Classification of Elements and Periodicity in Properties

In 1800 only 31 elements were known; by 1865 there were 63, and today there are 118. Studying each one on its own is unworkable, so chemists looked for a way to group elements that behave alike. Once similar elements sit together, the chemistry of one member of a family tells you about the others.

The periodic table is chemistry's most successful classification, and it is a tool for prediction rather than a chart to memorise. From an element's position you can predict its valence, the formulas of its oxide and hydride, whether it is a metal or a non-metal, and how its size and reactivity compare with its neighbours. Every earlier attempt was a search for the property whose repetition is truly periodic; it turned out to be electronic configuration.

Döbereiner's triads: by 1829 Johann Döbereiner had noticed groups of three similar elements in which the atomic mass of the middle one is roughly the mean of the other two: Li 7, Na 23, K 39; also Ca–Sr–Ba and Cl–Br–I. Limitation: only a few such triads could be found.

Newlands' law of octaves (1865): arranged by increasing atomic mass, every eighth element resembled the first. Limitation: it held only for the lighter elements, up to calcium.

Mendeleev's periodic table (1869): Mendeleev (and independently Lothar Meyer) arranged the 63 known elements by increasing atomic mass, grouping those with similar properties, judged by the formulas of their hydrides and oxides. His periodic law: the properties of the elements are a periodic function of their atomic weights. He left gaps for undiscovered elements and predicted their properties, naming them eka-aluminium (gallium, found 1875) and eka-silicon (germanium, 1886). Where properties demanded it he ignored mass order, placing iodine after the heavier tellurium to keep it with the halogens.

Limitations: such reversed pairs (Te–I, Co–Ni, later Ar–K) had no explanation in a mass-based law, isotopes had no clear place, and hydrogen's position was uncertain.

In 1913 Henry Moseley found that the square root of the frequency of the characteristic X-rays of an element rises linearly with its atomic number, not its atomic mass. Atomic number, the number of protons, is therefore the fundamental property, giving the modern periodic law: the physical and chemical properties of the elements are periodic functions of their atomic numbers. Properties repeat because electronic configuration repeats, and Ar (Z = 18) simply precedes K (Z = 19).

The long form of the periodic table has 7 horizontal periods and 18 vertical groups, numbered 1 to 18 by IUPAC. Each period begins when electrons start entering a new principal shell and ends with a noble gas. Elements in a group have the same number of valence electrons and similar chemistry; elements in a period have the same number of shells but a nuclear charge that rises by one at each step, so their properties change steadily from left to right.

Period lengths are 2, 8, 8, 18, 18, 32 and 32: twice the number of orbitals filled in the period. Period 6 fills 6s, 4f, 5d, 6p (16 orbitals, 32 elements).

New heavy elements are made in a few laboratories, and disputes over priority made naming slow. IUPAC therefore gives each new element a temporary systematic name and three-letter symbol built from the digits of its atomic number, used until a permanent name is approved. The digit roots are 0 nil, 1 un, 2 bi, 3 tri, 4 quad, 5 pent, 6 hex, 7 sept, 8 oct, 9 enn. Write the roots for the digits of Z in order, add -ium, and take the first letter of each root as the symbol.

Examples: Z = 104 is un-nil-quad-ium, Unnilquadium (Unq), now rutherfordium (Rf). Z = 106 is Unnilhexium (Unh), now seaborgium (Sg). Z = 118 is Ununoctium (Uuo), now oganesson (Og). NCERT's worked example is Z = 120: un-bi-nil-ium, Unbinilium (Ubn). A double letter is dropped where roots meet, so Z = 109 is Unnilennium (Une) and Z = 102 is Unnilbium (Unb).

The shape of the periodic table is a picture of the order in which orbitals fill: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. A new period starts every time electrons begin to enter a new principal shell. Periods 1, 2 and 3 fill 1s, then 2s and 2p, then 3s and 3p; 3d waits because 4s is lower in energy. Period 4 starts with 4s (K, Ca), then the ten 3d elements Sc to Zn (the first transition series), then 4p (Ga to Kr). Period 6 fills 6s, 4f (the lanthanoids), 5d and 6p; period 7 fills 7s, 5f (the actinoids), 6d and 7p and ends at Og (Z = 118).

