Classification of Elements and Periodicity in Properties: NEET notes

1. Why classify, and the first attempts

NCERT §3.1, §3.2

• Only 31 elements were known in 1800 and 63 by 1865; today elements up to Z = 118 are known and named. Studying each one and its compounds separately is impractical, so chemists looked for a system that would organise known facts and predict new ones.
• Dobereiner (German, by 1829) found groups of three similar elements, triads, in which the middle element's atomic weight is roughly the average of the other two and its properties lie in between: Li 7, Na 23, K 39; Ca 40, Sr 88, Ba 137; Cl 35.5, Br 80, I 127.
• Check a triad: (7 + 39)/2 = 23 for Na, and (35.5 + 127)/2 = 81.25, close to Br's 80. The law of triads held for only a few elements and was set aside as coincidence.
• de Chancourtois, a French geologist, arranged the known elements by increasing atomic weight on a cylinder in 1862 to show recurring properties; it drew little attention.
• Newlands (English, 1865), Law of Octaves: listed by atomic weight, every eighth element resembles the first, like the eighth note of a musical scale. It worked only up to calcium, but the Royal Society later gave him the Davy Medal (1887).
• Lothar Meyer (German) graphed physical properties such as atomic volume and the melting and boiling points against atomic weight. The pattern repeated, but its length grew along the list. His 1868 table resembled the modern one closely, yet appeared in print only after Mendeleev's.

2. Mendeleev's periodic table

NCERT §3.2

• Mendeleev (Russian, 1834-1907) and Lothar Meyer, working independently, both proposed in 1869 that properties recur at regular intervals when elements are listed by atomic weight. Mendeleev is credited with publishing the Periodic Law first.
• Mendeleev's Periodic Law: the properties of elements are a periodic function of their atomic weights.
• He placed elements in rows and columns so that similar ones fell in the same vertical group, and used a wide range of properties, especially the formulas and behaviour of their compounds.
• He broke atomic-weight order when properties demanded it, suspecting faulty weights: iodine, lighter than tellurium, was put in Group VII with fluorine, chlorine and bromine.
• He left gaps for undiscovered elements and predicted their properties: eka-aluminium (later gallium) under aluminium and eka-silicon (later germanium) under silicon.
• Predicted vs found, eka-aluminium/gallium: atomic weight 68 vs 70, density 5.9 vs 5.94 g/cm³, low melting point vs 302.93 K, oxide E₂O₃ vs Ga₂O₃, chloride ECl₃ vs GaCl₃.
• Predicted vs found, eka-silicon/germanium: atomic weight 72 vs 72.6, density 5.5 vs 5.36 g/cm³, high melting point vs 1231 K, oxide EO₂ vs GeO₂, chloride ECl₄ vs GeCl₄. The success of these bold numerical forecasts made his table famous.

3. Modern periodic law and the long form

NCERT §3.3

• In 1913 Moseley (English physicist) found that √ν of an element's characteristic X-rays gives a straight line against atomic number Z, not against atomic mass. Z is therefore the more fundamental property.
• Modern Periodic Law: the physical and chemical properties of elements are periodic functions of their atomic numbers.
• Z equals the number of protons, which equals the number of electrons in a neutral atom. Periodicity is really the repeating pattern of electronic configurations, which fix the properties of elements and their compounds.
• Nature supplies 94 elements. Pitchblende, a uranium ore, contains traces of neptunium and plutonium as well as actinium and protactinium.
• The long form is the most widely used table. Horizontal rows are periods; vertical columns, holding elements with similar outer configurations, are groups or families. IUPAC numbers the groups 1 to 18, replacing IA…VIIA, VIII, IB…VIIB and 0.
• There are seven periods, and the period number is the highest principal quantum number n present. Period lengths are 2, 8, 8, 18, 18 and 32; the seventh, like the sixth, can hold at most 32.
• The 14 lanthanoids (period 6) and 14 actinoids (period 7) sit in two separate rows at the bottom. Seaborg's work on plutonium (1940) and the transuranium elements 94-102 led to this layout; he received the 1951 Nobel Prize and element 106 is named seaborgium (Sg).

