The s-Block Elements covers Group 1 (alkali metals: Li, Na, K, Rb, Cs, Fr) and Group 2 (alkaline earth metals: Be, Mg, Ca, Sr, Ba, Ra), tracing how their ns1 and ns2 configurations drive trends in atomic size, ionization enthalpy, reactivity with water and oxygen, flame colours, and the solubility of their hydroxides, carbonates and sulphates. It also explains the anomalous behaviour of lithium and beryllium, their diagonal relationships with magnesium and aluminium, and key industrial compounds like NaOH, washing soda and Plaster of Paris. For NEET this is a high-yield inorganic chapter where flame colours, oxide types and the anomalous-element facts are repeatedly tested as direct one-mark questions.
Table of Contents
- Key Concepts - configuration, periodic trends, anomalies, reactions, compounds, biology
- Weightage in Board & Entrance Exams
- Important Definitions
- Solved Examples
- Important Questions for Board Exams
- Quick Revision Points
Key Concepts
1. What Are s-Block Elements?
The s-block consists of the elements in which the last electron enters the outermost s-orbital. Because an s-orbital holds at most two electrons, the block is just two groups wide.
- Group 1 - Alkali metals: Li, Na, K, Rb, Cs, Fr. General configuration ns¹.
- Group 2 - Alkaline earth metals: Be, Mg, Ca, Sr, Ba, Ra. General configuration ns².
They are called “alkali” metals because their hydroxides are strong alkalis, and “alkaline earth” metals because their oxides are alkaline and were found in the earth’s crust. Hydrogen, despite being 1s¹, is not a true alkali metal - it is a non-metal placed separately.
2. Electronic Configuration
Every s-block element has a noble-gas core plus one or two outer s-electrons, which it readily loses to form +1 or +2 ions.
| Element | Configuration | Common Ion |
|---|---|---|
| Li | [He] 2s¹ | Li⁺ |
| Na | [Ne] 3s¹ | Na⁺ |
| K | [Ar] 4s¹ | K⁺ |
| Be | [He] 2s² | Be²⁺ |
| Mg | [Ne] 3s² | Mg²⁺ |
| Ca | [Ar] 4s² | Ca²⁺ |
Key idea: Losing the loosely held ns electron(s) gives a stable noble-gas configuration, which is why these metals are so reactive and almost always found as +1 or +2 ions in nature.
3. Trends in Atomic and Ionic Radii
Down a group, a new shell is added at each step, so both atomic and ionic radii increase down Group 1 and Group 2.
- Alkali metals have the largest atomic radii in their respective periods.
- For the same period, a Group 2 atom is smaller than the corresponding Group 1 atom, because Group 2 has a higher nuclear charge pulling the same shell inward.
- Cations are always smaller than their parent atoms (Na⁺ < Na) because a whole shell is lost and the remaining electrons feel a stronger effective nuclear charge.
4. Trends in Ionization Enthalpy
Ionization enthalpy is the energy needed to remove the outermost electron from a gaseous atom. Since the outer electron gets farther from the nucleus down a group, ionization enthalpy decreases down both groups.
- Alkali metals have the lowest ionization enthalpies in their periods - they lose one electron very easily.
- Group 2 elements have higher first ionization enthalpies than Group 1 (smaller size, higher nuclear charge).
- However, the second ionization enthalpy of Group 2 is far lower than that of Group 1, because removing a second electron from Group 1 means breaking into a stable noble-gas core - so Group 2 readily forms +2 ions.
5. Hydration Enthalpy
Hydration enthalpy is the energy released when one mole of gaseous ions dissolves in water and gets surrounded by water molecules. Smaller, more highly charged ions are hydrated more strongly.
- Hydration enthalpy decreases down a group as ionic size increases: Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺.
- Group 2 ions (M²⁺) have larger hydration enthalpies than Group 1 ions (M⁺) of comparable size because of their higher charge.
- Li⁺ is the most heavily hydrated ion, so hydrated Li⁺ is effectively the largest - that is why lithium is the least mobile alkali metal ion in solution despite being the smallest atom.
6. Physical Properties
s-Block metals are soft, light, silvery and excellent conductors, with low melting and boiling points that fall down the group as the metallic bond weakens.
- Softness: alkali metals are so soft they can be cut with a knife; Group 2 metals are harder.
- Density: Li, Na and K are lighter than water; Group 2 metals are denser than Group 1.