Two rules follow. The period number equals the principal quantum number n of the outermost shell: sodium, 1s²2s²2p⁶3s¹, has highest n = 3, so period 3. The group is fixed by the number and type of valence electrons, the same for every member: all alkali metals are ns¹, all halogens ns²np⁵, which is why a group behaves alike.

Elements are sorted into four blocks by the subshell that receives the last electron.

s-block (groups 1, 2): ns¹ alkali metals, ns² alkaline earth metals: reactive metals with low ionization enthalpy. Group = number of valence electrons.

p-block (groups 13 to 18): ns²np¹ to ns²np⁶. With the s-block these are the representative (main-group) elements. It ends with the noble gases, whose complete ns²np⁶ shell (1s² for He) makes them very unreactive. Group = 10 + valence electrons.

d-block, the transition elements (groups 3 to 12): (n−1)d^1–10ns^0–2. Metals with variable oxidation states and coloured ions; the series are 3d (Sc–Zn), 4d (Y–Cd), 5d (La, Hf–Hg). Zn, Cd and Hg (full d¹⁰) lack most transition properties. Group = (n−1)d electrons + ns electrons.

f-block, the inner transition elements: (n−2)f^1–14(n−1)d^0–1ns². The lanthanoids (Ce–Lu, 4f) and actinoids (Th–Lr, 5f) all belong to group 3 and are chemically very uniform because the differentiating electron enters a deep f subshell.

Metals (over 78% of all elements) lie on the left and centre, non-metals at the top right, and a zig-zag line separates them; elements on the line such as Si, Ge, As, Sb and Te are metalloids. Hydrogen (1s¹) fits group 1 or 17 and is shown separately.

An atom has no sharp edge, so radius is defined operationally. Covalent radius is half the internuclear distance between two like atoms joined by a single bond. Metallic radius is half the distance between adjacent atoms in a metal crystal. Van der Waals radius, half the distance between nuclei of two non-bonded neighbouring atoms, is used for noble gases and is larger than the covalent radius.

Across a period atomic radius decreases (Li 152 pm to F 64 pm): the shells stay the same while the nuclear charge rises by one proton at each step, so the outer electrons are pulled in more strongly. (Neon looks larger only because its van der Waals radius is quoted.) Down a group it increases (Li 152 pm to Cs 262 pm): a new shell is added, and the extra shielding by inner electrons outweighs the rise in nuclear charge.

Ionic radius. A cation is smaller than its parent atom: removing electrons often empties the outer shell, and the same nuclear charge acts on fewer electrons (Na 186 pm, Na⁺ 95 pm). An anion is larger: the added electron raises electron–electron repulsion while the nuclear charge is unchanged. For isoelectronic species, which have the same number of electrons, radius decreases as nuclear charge increases: N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺, all with 10 electrons.

Ionization enthalpy (Δ_iH) is the energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state: X(g) → X⁺(g) + e⁻. It is always positive, and each successive value is larger because the electron leaves a more positive ion.

The value is set by the effective nuclear charge: the nuclear attraction less the shielding by inner electrons. Within a shell s electrons penetrate most, so shielding effectiveness is s > p > d > f.

Across a period Δ_iH increases (Li 520, Ne 2080 kJ mol⁻¹): nuclear charge rises while shielding hardly changes. Down a group it decreases (Li 520, Cs 374): the outer electron is farther away and shielded by more shells.

Exceptions in period 2 (actual order Li < B < Be < C < O < N < F < Ne): Be (2s², 899) is higher than B (2s²2p¹, 801) because B's electron leaves a 2p orbital, higher in energy, less penetrating and shielded by the 2s pair. N (2p³, 1402) is higher than O (2p⁴, 1314) because N's 2p electrons occupy separate orbitals with parallel spins, a stable half-filled subshell, while O's fourth electron is paired and its repulsion makes it easier to remove.

Electron gain enthalpy (Δ_egH) is the enthalpy change when an electron is added to a neutral gaseous atom to form a negative ion: X(g) + e⁻ → X⁻(g). It is negative when energy is released, positive when energy must be supplied. A second electron gain enthalpy is always positive because the electron is forced onto a negative ion: O⁻(g) + e⁻ → O²⁻(g) is about +780 kJ mol⁻¹.

Across a period Δ_egH becomes more negative as the atom shrinks and its effective nuclear charge grows. The halogens have the most negative values of all, since one electron completes a stable ns²np⁶ shell. Noble gases have positive values: the added electron would have to enter the next higher shell. Down a group the value becomes less negative as the atom grows.