4. Naming elements with Z above 100

NCERT §3.4

• New heavy elements are so unstable that sometimes only a few atoms are made, in a handful of laboratories. Rival claims followed: for element 104 the Americans proposed rutherfordium and the Soviets kurchatovium.
• IUPAC's rule: until a discovery is confirmed and named, the element gets a systematic name built from its digits. Roots: 0 nil, 1 un, 2 bi, 3 tri, 4 quad, 5 pent, 6 hex, 7 sept, 8 oct, 9 enn.
• Join the roots in the order of the digits and add -ium. The temporary symbol has three letters, the first letter of each root, with the first capitalised.
• Example: Z = 120 → un + bi + nil + ium = unbinilium, symbol Ubn. Z = 101 → unnilunium (Unu); Z = 118 → ununoctium (Uuo).
• A final name and symbol follow a vote among IUPAC delegates of the member countries. They may recall the place of discovery (a country or one of its states) or honour a well-known scientist.
• Elements up to Z = 118 have been discovered and all have official IUPAC names, for example 101 mendelevium (Md), 104 rutherfordium (Rf), 106 seaborgium (Sg), 113 nihonium (Nh), 114 flerovium (Fl), 117 tennessine (Ts) and 118 oganesson (Og).

5. Electronic configuration, periods and groups

NCERT §3.5

• Where an element sits in the table reflects the quantum numbers of the last orbital to be filled.
• The period number is n of the outermost (valence) shell. Each period holds twice as many elements as there are orbitals being filled in it.
• Period 1 (1s): H and He, 2 elements. Period 2 fills 2s and 2p, Li to Ne, 8 elements. Period 3 fills 3s and 3p, Na to Ar, 8 elements.
• Period 4 begins at K (4s). Before 4p, the 3d orbitals fill: the 3d transition series runs from Sc (Z = 21, 3d¹4s²) to Zn (Z = 30, 3d¹⁰4s²). The period ends at Kr after 4p, 18 elements in all.
• Period 5 begins at Rb, contains the 4d series starting at Y (Z = 39) and ends at Xe after 5p: 18 elements, because 5s, 4d and 5p offer 9 orbitals, room for 18 electrons.
• Period 6 has 32 elements, filling 6s, 4f, 5d and 6p. The 4f series (lanthanoids) runs from Ce (Z = 58) to Lu (Z = 71).
• Period 7 fills 7s, 5f, 6d and 7p and holds most of the man-made radioactive elements; it closes with element 118, a noble gas. Filling 5f after Ac (Z = 89) gives the actinoid series.
• Elements in one group have the same number and arrangement of outer electrons, hence similar properties. Group 1 all end in ns¹: Li [He]2s¹, Na [Ne]3s¹, K [Ar]4s¹, Rb [Kr]5s¹, Cs [Xe]6s¹, Fr [Rn]7s¹.
• Conclusion: properties depend periodically on atomic number, not on relative atomic mass.

