Modern Periodic Table: History, Groups, Blocks and Periodic Trends

History of the Periodic Table: Dobereiner to Moseley

Dobereiner's Triads, Chancourtois and Newlands' Law of Octaves

The need for an order among the elements grew with their number. In 1800 only 31 elements were known, and by 1865 the figure had more than doubled to 63. Studying each element and its many compounds separately had become impractical, so chemists looked for a system that would organise known facts and also predict new ones. The search moved from atomic weight to atomic number over about a century.

The idea that elements show trends in their properties began with the German chemist Johann Dobereiner in the early 1800s. By 1829 he had noticed that several groups of three similar elements, which he called triads, shared a pattern. The middle element had an atomic weight about half way between those of the other two, and its properties also lay between theirs. Lithium, sodium and potassium (7, 23 and 39) form one triad, since the average of 7 and 39 is 23. Calcium, strontium and barium (40, 88 and 137) and chlorine, bromine and iodine (35.5, 80 and 127) form two more.

The rule worked for only a few elements, so it was dismissed as coincidence. In 1862 the French geologist A. E. B. de Chancourtois arranged the then known elements by increasing atomic weight on a cylindrical table to display the periodic recurrence of properties, but this too drew little attention.

In 1865 the English chemist John Alexander Newlands proposed the law of octaves. Arranging the elements by increasing atomic weight, he noted that every eighth element had properties similar to the first, like the eighth note of a musical scale. The law held only up to calcium, and it was not widely accepted at the time. The Royal Society, London, awarded Newlands the Davy Medal in 1887 for the work. NCERT dates the law of octaves to 1865; other accounts place Newlands' first classification in 1864.

Mendeleev's Periodic Law, Gaps and Predictions

In 1869 the Russian chemist Dmitri Mendeleev (1834 to 1907) and the German chemist Lothar Meyer (1830 to 1895) proposed, working independently, that when elements are arranged by increasing atomic weight, similarities in physical and chemical properties appear at regular intervals. Meyer plotted atomic volume, melting point and boiling point against atomic weight and obtained a periodically repeating pattern. By 1868 he had a table closely resembling the modern one, but it was published only after Mendeleev's work, so Mendeleev is credited with the first published periodic law. It states that the properties of the elements are a periodic function of their atomic weights.

Mendeleev set elements in horizontal rows and vertical columns so that those with similar properties fell in one column, called a group. He relied on the formulas and properties of the compounds each element formed, a broader basis than Meyer's. Where strict order of atomic weight broke the pattern, he put similar elements together instead. Iodine, with a lower atomic weight than tellurium of Group VI, was placed in Group VII with fluorine, chlorine and bromine because its properties match theirs.

His boldest step was to leave gaps for elements not yet discovered. He named the missing element beneath aluminium eka-aluminium and the one beneath silicon eka-silicon; the prefix is the Sanskrit numeral for one, used for an element one place below a known element of the same group. He described their properties in advance, and gallium and germanium later matched.

Gallium was discovered in 1875 and germanium was isolated in 1886. A third prediction, eka-boron with atomic mass 44, matched scandium (44.955907), and eka-manganese with atomic mass 100 matched technetium. The accuracy of these quantitative predictions made Mendeleev and his table famous. The noble gases, which he had not foreseen, fitted when argon, isolated in 1894 by Ramsay and Rayleigh, was followed by its relatives; Mendeleev placed them in a new Group 0 in 1902.

Moseley's Atomic Number and the Modern Periodic Law

When Mendeleev built his table, chemists knew nothing of the internal structure of the atom. In 1913 the English physicist Henry Moseley studied the characteristic X‑ray spectra of the elements and found regularities. A plot of the square root of the X‑ray frequency against atomic number (Z) gave a straight line, whereas the same plot against atomic mass did not. He thereby showed that atomic number is a more fundamental property of an element than its atomic mass.

Mendeleev's law was modified accordingly. The atomic number equals the nuclear charge, that is the number of protons, and also the number of electrons in a neutral atom. This made it easy to see why electron configuration controls periodicity. The periodic law is the consequence of the periodic variation in the electronic configurations that fix an element's properties. Moseley's measurements put argon before potassium, cobalt before nickel and tellurium before iodine, settling the reversed pairs of the older table. He was shot and killed at Gallipoli on 10 August 1915, at the age of 27.