- Flame colour: their loosely held electrons are easily excited, giving characteristic flame colours - Li (crimson red), Na (golden yellow), K (lilac/violet), Ca (brick red), Sr (crimson), Ba (apple green). Be and Mg do not impart colour (electrons too tightly bound).
7. Chemical Reactivity
Reactivity increases down both groups because ionization enthalpy falls. Group 1 metals are more reactive than Group 2 metals.
Reaction with Water
Alkali metals react vigorously with water to give a hydroxide and hydrogen gas:
2Na + 2H₂O → 2NaOH + H₂↑
The reaction gets more violent down the group (K catches fire, Cs explodes). Group 2 metals react less readily - Be does not react, Mg reacts only with hot water/steam, while Ca, Sr and Ba react with cold water:
Ca + 2H₂O → Ca(OH)₂ + H₂↑
Reaction with Oxygen
The product of burning depends on the metal - a favourite exam point:
- Lithium forms mainly the oxide (Li₂O).
- Sodium forms mainly the peroxide (Na₂O₂).
- K, Rb, Cs form superoxides (e.g. KO₂).
- Group 2 metals form normal oxides (MO); Ba can form BaO₂.
Reaction with Hydrogen
Except Be, all form ionic (saline) hydrides on heating:
2Na + H₂ → 2NaH Ca + H₂ → CaH₂
These hydrides contain the H⁻ ion and act as strong reducing agents.
8. Anomalous Behaviour of Lithium and Beryllium
The first member of each group differs from the rest because of its very small size, high charge density (high polarising power), and absence of d-orbitals.
Anomalies of Lithium
- Li is much harder, with the highest melting/boiling point of the alkali metals.
- LiCl is deliquescent and somewhat covalent; Li forms only the oxide (not peroxide/superoxide).
- Li₂CO₃, LiOH, LiF and Li₃PO₄ are sparingly soluble - unlike the very soluble salts of Na, K.
- Li reacts directly with N₂ to form Li₃N (other alkali metals do not).
Anomalies of Beryllium
- Be is harder, has a high melting point and forms predominantly covalent compounds (e.g. BeCl₂ is covalent and polymeric).
- Be does not react with water; its oxide and hydroxide are amphoteric, while those of other Group 2 metals are basic.
- Be forms complexes such as [BeF₄]²⁻ owing to its small size and high charge.
9. Diagonal Relationship
Because moving right increases charge density and moving down decreases it, the first element of a group resembles the second element of the next group placed diagonally. This is the diagonal relationship.
- Li resembles Mg: both form nitrides with N₂, their carbonates decompose on heating, both have sparingly soluble fluorides and carbonates, and both form covalent organometallic compounds.
- Be resembles Al: both are amphoteric, both form covalent halides that dimerise/polymerise, and both are passivated by concentrated HNO₃.
The cause is the similar charge-to-radius ratio (polarising power) of the diagonally placed elements.
10. Important Compounds of Sodium
Sodium Hydroxide (NaOH) - Caustic Soda
Made by the electrolysis of brine (Castner–Kellner process). It is a deliquescent white solid, strongly alkaline, used in soap, paper, and the petroleum industry.
2NaCl + 2H₂O → 2NaOH + Cl₂ + H₂
Sodium Carbonate (Na₂CO₃·10H₂O) - Washing Soda
Made by the Solvay (ammonia-soda) process. Used in water softening, glass, soap and detergents. On heating, washing soda loses water of crystallisation to give the white anhydrous powder (soda ash).
NaCl + NH₃ + CO₂ + H₂O → NaHCO₃ + NH₄Cl; 2NaHCO₃ → Na₂CO₃ + CO₂ + H₂O
Sodium Bicarbonate (NaHCO₃) - Baking Soda
A mild, non-toxic alkali used in baking (releases CO₂), as an antacid, and in fire extinguishers. It is a Solvay-process intermediate.
11. Important Compounds of Calcium
Calcium Oxide (CaO) - Quick Lime
Made by heating limestone: CaCO₃ → CaO + CO₂. Reacts vigorously with water (slaking) to give slaked lime, Ca(OH)₂. Used in cement, mortar and as a drying agent.
Calcium Carbonate (CaCO₃) - Limestone/Marble
Occurs as marble, chalk and limestone. Used in cement manufacture, as a building stone, and an antacid. It decomposes on heating to give CaO.