6. The s, p, d and f blocks

NCERT §3.6.1-3.6.4

• The block is named after the kind of orbital receiving the last electron. Two exceptions: helium (1s²) belongs to the s-block but sits with group 18 because its shell is full; hydrogen (1s¹) can lose an electron like group 1 or gain one like group 17, so it is placed on its own at the top.
• s-block: groups 1 (alkali metals, ns¹) and 2 (alkaline earth metals, ns²). Reactive metals with low ionization enthalpies, forming 1+ or 2+ ions; metallic character and reactivity rise down the group, so they never occur free. Their compounds are mainly ionic, except those of Li and Be.
• p-block: groups 13 to 18, outer configuration ns²np¹ to ns²np⁶. The s- and p-blocks together are the representative or main group elements.
• Each period ends in a noble gas, ns²np⁶; its full valence shell is hard to disturb, so reactivity is very low. Before them come the halogens (group 17) and chalcogens (group 16), with strongly negative electron gain enthalpies; they readily take one or two electrons to reach a noble gas configuration.
• d-block (transition elements): groups 3 to 12, filling inner d orbitals; general configuration (n−1)d¹⁻¹⁰ns⁰⁻², with Pd (4d¹⁰5s⁰) the exception. All are metals, mostly with coloured ions, variable oxidation states and paramagnetism, and are often catalysts.
• Zn, Cd and Hg, with (n−1)d¹⁰ns², miss most typical transition-metal features. Sitting between the very active s-block metals and the less active groups 13 and 14, the d-block acts as a bridge, which explains the name transition elements.
• f-block (inner transition elements): lanthanoids Ce (58) to Lu (71) and actinoids Th (90) to Lr (103), outer configuration (n−2)f¹⁻¹⁴(n−1)d⁰⁻¹ns². All are metals, and members of one series are very alike.
• Early actinoids have more complex chemistry than the matching lanthanoids because they show many oxidation states. Actinoids are radioactive, many made only in nanogram amounts or less; elements after uranium are the transuranium elements.
• Predicting from the table: Z = 117 falls in group 17, [Rn]5f¹⁴6d¹⁰7s²7p⁵; Z = 120 would be in group 2, [Uuo]8s².

7. Metals, non-metals and metalloids

NCERT §3.6.5

• Metals make up more than 78% of the known elements and lie on the left of the table. Non-metals, fewer than twenty, lie at the top right.
• Metals are usually solid at room temperature (mercury is liquid; gallium melts at 303 K and caesium at 302 K), with high melting and boiling points. They conduct heat and electricity and are malleable (beaten into sheets) and ductile (drawn into wires).
• Non-metals are mostly gases or solids that melt and boil at low temperatures (boron and carbon are exceptions). They conduct poorly, and the solid ones are mostly brittle, neither malleable nor ductile.
• Across a period, character changes from metallic to non-metallic; down a group, elements become more metallic.
• The change is gradual. Elements along the diagonal zig-zag line, such as silicon, germanium, arsenic, antimony and tellurium, show properties of both and are called metalloids or semi-metals.
• Worked order of increasing metallic character: P < Si < Be < Mg < Na (metallic character falls left to right and rises down a group).

8. Atomic and ionic radii

NCERT §3.7.1(a), (b)

• Periodic trends: across a period reactivity is high in group 1, lower towards the middle and highest in group 17. Down the alkali metals reactivity rises; down the halogens it falls. Trends are explained through nuclear charge and energy levels.
• An atom is tiny (about 1.2 Å = 1.2 × 10⁻¹⁰ m) and its electron cloud has no sharp edge, so its radius is estimated from distances between bonded atoms.
• Covalent radius (non-metals): half the single-bond length, e.g. Cl–Cl in Cl₂ is 198 pm, so r(Cl) = 99 pm. Metallic radius: half the distance between neighbouring metal cores, e.g. Cu–Cu 256 pm, so r(Cu) = 128 pm. NCERT calls either one the atomic radius.
• Across a period the radius falls: electrons enter the same shell while effective nuclear charge rises and pulls them in. Period 2 (pm): Li 152, Be 111, B 88, C 77, N 74, O 66, F 64. Period 3: Na 186, Mg 160, Al 143, Si 117, P 110, S 104, Cl 99.
• Down a group the radius grows: n increases, and filled inner shells shield the outer electrons. Group 1 (pm): Li 152, Na 186, K 231, Rb 244, Cs 262. Group 17: F 64, Cl 99, Br 114, I 133, At 140.
• Noble gases are left out: they are monoatomic, so only non-bonded (van der Waals) radii exist, and these are much larger than covalent radii.
• A cation is smaller than its atom (fewer electrons, same nuclear charge); an anion is larger (extra repulsion, lower effective nuclear charge). F 64 pm but F⁻ 136 pm; Na 186 pm but Na⁺ 95 pm.
• Isoelectronic species have the same number of electrons, e.g. O²⁻, F⁻, Na⁺ and Mg²⁺ (10 each). The more protons, the smaller the ion: O²⁻ > F⁻ > Na⁺ > Mg²⁺.
• Among Mg, Mg²⁺, Al and Al³⁺ the largest is Mg and the smallest is Al³⁺.

