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Boron and aluminium are always +3, so why is thallium most stable as +1? Carbon is +4, so why is lead mostly +2? The usual answer, "the inert pair effect", can sound like a label. This post explains what the effect is, what causes it, and how to use it, with three worked examples.
Which exams test it, and where is it in the books?
The JEE (Main) 2026 syllabus and the NEET (UG) 2026 syllabus include a p-block unit on Groups 13 to 18 with "general trends in physical and chemical properties ... down the groups". The JEE (Advanced) 2026 brochure asks for "oxidation state and trends in chemical reactivity of elements of groups 13-17". The current NCERT Class 11 chapter on classification of elements names the inert pair effect and says it will be treated with each group, but the current NCERT books and the CBSE 2026-27 chemistry curriculum no longer have p-block chapters. The NIOS senior secondary chemistry lesson on the p-block is a free official second source. Check your exam's syllabus for the current year.
What is the inert pair effect?
NIOS describes the trend: in groups 13, 14 and 15, higher oxidation states become less stable going down the group. Boron and aluminium are trivalent, but gallium, indium and thallium also show +1, and "+1 state of thallium is very stable". Divalent germanium, tin and lead compounds exist, and +3 is the stable state of antimony and bismuth. The outer configurations are ns²np¹, ns²np² and ns²np³, so the preferred states (+1, +2, +3) suggest that the two s electrons do not take part in bonding. That reluctance is what the name describes.
A 2025 analysis by Norman and Winter traces the idea to Nevil Sidgwick in 1927 and uses a large database of known compounds to show "a preponderance of lower valent compounds for the elements thallium, lead and bismuth".
What actually causes it?
The name implies the 6s electrons are hard to remove. Norman and Winter point out that this is something of a misnomer: as Russell Drago argued in 1958, the 6s ionisation energies of thallium and lead are comparable to, or somewhat lower than, those of the ns electrons of some lighter elements in their groups.
The explanation they summarise works like an energy budget. To form MX₃ instead of MX, a group 13 atom must first promote an electron from s to p, which costs energy. For light elements, the two extra bonds release more than enough to pay for it. Down the group, covalent bonds get weaker faster than the promotion energy changes, so for the 6p elements the two extra, weaker bonds no longer pay. NIOS gives the same two factors, promotion energy and poorer orbital overlap in large atoms, and adds that once these energies are considered, the term "inert pair effect" loses its significance.
Worked example 1: count the electrons
Thallium, lead and bismuth have atomic numbers 81, 82 and 83 (see NCERT's classification chapter).
Tl: [Xe]4f¹⁴5d¹⁰6s²6p¹. Tl⁺ has 81 − 1 = 80 electrons: [Xe]4f¹⁴5d¹⁰6s².
Pb: [Xe]4f¹⁴5d¹⁰6s²6p². Pb²⁺ has 82 − 2 = 80 electrons: [Xe]4f¹⁴5d¹⁰6s².
Bi: [Xe]4f¹⁴5d¹⁰6s²6p³. Bi³⁺ has 83 − 3 = 80 electrons: [Xe]4f¹⁴5d¹⁰6s².
The preferred ions are isoelectronic. Each has lost only its 6p electrons and kept the 6s² pair. The higher states Tl³⁺, Pb⁴⁺ and Bi⁵⁺ would each have 78 electrons, with the 6s pair gone too.
Worked example 2: the lead–acid battery
The NCERT Class 12 chapter Electrochemistry gives the reactions in a lead storage battery:
Anode: Pb(s) + SO₄²⁻(aq) → PbSO₄(s) + 2e⁻
Cathode: PbO₂(s) + SO₄²⁻(aq) + 4H⁺(aq) + 2e⁻ → PbSO₄(s) + 2H₂O(l)
Assign oxidation numbers to lead. At the anode it goes from 0 in Pb to +2 in PbSO₄: oxidation by 2. At the cathode it goes from +4 in PbO₂ to +2 in PbSO₄: reduction by 2. Both electrodes end at Pb(II). PbO₂ works as the oxidising agent because Pb(IV) readily accepts two electrons to reach the more stable +2 state, which fits Norman and Winter's note that inorganic Pb(IV) compounds are generally less stable and often strongly oxidising.
Worked example 3: which halides exist?
Norman and Winter note that all four lead(II) halides PbX₂ are known, but of the lead(IV) halides only PbF₄ is stable at room temperature. For carbon and silicon the tetrahalides are far more stable than the dihalides, and for germanium and tin both kinds can be isolated. The +4 state gets less stable down group 14, and lead is where +2 takes over. Fluorine is the exception because Pb–F bonds are strong enough to pay for using the 6s pair. Bismuth follows the same logic: of its pentahalides, only BiF₅ is thermally stable.
Which mistakes cost marks?
Saying the s electrons "cannot" be removed. They can; Pb(IV) and Tl(III) compounds exist. They are less stable, not impossible.
Applying the effect to light elements. Boron and aluminium are firmly +3.
Forgetting it makes higher states oxidising. Tl(III), Pb(IV) and Bi(V) compounds tend to be reduced.
For the rest of the p-block, see our post on groups 15 to 17, and for the trends behind it, periodic trends and their exceptions.
If oxidation numbers themselves are the weak spot, the Fermi Concept Library covers every concept in chemistry, each with its own lesson, so you can fix that first.
A practice routine
Day 1: Write the outer configurations for groups 13 to 15 and the two common oxidation states of each heavy element.
Day 2: Redo Worked example 1 for indium and tin, then explain why their lower states are less dominant.
Day 3: Assign oxidation numbers in five lead and thallium compounds and say which would act as oxidants.
Day 4: Explain the energy-budget argument in five sentences without notes.
Today, explain to yourself in two sentences why PbO₂ is a strong oxidising agent while PbO is far less so.