The F–Cl exception: fluorine (−328 kJ mol⁻¹) is less negative than chlorine (−349), and oxygen (−141) less negative than sulphur (−200). The 2p subshell of a period-2 atom is so small that the added electron is strongly repelled by the electrons already there, which partly cancels the nuclear attraction; the larger 3p subshell of Cl or S has more room.

Electronegativity is the ability of an atom in a chemical compound to attract the shared electrons of a bond towards itself. Unlike ionization enthalpy or electron gain enthalpy, which are measurable energies of an isolated atom, it is a relative number assigned to a bonded atom, and it varies a little with oxidation state and bonding partner. The usual scale is Pauling's, which fixes fluorine at 4.0, the highest value; O is 3.5, N and Cl 3.0, C 2.5, and the alkali metals lie below 1.0.

Across a period electronegativity increases (Li 1.0 to F 4.0) as the atom becomes smaller and its effective nuclear charge larger. Down a group it decreases (F 4.0, Cl 3.0, Br 2.8, I 2.5) as the atom grows.

It also fixes oxidation states: in OF₂ the more electronegative fluorine takes the electrons and oxygen is +2, while in Na₂O oxygen is −2.

The valence of a representative element is usually the number of valence electrons, or eight minus that number. Across period 3 the valence electrons run 1 to 8 for Na to Ar, so the valence towards oxygen runs 1 to 7 and the oxides are Na₂O, MgO, Al₂O₃, SiO₂, P₂O₅, SO₃, Cl₂O₇. The valence towards hydrogen rises to 4 and then falls: NaH, MgH₂, AlH₃, SiH₄, PH₃, H₂S, HCl. The same valence repeats down a group, so all group 14 hydrides are XH₄.

Chemists now prefer oxidation state, which is fixed by electronegativity. In OF₂ fluorine is the more electronegative atom, so each F is −1 and O is +2; in Na₂O oxygen is −2 and Na is +1. Transition elements show variable oxidation states because their (n−1)d and ns electrons are close in energy, and heavier p-block elements show two states differing by two, such as Pb(II) and Pb(IV).

The first element of each group (Li, Be, B, C, N, O, F) differs in several ways from the rest of its group. NCERT gives four causes: small size, large charge-to-radius ratio, high electronegativity, and the absence of d orbitals in the valence shell.

Maximum covalency of four. A period-2 atom has only four valence orbitals (one 2s, three 2p), so it forms at most four covalent bonds: boron forms only [BF₄]⁻, while aluminium, with empty 3d orbitals, forms [AlF₆]³⁻. Multiple bonding. Small period-2 atoms readily form pπ–pπ multiple bonds (C=C, C≡C, N≡N, C=O), which larger period-3 atoms rarely do; N₂ is a triple-bonded gas while phosphorus is single-bonded P₄.

Diagonal relationship. Because it is unusually small, a period-2 element often resembles the period-3 element diagonally below and to its right more than its own group, since their charge-to-radius ratios are comparable. The pairs are Li–Mg, Be–Al and B–Si. Like Mg, lithium burns to a normal oxide rather than a peroxide and its carbonate decomposes on heating; like Al, beryllium has an amphoteric oxide and covalent chlorides.

Chemical reactivity is highest at the two ends of a period and lowest in the middle. On the left, low ionization enthalpy lets the metal lose electrons: it is a good reducing agent. On the right, a very negative electron gain enthalpy lets the non-metal gain electrons: it is a good oxidising agent. Metallic character decreases across a period and increases down a group; non-metallic character does the opposite. In group 1 reactivity increases downward (Li < Na < K < Rb < Cs) as ionization enthalpy falls; in group 17 it decreases downward (F > Cl > Br > I) as electronegativity falls.

The nature of oxides shows this directly. Across period 3, Na₂O and MgO are basic (Na₂O + H₂O → 2NaOH), Al₂O₃ is amphoteric, and SiO₂, P₄O₁₀, SO₃ and Cl₂O₇ are increasingly acidic (Cl₂O₇ + H₂O → 2HClO₄). Metal oxides are basic, non-metal oxides acidic, oxides of elements near the dividing line amphoteric. Down a group oxides become more basic: BeO is amphoteric, BaO strongly basic.