The law revealed analogies among the 94 naturally occurring elements, and it has carried on to the artificially produced, short-lived elements. In the middle of the 20th century the American chemist Glenn T. Seaborg's work on the transuranium elements, beginning with plutonium, led to the actinoids being placed below the lanthanoids. He received the 1951 Nobel Prize in Chemistry, and element 106 is named seaborgium in his honour.

Structure of the Long Form Periodic Table

Groups, Periods and Electronic Configuration

The most convenient and widely used layout is the long form of the table. Its horizontal rows, which Mendeleev called series, are periods, and its vertical columns are groups or families. Elements with similar outer electronic configurations stand in one group. On the recommendation of the International Union of Pure and Applied Chemistry (IUPAC), groups are numbered 1 to 18, replacing the older notation IA to VIIA, VIII, IB to VIIB and 0.

There are seven periods, and the period number equals the highest principal quantum number (n) of the elements in it. The number of elements in a period is twice the number of orbitals available in the energy level being filled, which gives 2, 8, 8, 18, 18 and 32 for the first six periods. The seventh period would also hold a maximum of 32, and with the discovery of element 118 it is complete.

The fourth period starts at potassium and fills 4s before the 3d orbitals, giving the 3d transition series from scandium (Z = 21) to zinc (Z = 30), and ends at krypton. The fifth period begins with rubidium and contains the 4d series starting at yttrium (Z = 39), ending at xenon. In the sixth period, filling of 4f begins with cerium (Z = 58) and ends at lutetium (Z = 71), the lanthanoid series. In the seventh, filling of 5f follows actinium (Z = 89), the actinoid series. These 14 lanthanoids and 14 actinoids are set in two panels at the bottom, which keeps the table compact and keeps elements of similar properties in a single column.

Elements in one group share the same number and arrangement of outer electrons, so they behave alike. All Group 1 elements have an ns¹ valence configuration. Of the 118 elements, only those up to atomic number 94 exist in nature; the rest are synthetic. Elements after uranium are called transuranium elements.

s Block and p Block Elements: Main Group Elements

The elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) have ns¹ and ns² outermost configurations and make up the s block. All are reactive metals with low ionisation enthalpies. They lose the outermost electrons readily to form ions with a single positive charge (alkali metals) or a double positive charge (alkaline earth metals). Metallic character and reactivity increase down the group, and because of their high reactivity they are never found pure in nature. Their compounds are predominantly ionic, except those of lithium and beryllium.

The p block comprises Groups 13 to 18. With the s block, these are the representative or main group elements. The outermost configuration varies from ns² np¹ to ns² np⁶ in each period, and each period ends in a noble gas with a closed ns² np⁶ shell that is very hard to alter, so the noble gases show very low reactivity. Just before them stand two chemically important groups of non-metals, the halogens (Group 17) and the chalcogens (Group 16). Both have highly negative electron gain enthalpies and readily add one or two electrons respectively to reach the noble gas arrangement.

Two elements break the block rule. Helium belongs to the s block by configuration (1s²) but sits in Group 18, because its filled valence shell gives it the properties of the noble gases. Hydrogen has one s electron, so it can sit in Group 1, yet it can also gain an electron and behave like a Group 17 element. Because it is a special case, the table places it separately at the top.

d Block and f Block Elements: Transition and Inner Transition Metals

Groups 3 to 12 in the centre of the table form the d block, the elements in which inner d orbitals fill with electrons. Their general outer configuration is (n‑1)d¹⁻¹⁰ ns⁰⁻², except palladium (4d¹⁰ 5s⁰). They are all metals, and they commonly form coloured ions, show variable valence (oxidation states), are paramagnetic and often serve as catalysts. Finely divided iron in the Haber process and nickel in catalytic hydrogenation are examples. Zinc, cadmium and mercury, with the configuration (n‑1)d¹⁰ ns², do not show most of these properties. The d block is called the transition series because it forms a bridge between the chemically active s block metals and the less active elements of Groups 13 and 14.