9. Ionization enthalpy

NCERT §3.7.1(c)

• Ionization enthalpy ΔᵢH is the energy needed to remove an electron from an isolated gaseous atom in its ground state: X(g) → X⁺(g) + e⁻. It is measured in kJ mol⁻¹ and, unqualified, means the first ionization enthalpy.
• It is always positive. The second, X⁺(g) → X²⁺(g) + e⁻, is larger than the first because the electron leaves a positive ion; the third is larger still.
• Plotted against Z (up to 60), ΔᵢH peaks at the noble gases (closed shells) and dips at the alkali metals, whose low values match their high reactivity. Generally it rises across a period and falls down a group, the reverse of atomic radius.
• Shielding: inner electrons screen the valence electron, so it feels an effective nuclear charge smaller than the full charge. Lithium's 2s electron, screened by 1s², feels much less than +3. Shielding works best when inner shells are complete.
• Across a period, electrons join the same shell and shielding barely grows, so rising nuclear charge wins and ΔᵢH rises. Down a group, the electron is farther out and better shielded, so ΔᵢH falls.
• Exception 1: ΔᵢH(B) < ΔᵢH(Be). Boron loses a 2p electron, which penetrates less and is more shielded than beryllium's 2s electron.
• Exception 2: ΔᵢH(O) < ΔᵢH(N). Nitrogen's three 2p electrons are unpaired in separate orbitals (Hund's rule); in oxygen two 2p electrons share an orbital, and their repulsion makes one easier to remove.
• Actual period 2 order: Li < B < Be < C < O < N < F < Ne. In period 3, with Na 496, Mg 737 and Si 786 kJ mol⁻¹, Al should be near 575 kJ mol⁻¹, below Mg, because its 3p electron is shielded by 3s electrons.
• Per atom to per mole: hydrogen's ground-state electron has E = −2.18 × 10⁻¹⁸ J, so ΔᵢH = 2.18 × 10⁻¹⁸ × 6.022 × 10²³ ≈ 1.31 × 10⁶ J mol⁻¹.

10. Electron gain enthalpy and electronegativity

NCERT §3.7.1(d), (e)

• Electron gain enthalpy ΔegH is the enthalpy change when a neutral gaseous atom takes an electron: X(g) + e⁻ → X⁻(g). It can be negative (energy released) or positive (energy needed).
• Halogens have strongly negative ΔegH because one electron completes a noble gas configuration. Noble gases have large positive values because the electron must enter a new, higher shell. The most negative values are at the upper right, just before the noble gases.
• Trend, less regular than ionization enthalpy: ΔegH usually becomes more negative across a period (smaller atom, higher effective nuclear charge) and less negative down a group (larger atom).
• Exception: O and F are less negative than S and Cl. The electron entering the compact n = 2 shell is strongly repelled; in n = 3 there is more room. So Cl (−349) is more negative than F (−328), and of P, S, Cl and F, Cl is most negative and P least.
• NCERT values (kJ mol⁻¹): H −73; Li −60, Na −53, K −48, Rb −47, Cs −46; O −141, S −200, Se −195, Te −190, Po −174; F −328, Cl −349, Br −325, I −295, At −270; He +48, Ne +116, Ar +96, Kr +96, Xe +77, Rn +68.
• Many books use electron affinity Aₑ, the negative of ΔegH, so a release of energy gives a positive Aₑ. Defined at absolute zero, it relates as ΔegH = −Aₑ − (5/2)RT.
• Electronegativity is a qualitative measure of how strongly an atom in a compound pulls shared electrons to itself. It cannot be measured directly and changes with the bonding partner. Scales: Pauling (most used), Mulliken-Jaffe, Allred-Rochow; Pauling (1922) fixed fluorine at 4.0.
• Pauling values: Li 1.0, Be 1.5, B 2.0, C 2.5, N 3.0, O 3.5, F 4.0; Na 0.9, Mg 1.2, Al 1.5, Si 1.8, P 2.1, S 2.5, Cl 3.0; down group 1: Li 1.0, Na 0.9, K 0.8, Rb 0.8, Cs 0.7; down group 17: F 4.0, Cl 3.0, Br 2.8, I 2.5, At 2.2.
• Electronegativity rises across a period and falls down a group, following the inverse of atomic radius, like ionization enthalpy. It goes with non-metallic character and against metallic character.

