner (triads, 1829), Newlands (octaves, 1866), Mendeleev (1869), Moseley (atomic number, 1913).

Rule 1: Döbereiner's triads — Elements with similar properties come in threes, and the atomic mass of the middle one is about the average of the other two.

• Easy : Ca (40), Sr (88), Ba (137): average of Ca and Ba = 88.5, close to Sr.
• Medium : Cl (35.5), Br (80), I (127): average of Cl and I = 81.25, close to Br.
• Exam-type : Only a few triads were found, so the idea could not classify all known elements.

Rule 2: Newlands and Mendeleev — Newlands arranged elements by atomic mass and found that every eighth element was similar (law of octaves), but it failed beyond calcium. Mendeleev put elements in order of atomic mass, left gaps for undiscovered elements and predicted their properties (eka-aluminium became gallium, eka-silicon became germanium).

• Easy : Mendeleev's periodic law is based on atomic mass.
• Medium : His predictions for germanium matched the real element, which made his table widely accepted.
• Exam-type : Mendeleev placed some elements out of mass order (for example Te before I) to keep chemical similarity, a sign that atomic mass was not the true basis.

Rule 3: Moseley and the modern law — Moseley's X-ray work (1913) showed that atomic number, not atomic mass, is the basic property. The modern table has 7 periods and 18 groups.

• Easy : The modern periodic law is based on atomic number.
• Medium : With atomic number as the basis, Te (Z = 52) correctly comes before I (Z = 53).
• Exam-type : Isotopes of an element have different mass but the same Z, so they share one place in the table.

Rule 4: IUPAC names for Z above 100 — Use the digit roots: 0 nil, 1 un, 2 bi, 3 tri, 4 quad, 5 pent, 6 hex, 7 sept, 8 oct, 9 enn; add "ium" at the end. The symbol uses the first letter of each root.

• Easy : Z = 101 is un-nil-un-ium = unnilunium (Unu).
• Medium : Z = 120 is un-bi-nil-ium = unbinilium (Ubn).
• Exam-type : Z = 119 is un-un-enn-ium = ununennium (Uue).

Spot the difference

Link to Std 9–10 : You located elements by group and period; now we use electronic configuration to find the position without a table.

In the figure : Blocks of the modern table: s (groups 1–2), d (groups 3–12), p (groups 13–18) and f (two rows below).

Rule 1: Length of periods — Period 1 has 2 elements; periods 2 and 3 have 8 each; periods 4 and 5 have 18 each; period 6 has 32; period 7 is meant to hold 32.

• Easy : Period 2 (Li to Ne) has 8 elements.
• Medium : Period 4 has 18 because 4s, 3d and 4p (2 + 10 + 6) fill in turn.
• Exam-type : Lanthanoids (Z = 57–71) and actinoids (Z = 89–103) are placed below the main table as the f-block.

Rule 2: Finding group and period — s-block: group = number of valence electrons (ns¹ is group 1, ns² is group 2). p-block: group = 10 + number of electrons in the outermost shell (ns² np¹ is group 13). d-block: group = number of (n − 1)d electrons + number of ns electrons.

• Easy : Z = 20: [Ar] 4s², so period 4, group 2.
• Medium : Z = 35: [Ar] 3d¹⁰ 4s² 4p⁵, so period 4, group 10 + 7 = 17.
• Exam-type : Z = 26: [Ar] 3d⁶ 4s², so period 4, group 6 + 2 = 8 (d-block).

Rule 3: Metals, non-metals and metalloids — Metals (about three-quarters of the elements) lie on the left and centre, are good conductors and form cations. Non-metals lie at the upper right and form anions. Metalloids (B, Si, Ge, As, Sb, Te) lie along the zig-zag boundary and have intermediate properties.

• Easy : Na is a metal; Cl is a non-metal.
• Medium : Silicon is a metalloid and a semiconductor.
• Exam-type : Hydrogen is placed in group 1 for its ns¹ configuration, though it is not a metal at ordinary conditions.

Spot the difference

Link to Std 9–10 : You knew that atoms get bigger down a group; now we explain the trend using nuclear charge and shielding.

In the figure : Atomic radius (pm) across the third period: Na 186, Mg 160, Al 143, Si 117, P 110, S 104, Cl 99.

Rule 1: Trend in a period — Across a period (left to right) the nuclear charge increases but the electrons enter the same shell, so the pull increases and atomic radius decreases.

• Easy : Li 152 pm, Be 111 pm, B 88 pm, C 77 pm.
• Medium : Bond distance in Br₂ is 228 pm, so covalent radius of Br = 228 ÷ 2 = 114 pm.
• Exam-type : Distance between two adjacent sodium atoms in solid sodium is 372 pm, so metallic radius = 186 pm.

Rule 2: Trend in a group — Down a group a new shell is added and inner shells shield the nucleus, so atomic radius increases.

• Easy : Li 152, Na 186, K 231 pm.
• Medium : F 64, Cl 99, Br 114, I 133 pm.
• Exam-type : Noble gas radii are not compared with covalent radii, since they form no bonds and their van der Waals radii are larger.