The two rows at the bottom, the lanthanoids cerium (58) to lutetium (71) and the actinoids thorium (90) to lawrencium (103), form the f block, or inner transition elements. Their outer configuration is (n‑2)f¹⁻¹⁴ (n‑1)d⁰⁻¹ ns², and the last electron added to each enters an f orbital. All are metals, and properties within each series are quite similar. The chemistry of the early actinoids is more complicated than that of the lanthanoids because of the large number of oxidation states they can show. Actinoids are radioactive, and many have been made only in nanogram quantities or less by nuclear reactions.

A short exercise shows how position follows from configuration. Element 117, now named tennessine, belongs to Group 17, the halogen family, with the configuration [Rn] 5f¹⁴ 6d¹⁰ 7s² 7p⁵, because its 7p subshell is one electron short of the filled shell of the noble gas 118.

Metals, Non-metals and Metalloids

A second way to read the table is by character. Metals make up more than 78 per cent of the known elements and lie on the left. They are usually solids at room temperature, with mercury as the exception, and gallium and caesium melt at only 303 K and 302 K. They typically have high melting and boiling points, conduct heat and electricity well, and are malleable and ductile.

Non-metals lie at the top right and number fewer than twenty. They are usually solids or gases at room temperature with low melting and boiling points, though boron and carbon are exceptions. They conduct poorly, and most solid non-metals are brittle, neither malleable nor ductile. Along a row the character changes from metallic on the left to non-metallic on the right, and the metallic character increases down a group.

The change is not abrupt. A thick zig-zag line separates the two regions, and the elements bordering it, silicon, germanium, arsenic, antimony and tellurium, show properties of both. These are the metalloids, also called semi-metals. NCERT names these five; other accounts add boron to make six. Metalloids are typically semiconductors, which is why silicon and germanium matter so much in electronics.

A worked comparison shows how to rank metallic character. Take Si, Be, Mg, Na and P and put them in order of increasing metallic character.

1. Metallic character increases down a group, so magnesium is more metallic than beryllium.
2. Metallic character decreases from left to right across a period, so sodium outranks magnesium, which outranks silicon, which outranks phosphorus.
3. Place beryllium, in period 2, below magnesium but above silicon, which gives the order phosphorus, silicon, beryllium, magnesium and sodium, from least to most metallic.

Naming of New Elements and the 2016 Names

IUPAC Systematic Names for Elements Above 100

Naming a new element was traditionally the privilege of its discoverers, with IUPAC ratifying the suggestion. This caused disputes. For element 104, American scientists proposed rutherfordium and Soviet scientists kurchatovium. The elements with very high atomic numbers are so unstable that only minute quantities, sometimes a few atoms, are obtained, and work on them needs costly equipment in only a few laboratories, so rival claims arise before reliable data exist.

To avoid such disputes, IUPAC recommended that until a discovery is proved and a name is officially recognised, an element carries a systematic name derived from its atomic number. Each digit is replaced by a numerical root, the roots are joined in the order of the digits, and "ium" is added at the end. The symbol has three letters.

Element 120, if made, would take un, bi and nil, giving unbinilium with the symbol Ubn. Element 113 took un, un and tri, giving ununtrium, Uut. Later the permanent name and symbol are decided by a vote of IUPAC representatives from each country; the name may reflect the country or state of discovery or honour a notable scientist. The systematic names always end in "ium", even for elements expected to be halogens or noble gases, which is why 117 and 118 were ununseptium and ununoctium before their permanent names. Recommendations published in 2016 restored the "-ine" ending for new Group 17 elements and the "-on" ending for new Group 18 elements.

The Four Elements Named in 2016

The four newest elements completed the seventh period. In December 2015 the IUPAC and IUPAP Joint Working Party recognised elements 113, 115, 117 and 118, and their permanent names were adopted in 2016; moscovium and oganesson were formally named on 28 November 2016. All four are synthetic and extremely radioactive.

Oganesson has the highest atomic number and atomic mass of all known elements, and its only known isotope, oganesson‑294, has a half-life of 0.7 milliseconds. Moscovium and oganesson were first made at the Joint Institute for Nuclear Research at Dubna by Russian and American scientists working together, and nihonium was recognised as a discovery of the Riken team in Japan.