11. Valence and the anomalous second period

NCERT §3.7.2

• Valence of representative elements usually equals the number of outer electrons or eight minus it. Groups 1, 2, 13, 14, 15, 16, 17, 18 have 1 to 8 valence electrons and valences 1, 2, 3, 4, 3 or 5, 2 or 6, 1 or 7, 0 or 8.
• Oxidation state is the charge an atom takes on when shared electrons are assigned by electronegativity. In OF₂ (F > O) each F is −1, so O is +2; in Na₂O, O takes an electron from each Na and is −2, and each Na is +1.
• Predicting formulas: Si (group 14, valence 4) with Br (valence 1) gives SiBr₄; Al (group 13, valence 3) with S (group 16, valence 2) gives Al₂S₃.
• Hydrides and oxides follow the valence: LiH, CaH₂, B₂H₆, CH₄, NH₃, H₂O, HF; Na₂O, MgO, Al₂O₃, SiO₂, P₄O₁₀, SO₃, Cl₂O₇. Transition elements and actinoids commonly show variable valence.
• The first member of groups 1, 2 and 13-17 (Li, Be, B to F) differs from the rest of its group. Li and Be form compounds with clear covalent character, while their heavier relatives are mainly ionic.
• Diagonal relationship: Li resembles Mg, and Be resembles Al, the second element of the next group. Metallic radii (pm): Li 152, Mg 160; Be 111, Al 143.
• Causes: small size, large charge/radius ratio and high electronegativity; also, a second-period atom has only four valence orbitals (2s, 2p) against nine (3s, 3p, 3d) for the next member. So its maximum covalency is 4 (B forms BF₄⁻) while Al can form AlF₆³⁻.
• Second-period p-block atoms form pπ–pπ multiple bonds readily, with themselves (C=C, C≡C, N=N, N≡N) and with other second-period atoms (C=O, C=N, C≡N, N=O).
• Oxidation state and covalency differ: in [AlCl(H₂O)₅]²⁺, Al has oxidation state +3 but covalency 6.

12. Periodic trends and chemical reactivity

NCERT §3.7.3

• Every chemical and physical property shows up the electronic configuration. Across a period radii shrink, ionization enthalpy mostly rises and electron gain enthalpy becomes more negative.
• So reactivity is highest at the two ends of a period and lowest in the middle: on the left by losing an electron to form a cation (alkali metals), on the right by gaining one to form an anion (halogens). This links to reducing and oxidising behaviour.
• Metallic character is greatest at the far left and gives way to non-metallic character towards the right.
• Reaction with oxygen shows it: the oxide at the left end is the most basic (Na₂O) and at the right end the most acidic (Cl₂O₇). Oxides in the middle are amphoteric (Al₂O₃, As₂O₃) or neutral (CO, NO, N₂O).
• Amphoteric oxides act as acids towards bases and as bases towards acids; neutral oxides do neither. With water: Na₂O + H₂O → 2NaOH (strong base) and Cl₂O₇ + H₂O → 2HClO₄ (strong acid); litmus confirms each.
• In the 3d transition series, radii change much less across the period than for main group elements, and even less in the 4f series. Their ionization enthalpies lie between those of the s- and p-blocks, so they are less electropositive than groups 1 and 2.
• Down a main group, radii grow, ionization enthalpy falls gradually and electron gain enthalpy becomes less negative (except that O and F are less negative than S and Cl). Metallic character rises and non-metallic character falls. Transition elements show the reverse trend down a group, explained by size and ionization enthalpy.