Plaster of Paris - CaSO₄·½H₂O
Made by heating gypsum (CaSO₄·2H₂O) to about 393 K (120 °C):
2(CaSO₄·2H₂O) → 2(CaSO₄·½H₂O) + 3H₂O
When mixed with water it sets into a hard mass of gypsum, expanding slightly - which is why it is used for plaster casts (fractured bones), moulds, statues and blackboard chalk.
12. Biological Importance
s-Block ions are essential to life - a small but exam-frequent topic.
- Sodium (Na⁺): the chief extracellular cation; controls fluid balance, blood pressure, and nerve-impulse transmission (the sodium–potassium pump).
- Potassium (K⁺): the chief intracellular cation; needed for nerve transmission, muscle contraction and enzyme activation.
- Magnesium (Mg²⁺): the central metal ion in chlorophyll; activates many enzymes and stabilises DNA/ATP.
- Calcium (Ca²⁺): builds bones and teeth (as phosphate), and is vital for blood clotting, muscle contraction and nerve signalling.
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Weightage in Board & Entrance Exams
| Exam | Typical Weightage | Most-Tested Areas |
|---|---|---|
| CBSE Board (Class 11) | 6–8 marks | Periodic trends, anomalies of Li/Be, diagonal relationship, important compounds |
| JEE Main | 1–2 questions | Reactions with O₂/water, oxides vs peroxides vs superoxides, properties of compounds |
| NEET | 1–2 questions | Biological importance, hydration enthalpy, flame colours, plaster of Paris |
[TABLE: Question-type split - VSA (1 mark): definitions, flame colours, formulae; SA (2–3 marks): trends, anomalies, diagonal relationship; LA (5 marks): preparation/uses of NaOH, Na₂CO₃, plaster of Paris with equations.]
Important Definitions
| Term | Definition |
|---|---|
| s-Block element | Element whose last electron enters the outermost s-orbital (Groups 1 and 2) |
| Alkali metals | Group 1 elements (ns¹) whose hydroxides are strong alkalis |
| Alkaline earth metals | Group 2 elements (ns²) whose oxides are alkaline |
| Ionization enthalpy | Energy needed to remove the outermost electron from a gaseous atom |
| Hydration enthalpy | Energy released when one mole of gaseous ions is surrounded by water |
| Diagonal relationship | Resemblance of an element to the diagonally placed element of the next group (Li–Mg, Be–Al) |
| Superoxide | Compound containing the O₂⁻ ion, e.g. KO₂ |
| Slaking of lime | Reaction of CaO with water to give Ca(OH)₂ |
| Plaster of Paris | CaSO₄·½H₂O, made by partially dehydrating gypsum |
| Polarising power | Ability of a cation to distort an anion; high for small, highly charged ions |
Solved Examples
Example 1
Write the products formed when Li, Na and K are each burnt in excess oxygen.
Answer: Li → Li₂O (oxide); Na → Na₂O₂ (peroxide); K → KO₂ (superoxide). The tendency to form larger anions increases as the cation gets larger.
Example 2
Why is lithium the strongest reducing agent among alkali metals in aqueous solution despite having the highest ionization enthalpy in the group?
Answer: Reducing power in solution depends on the standard electrode potential, which includes the very high hydration enthalpy of the small Li⁺ ion. The large energy released on hydrating Li⁺ outweighs its high ionization enthalpy, making Li the strongest reducing agent in water.
Example 3
Identify the gas evolved and the salt formed when sodium reacts with water. Write the equation.
Answer: Gas evolved is hydrogen (H₂) and the salt is sodium hydroxide. 2Na + 2H₂O → 2NaOH + H₂↑.
Example 4
Give two reasons for the diagonal relationship between lithium and magnesium.
Answer: (i) Both have nearly the same charge-to-radius ratio (polarising power). (ii) Both react directly with N₂ to form nitrides (Li₃N, Mg₃N₂) and both have sparingly soluble carbonates that decompose on heating.
Example 5
How is plaster of Paris prepared, and why must the temperature be controlled?
Answer: Gypsum is heated to ~393 K: 2(CaSO₄·2H₂O) → 2(CaSO₄·½H₂O) + 3H₂O. If heated above 393 K, all water is lost giving anhydrous CaSO₄ (dead burnt plaster), which no longer sets with water.
Example 6
Arrange Li⁺, Na⁺, K⁺ in increasing order of (a) ionic radius and (b) hydration enthalpy.