Periodic Trends in Properties of Elements

Atomic Radius and Ionic Radius

An atom is about 1.2 angstrom (1.2 × 10⁻¹⁰ m) in radius and its electron cloud has no sharp boundary, so size can only be estimated from the distance between atoms in the combined state. The covalent radius of a non-metal is half the distance between two bonded atoms, so the Cl-Cl bond distance of 198 pm gives 99 pm for chlorine. The metallic radius is half the distance between metal cores in the crystal, so copper atoms 256 pm apart give 128 pm. Atomic radius here covers both. Noble gases are left out, since their non-bonded radii are very large and are compared with van der Waals radii instead.

Across a period the radius generally decreases. The outer electrons stay in the same valence shell while the effective nuclear charge rises with atomic number, pulling them in. Down a group it increases, because each step adds a shell, the principal quantum number grows, and the inner filled shells screen the outer electrons from the nucleus.

Ionic radius follows the same trends as atomic radius. A cation is smaller than its parent atom because it has fewer electrons for the same nuclear charge; sodium is 186 pm and Na⁺ is 95 pm. An anion is larger because the added electrons repel one another and the effective nuclear charge per electron falls; fluorine is 64 pm and the fluoride ion 136 pm.

Species with the same number of electrons are isoelectronic. O²⁻, F⁻, Na⁺ and Mg²⁺ each have 10 electrons, yet their radii differ because their nuclear charges differ. The cation with the greater positive charge is smaller, and the anion with the greater negative charge is larger. For Mg, Mg²⁺, Al and Al³⁺, magnesium is the largest and Al³⁺ the smallest.

Across the 3d transition series the change in atomic radius is much smaller than among representative elements, and it is smaller still across the 4f series.

Ionisation Enthalpy and Its Exceptions

Ionisation enthalpy (ionization energy) is the energy needed to remove an electron from an isolated gaseous atom in its ground state. The first ionisation enthalpy is the enthalpy change for X(g) becoming X⁺(g) plus an electron, in kJ per mol. It is always positive, since energy is always needed, and each later ionisation enthalpy is higher than the one before because an electron is harder to remove from a positive ion. Unqualified, the term means the first.

Plotted against atomic number from 1 to 60, the first ionisation enthalpy shows maxima at the noble gases, whose closed shells are very stable, and minima at the alkali metals, whose low values explain their high reactivity. Two factors decide it, namely the attraction of the electron to the nucleus and the repulsion among electrons. The effective nuclear charge a valence electron feels is less than the actual nuclear charge, because inner electrons shield it, and shielding is most effective when the inner shells are completely filled, as in the alkali metals.

• Across a period: Added electrons enter the same principal level, shielding hardly increases, and the growing nuclear charge wins, so ionisation enthalpy generally rises.
• Down a group: The outer electron is farther away and shielding increases, which outweighs the growing nuclear charge, so ionisation enthalpy generally falls.

Two exceptions in period 2 are asked often. Boron's first ionisation enthalpy is slightly less than beryllium's even though its nuclear charge is greater, because beryllium loses a 2s electron, which penetrates closer to the nucleus, whereas boron loses a 2p electron, which is better shielded. Oxygen's is less than nitrogen's, because nitrogen's three 2p electrons occupy separate orbitals (Hund's rule), whereas two of oxygen's four 2p electrons share one orbital and repel, so it is easier to remove one. In period 3 the values for sodium, magnesium and silicon are 496, 737 and 786 kJ per mol, and aluminium's is lower than magnesium's, because its 3p electron is shielded by the 3s electrons.

Electron Gain Enthalpy and Electronegativity

When an electron is added to a neutral gaseous atom to form a negative ion, the enthalpy change is the electron gain enthalpy. It can be negative, when energy is released, or positive. Halogens have very negative values because one electron completes a noble gas configuration, while noble gases have large positive values, since the electron would have to enter the next higher principal level. Many books define electron affinity as the negative of this change, so a large positive electron affinity matches a very negative electron gain enthalpy.

Across a period the value generally becomes more negative, and down a group less negative, as the atom grows and the added electron sits farther from the nucleus. There is an exception in the second period. The values for oxygen and fluorine are less negative than those of sulfur and chlorine, since the added electron enters the small n = 2 level and meets strong repulsion from the electrons already there, whereas in the n = 3 level it has more room. Chlorine, not fluorine, therefore has the most negative value of all.