Rule 3: Ionic radius — A cation is smaller than its atom (fewer electrons, same nuclear charge); an anion is larger (more electrons, more repulsion).

• Easy : Na 186 pm but Na⁺ 95 pm.
• Medium : F 64 pm but F⁻ 136 pm.
• Exam-type : Of Fe, Fe²⁺ and Fe³⁺ the order of size is Fe > Fe²⁺ > Fe³⁺.

Rule 4: Isoelectronic species — With the same number of electrons, a larger nuclear charge gives a smaller radius.

• Easy : O²⁻, F⁻, Na⁺, Mg²⁺ all have 10 electrons.
• Medium : Order of size: O²⁻ > F⁻ > Na⁺ > Mg²⁺.
• Exam-type : K⁺ and Cl⁻ both have 18 electrons; Cl⁻ (17 protons) is larger than K⁺ (19 protons).

Spot the difference

Link to Std 9–10 : You know metals lose and non-metals gain electrons; now we measure how easily.

In the figure : First ionisation enthalpy (kJ/mol) of the second period: Li 520, Be 899, B 801, C 1086, N 1402, O 1314, F 1681, Ne 2081.

Rule 1: Trends of ionisation enthalpy — It increases across a period (more nuclear charge, smaller size) and decreases down a group (larger size, more shielding). Successive ionisation enthalpies always increase: ΔᵢH₁ < ΔᵢH₂ < ΔᵢH₃.

• Easy : Na 496 kJ/mol is lower than Mg 737 kJ/mol.
• Medium : Li 520 is higher than Na 496 and K 419 (group 1 decreases).
• Exam-type : For an element, ΔᵢH values of 577, 1817, 2745 and 11 577 kJ/mol show a big jump after the third, so it has 3 valence electrons (group 13, aluminium).

Rule 2: Exceptions — Be (899) is higher than B (801) because a 2s electron is held more tightly than a 2p electron. N (1402) is higher than O (1314) because N has a stable half-filled 2p³, while in O a paired electron in 2p⁴ is easier to remove.

• Easy : B < Be in ionisation enthalpy.
• Medium : O < N in ionisation enthalpy.
• Exam-type : Likewise Al (577) is less than Mg (737).

Rule 3: Electron gain enthalpy — It is generally more negative across a period and less negative down a group. Halogens have large negative values; noble gases have positive values because the new electron must enter a new shell. Cl (−349 kJ/mol) is more negative than F (−328 kJ/mol), because the small 2p subshell of F suffers more repulsion.

• Easy : Cl −349, F −328, Br −325, I −295 kJ/mol.
• Medium : O −141 and S −200 kJ/mol: S is more negative than O.
• Exam-type : Ne +116 kJ/mol: positive, so Ne does not accept an electron.

Spot the difference

Link to Std 9–10 : You wrote formulas using valency; here we see why valency changes in a regular way across the table.

In the figure : Summary arrows: across a period atomic size falls, while ionisation enthalpy and electronegativity rise; down a group the opposite happens.

Rule 1: Trend of electronegativity (Pauling scale) — Increases across a period and decreases down a group.

• Easy : Period 2: Li 1.0, Be 1.5, B 2.0, C 2.5, N 3.0, O 3.5, F 4.0.
• Medium : Group 17: F 4.0, Cl 3.0, Br 2.8, I 2.5.
• Exam-type : The difference in electronegativity tells the bond type: H–Cl is 3.0 − 2.1 = 0.9 (polar covalent); Na–Cl is 3.0 − 0.9 = 2.1 (largely ionic); Cl–Cl is 0 (pure covalent).

Rule 2: Valence and oxides across period 3 — For main-group elements, the valence is the number of valence electrons (up to 4) or 8 minus that number. Oxides change from basic to acidic across a period: Na₂O and MgO are basic, Al₂O₃ is amphoteric, SiO₂, P₄O₁₀, SO₃ and Cl₂O₇ are acidic.

• Easy : Valence of Na is 1, of Mg 2, of Al 3.
• Medium : Valence of Cl is 8 − 7 = 1 in HCl.
• Exam-type : Mg + N → Mg₃N₂ (3 Mg²⁺ balance 2 N³⁻); Al + Cl → AlCl₃; Ca + O → CaO.

Rule 3: Chemical reactivity — Reactivity of metals increases down a group (easier to lose electrons) and falls across a period; reactivity of non-metals decreases down a group (harder to gain electrons) and increases across a period. Group 1: Li < Na < K < Rb < Cs. Group 17: F > Cl > Br > I.

• Easy : Potassium is more reactive than sodium.
• Medium : Fluorine is the most reactive halogen.
• Exam-type : Noble gases are least reactive; they have complete shells.

Rule 4: Anomalous behaviour of the second-period elements — Li, Be and B differ from the rest of their groups because of small size, high ionisation enthalpy, and the absence of d-orbitals (maximum covalence 4). They resemble elements diagonally below: Li–Mg, Be–Al, B–Si (diagonal relationship).

• Easy : Li behaves like Mg.
• Medium : Be (like Al) forms amphoteric oxide.
• Exam-type : Boron cannot have covalence 6, but aluminium can (AlF₆³⁻), as Al has vacant 3d orbitals.

Spot the difference