Must-know facts

1. Mendeleev: properties are periodic in atomic weight. Modern law (after Moseley, 1913): periodic in atomic number.
2. Dobereiner's triads: middle atomic weight ≈ average of the ends (Li 7, Na 23, K 39). Newlands' octaves worked only up to calcium.
3. Eka-aluminium = gallium; eka-silicon = germanium.
4. Seven periods of 2, 8, 8, 18, 18, 32 and 32 (theoretical) elements; 18 IUPAC groups.
5. Period number = n of the valence shell; elements per period = 2 × orbitals being filled.
6. IUPAC roots: nil 0, un 1, bi 2, tri 3, quad 4, pent 5, hex 6, sept 7, oct 8, enn 9, then -ium. Z = 120 is unbinilium, Ubn.
7. Lanthanoids Ce (58) to Lu (71); actinoids Th (90) to Lr (103).
8. d-block general configuration (n−1)d¹⁻¹⁰ns⁰⁻²; Pd is 4d¹⁰5s⁰. Zn, Cd, Hg are not typical transition metals.
9. Metalloids: Si, Ge, As, Sb, Te. Metals exceed 78% of the elements.
10. Covalent radius of Cl = 198/2 = 99 pm; metallic radius of Cu = 256/2 = 128 pm.
11. Cation < parent atom < anion: Na⁺ 95 pm < Na 186 pm; F 64 pm < F⁻ 136 pm.
12. Isoelectronic size order: O²⁻ > F⁻ > Na⁺ > Mg²⁺.
13. Period 2 ionization enthalpy: Li < B < Be < C < O < N < F < Ne.
14. Most negative electron gain enthalpy of all: Cl (−349 kJ mol⁻¹), not F (−328).
15. Noble gases: positive electron gain enthalpy (Ne +116 kJ mol⁻¹).
16. Pauling scale: F = 4.0, the highest; Cs = 0.7 among group 1.
17. Diagonal pairs: Li–Mg and Be–Al.
18. Second-period maximum covalency 4 (BF₄⁻); Al reaches 6 (AlF₆³⁻).
19. Oxides: Na₂O basic, Cl₂O₇ acidic, Al₂O₃ and As₂O₃ amphoteric, CO, NO and N₂O neutral.

Common traps

Formulas

Key terms

Triad

Dobereiner's set of three similar elements whose middle atomic weight is about the mean of the other two.

Law of Octaves

Newlands' observation that, by atomic weight, every eighth element repeats the first element's properties.

Eka-aluminium

Mendeleev's name for the missing element below aluminium, later found as gallium.

Modern Periodic Law

Properties of elements repeat periodically with atomic number.

Period

A horizontal row; its number equals the principal quantum number of the valence shell.

Group

A vertical column of elements with similar outer electron configuration, numbered 1 to 18.

Representative elements

The s- and p-block elements together, also called main group elements.

Transition elements

d-block elements of groups 3 to 12, filling (n−1)d orbitals.

Inner transition elements

f-block lanthanoids and actinoids, filling (n−2)f orbitals.

Metalloid

An element along the zig-zag border, such as Si or As, with both metallic and non-metallic properties.

Covalent radius

Half the single-bond distance between two identical bonded atoms.

Isoelectronic species

Atoms or ions that carry the same number of electrons, such as Na⁺ and F⁻.

Effective nuclear charge

The net pull felt by an outer electron after inner electrons screen part of the nuclear charge.

Ionization enthalpy

Energy to remove the most loosely held electron from an isolated gaseous atom in its ground state.

Electron gain enthalpy

Enthalpy change when a gaseous atom accepts one electron to become a negative ion.

Electronegativity

The pull an atom in a compound exerts on a shared electron pair; a relative, not measured, quantity.

Diagonal relationship

Resemblance between a second-period element and the element one down and one to the right, as Li and Mg.

Amphoteric oxide

An oxide that reacts with both acids and bases, such as Al₂O₃.