Electronegativity is the ability of an atom in a compound to attract shared electrons to itself. It is qualitative and not directly measurable, and its value for an element varies with the element it is bonded to. Scales include those of Pauling, Mulliken-Jaffe and Allred-Rochow, and the Pauling scale is most widely used; it gives fluorine, the element with the greatest ability to attract electrons, the value 4.0. Electronegativity generally rises from left to right across a period and falls down a group, tracking the shrinking and growing of the atom.

Down Group 17 the values run F 4.0, Cl 3.0, Br 2.8, I 2.5 and At 2.2; down Group 1 they run Li 1.0, Na 0.9, K 0.8, Rb 0.8 and Cs 0.7.

Metallic Character, Valency and Chemical Reactivity

Electronegativity is directly related to non-metallic character, since non-metals tend to gain electrons, and inversely related to metallic character. Rising electronegativity across a period goes with rising non-metallic character and falling metallic character; falling electronegativity down a group goes with rising metallic character. Transition elements show the reverse of the main group trend down a group, which can be explained by their atomic size and ionisation enthalpy.

Valence is the combining capacity, now often expressed as oxidation state. For representative elements it equals the number of outer electrons or eight minus that number.

Using this, silicon (valence 4) and bromine (1) form SiBr₄, and aluminium (3) and sulfur (2) form Al₂S₃. Oxidation state depends on electronegativity. In OF₂, fluorine is the more electronegative and takes −1, so oxygen shows +2, while in Na₂O oxygen takes −2 and sodium +1.

Chemical reactivity is highest at the two extremes of a period and lowest in the centre. At the left, the low ionisation enthalpy lets Group 1 metals lose an electron readily; at the right, the very negative electron gain enthalpy lets Group 17 non-metals gain one. The noble gases, with positive electron gain enthalpies, are the exception to the right-hand rule. Among alkali metals reactivity increases down the group, while among halogens it decreases. Highly reactive elements are not found free in nature.

The oxide of the element at the far left is the most basic (Na₂O, which gives NaOH with water) and that at the far right the most acidic (Cl₂O₇, which gives HClO₄). Oxides in the centre are amphoteric (Al₂O₃, As₂O₃), behaving as acids with bases and as bases with acids, or neutral (CO, NO, N₂O), with no acidic or basic property.

Anomalous Second Period Elements and the Diagonal Relationship

The first element of Groups 1 and 2 (lithium and beryllium) and of Groups 13 to 17 (boron to fluorine) differs in many respects from the rest of its group. Lithium and beryllium form compounds of pronounced covalent character, whereas the other members form mainly ionic compounds. In fact lithium resembles magnesium, and beryllium resembles aluminium, a similarity called the diagonal relationship.

The anomalies come from the small size, large charge to radius ratio and high electronegativity of these elements. There is a second reason. The first member has only four valence orbitals (2s and 2p), whereas the second member has nine (3s, 3p and 3d). The maximum covalency of the first member of each group is therefore 4, so boron can form only BF₄⁻, while aluminium forms AlF₆³⁻ by expanding its valence shell. First members of the p block are also better at forming pi-bonded multiple bonds to themselves and to other second-period elements, as in C=C, C‑triple-C, N=N, C=O and N=O.

Oxidation state and covalency are different things. In the ion [AlCl(H₂O)₅]²⁺, aluminium has oxidation state +3 but covalency 6.

Element Families and Their Uses

Alkali Metals and Alkaline Earth Metals

The alkali metals are lithium, sodium, potassium, rubidium, caesium and francium, the metals of Group 1. Each has one outer electron, so each readily forms an ion with a single positive charge. They tarnish rapidly in air and react with water, the heavier ones more vigorously, so they are stored under oil, usually mineral oil or kerosene, and are found in nature only as salts. Caesium is the most reactive of the metals, and francium, extremely radioactive, occurs only in minute traces. Lithium has the lowest density of any metal, 0.534 g per cubic centimetre, and lithium, sodium and potassium are the only metals that float on water.

• Common uses: Sodium chloride is table salt, the sodium-vapour lamp gives light very efficiently, and lithium serves as the anode in lithium batteries.
• In the body: Sodium and potassium are essential electrolytes with major biological roles.

The alkaline earth metals are the six elements of Group 2, namely beryllium, magnesium, calcium, strontium, barium and radium. They are shiny and silvery-white, with two outer electrons, which they lose to form an ion with a double positive charge. Magnesium and calcium are essential to all known living organisms. Magnesium occurs in chlorophyll in plants, and calcium salts form the bones of vertebrates and the shells of some invertebrates, while calcium ions signal inside cells.

Halogens and Noble Gases

The halogens of Group 17 are fluorine, chlorine, bromine, iodine, astatine and tennessine. The word means salt former, because they react with metals to give salts such as sodium chloride. With seven outer electrons they gain one electron readily, and their reactivity falls down the group. Fluorine is the most reactive of all elements, the only element more electronegative than oxygen, and it attacks glass and forms compounds even with the usually inert noble gases. Chlorine, bromine and iodine are often used as disinfectants, and iodine is needed in trace amounts to make thyroid hormones such as thyroxine. Astatine is radioactive and has no stable isotope.

The noble gases of Group 18 are helium, neon, argon, krypton, xenon and radon, with synthetic oganesson as the seventh member. Their valence shells are complete, ns² np⁶ (1s² for helium), so they show very low reactivity. They were once called the inert gases, but that label is now deprecated because compounds are known. Neil Bartlett made the first noble gas compound, xenon hexafluoroplatinate, in 1962. Argon makes up about 0.94 per cent of the atmosphere by volume.

Chalcogens and Pnictogens: Groups 16 and 15

The chalcogens of Group 16 are oxygen, sulfur, selenium, tellurium, radioactive polonium and livermorium. The name means ore-forming, since these elements were involved in the formation of ores. Each has six valence electrons, two short of a full shell, which is why they have highly negative electron gain enthalpies and add two electrons. Oxygen, sulfur and selenium are non-metals, tellurium is a metalloid, and the nature of polonium is uncertain.

The pnictogens of Group 15 are nitrogen, phosphorus, arsenic, antimony, bismuth and moscovium. The name comes from the Greek for "to choke", a reference to nitrogen gas, and it was suggested by the Dutch chemist Anton Eduard van Arkel in the early 1950s. The lighter pnictogens, nitrogen, phosphorus and arsenic, tend to form 3- ions when reduced, and the group's valence of 3 and 5 follows from five outer electrons. Moscovium, element 115, is the heaviest pnictogen.

Lanthanoids, Lanthanide Contraction and Rare Earth Elements

The 4f series, cerium (58) to lutetium (71), is the lanthanoid series. Lanthanum, element 57, is usually counted with them, giving the fifteen lanthanides from lanthanum to lutetium. Across the series, atomic and ionic radii fall steadily from lanthanum to lutetium. This is the lanthanide contraction, caused by imperfect shielding of the nuclear charge by the 4f electrons.

A direct result is that period 6 transition metals have radii very similar to period 5 ones. Zirconium is 155 pm and hafnium 159 pm, and the ions Zr⁴⁺ and Hf⁴⁺ are 84 and 83 pm. The two therefore behave almost identically, and hafnium is found only with zirconium, which is much more abundant; hafnium was discovered as a separate element in 1923.

The rare earth elements are a set of 17, made up of the 15 lanthanides together with scandium and yttrium. The name misleads. They are not actually scarce; cerium is the 25th most abundant element in the Earth's crust at 68 parts per million, more abundant than copper. They are called rare because they occur only in compounds, not as pure metals, and are difficult to isolate and purify.

• Magnets and catalysts: Most rare earths were used in 2015 for catalysts and magnets, and electric vehicles and wind turbines have since raised demand. Lanthanum chloride is used in fluid catalytic cracking to make petrol and diesel, and cerium(III) oxide in catalytic converters.
• Displays and light: The unpaired 4f electrons emit light after excitation, so rare earths are widely used in applications where light emission matters, such as phosphor lamps and displays.
• Electronics: Cerium and neodymium are important in electronics and are used in LCD and plasma screens, fibre optics and lasers.

Supply is concentrated. China supplied around 90 per cent of global demand for the 17 rare-earth powders in 2019 and has restricted supply and sales since around 2010. In 2025 it mined about 70 per cent of the world supply and processed about 90 per cent.

India has the third-largest reserves in the world, 6.9 million tonnes, including almost 35 per cent of the world's sand mineral deposits; the monazite is found in coastal, inland and riverine sands of Andhra Pradesh, Odisha, Tamil Nadu, Kerala, West Bengal, Jharkhand, Gujarat and Maharashtra. The government-owned Indian Rare Earths Limited is a major player, and India launched its National Critical Mineral Mission in 2025 to build self-reliance. Rare earth elements are on the list of 30 critical minerals the government released on 24 July 2023.

Actinoids and India's Three-Stage Nuclear Programme

The actinoids are the fourteen elements from thorium (90) to lawrencium (103) in which the 5f orbitals fill. All are radioactive. Naturally occurring uranium and thorium and synthetically produced plutonium are the most abundant actinides on Earth, and they have been used in nuclear reactors. Monazite is an important ore of thorium, lanthanum and cerium, and it contains small amounts of uranium; the deposits in India are particularly rich.

India's three-stage nuclear power programme, planned by the physicist Homi J. Bhabha in the 1950s, is built on these actinoids. India has only around 1 to 2 per cent of the world's uranium reserves but about 25 per cent of the world's known thorium reserves. Thorium itself is not fissile and cannot undergo fission to produce energy; it must be transmuted to uranium‑233 in a reactor fuelled by other fissile material. The first two stages, heavy water reactors on natural uranium and plutonium-fuelled fast breeder reactors, are meant to generate enough fissile material from limited uranium so that thorium can be used fully in the third.

1. Stage 1: Pressurised heavy water reactors burn natural uranium and produce plutonium.
2. Stage 2: Fast breeder reactors fuelled with plutonium generate more fissile material than they consume.
3. Stage 3: Thorium‑232 is converted to uranium‑233, the fuel of thorium-based reactors.

Fast Breeder Reactors, Criticality and the Kalpakkam PFBR

A thermal reactor uses a moderator to slow the neutrons released by fission. A fast reactor needs no neutron moderator, runs on fast neutrons and therefore requires fuel comparatively rich in fissile material. Because water would slow the neutrons, a fast reactor uses a liquid metal coolant, and the Kalpakkam reactor uses liquid sodium. A breeder produces more fissile fuel than it burns, because fast neutrons convert fertile uranium‑238 into fissile plutonium‑239.

Criticality is the start of a controlled, self-sustaining fission chain reaction. The 500 MWe Prototype Fast Breeder Reactor (PFBR) at Kalpakkam in Tamil Nadu attained first criticality on 6 April 2026 at 08:25 pm, after meeting all the stipulations of the Atomic Energy Regulatory Board (AERB). It was designed by the Indira Gandhi Centre for Atomic Research (IGCAR) and built and commissioned by Bharatiya Nabhikiya Vidyut Nigam Ltd (BHAVINI), a public sector undertaking under the Department of Atomic Energy.

For energy policy the milestone has three implications. It lets India draw far more energy from its limited uranium reserves; it uses a closed fuel cycle that recycles nuclear material; and it is the vital bridge between the present fleet of pressurised heavy water reactors and future thorium-based reactors, offering low-carbon, base-load power with higher thermal efficiency.

Sources

• NCERT: Chemistry Part I (Class XI), Unit 3, Classification of Elements and Periodicity in Properties
• Press Information Bureau: List of 30 critical minerals (2023)
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• Department of Atomic Energy: PFBR attains first criticality
• Bhabha Atomic Research Centre: About BARC and the nuclear programme
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• Wikipedia: Tennessine
• Wikipedia: Oganesson
• Wikipedia: Moscovium
• Wikipedia: Nihonium
• Wikipedia: Mendelevium
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• Wikipedia: Lithium
• Wikipedia: Alkaline earth metal
• Wikipedia: Halogen
• Wikipedia: Noble gas
• Wikipedia: Chalcogen
• Wikipedia: Pnictogen
• Wikipedia: Rare-earth element
• Wikipedia: Lanthanide contraction
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• UPSC: Civil Services Examination question papers