Answer: (a) Ionic radius: Li⁺ < Na⁺ < K⁺. (b) Hydration enthalpy: K⁺ < Na⁺ < Li⁺ (smaller ion is hydrated more strongly).
Important Questions for Board Exams
1-Mark Questions (VSA)
- Why are Group 1 elements called alkali metals?
- Name the alkali metal that forms a nitride directly with nitrogen.
- What is the flame colour imparted by potassium?
- Write the formula of plaster of Paris.
- Which Group 2 element forms an amphoteric oxide?
2–3-Mark Questions (SA)
- Explain why ionization enthalpy decreases down Group 1 but the second ionization enthalpy of Group 2 is much lower than that of Group 1.
- Discuss the anomalous behaviour of beryllium in Group 2 with two examples.
- What is the diagonal relationship? Illustrate it with the Li–Mg pair.
- Why is Li⁺ the most heavily hydrated alkali metal ion? What is one consequence?
5-Mark Questions (LA)
- Describe the manufacture of sodium carbonate by the Solvay process, giving the equations and two uses.
- How is plaster of Paris prepared from gypsum? Give the equation, its setting reaction and two uses.
- Compare Group 1 and Group 2 elements with respect to atomic radii, ionization enthalpy, reactivity with water, and nature of oxides.
Quick Revision Points
- Group 1 = ns¹ (alkali metals); Group 2 = ns² (alkaline earth metals)
- Down the group: atomic/ionic radii ↑, ionization enthalpy ↓, reactivity ↑, hydration enthalpy ↓
- Burning in O₂: Li → oxide, Na → peroxide, K/Rb/Cs → superoxide
- Reaction with water: 2Na + 2H₂O → 2NaOH + H₂; Be no reaction, Mg with steam, Ca/Sr/Ba with cold water
- Anomalies: Li (covalent LiCl, forms Li₃N, sparingly soluble salts); Be (covalent, amphoteric oxide, forms complexes)
- Diagonal relationship: Li ~ Mg, Be ~ Al (same polarising power)
- Key sodium compounds: NaOH (caustic soda), Na₂CO₃·10H₂O (washing soda), NaHCO₃ (baking soda)
- Key calcium compounds: CaO (quick lime), CaCO₃ (limestone), CaSO₄·½H₂O (plaster of Paris)
- Biology: Na⁺/K⁺ → nerve impulses; Mg²⁺ → chlorophyll; Ca²⁺ → bones, clotting
- Flame colours: Li crimson, Na golden yellow, K lilac, Ca brick red, Ba apple green
Next Chapter: Chapter 8 - Hydrogen | Continue to Chapter 10 - The p-Block Elements
Class 11 Chemistry · Chapter 10 – swipe through all 9 cards to understand the whole chapter.
Meet the s-Block
The s-block is two columns of soft, reactive metals that lose their loose outer electron with ease.
Group 1: Li Na K Rb Cs Fr. Group 2: Be Mg Ca Sr Ba Ra.
- One/two electrons just outside a noble-gas core
- Low melting points, soft enough to cut with a knife
- Every trend traces back to how tightly the nucleus holds that loose electron
Size, IE & Hydration Down a Group
Going down a group, atoms swell, the outer electron escapes more easily, and ions hold water less tightly.
Smallest ion = highest charge density = most heavily hydrated (Li⁺).
- New shell each period → atomic & ionic radius increases
- Group 1 has the lowest first IE in the periodic table
- 2nd IE of Group 1 is huge (breaks the noble-gas core)
Across the Step: Na → Mg
Move one step right and the metal gets smaller, denser, harder, with higher IE and melting point.
Density rises down a group, but K is anomalously lighter than Na.
- Higher nuclear charge squeezes Group 2 electrons inward
- Two delocalised electrons make a tougher metallic bond
- Larger size → lower IE → greater reducing character down the group
Reaction with Water
Reactivity climbs down the group as IE falls; Li is calm, Cs is explosive.
Group 2 is milder: Be no reaction · Mg with steam · Ca, Sr, Ba with cold water.
- Metal donates electron to water, releasing H2 gas
- Alkali metals react vigorously; reactivity rises Li → Cs
- Beryllium does not react even with steam
Reaction with Oxygen: The Size Rule
Which oxide forms depends on cation size; a big cation can cradle a big, fragile oxygen anion.
‘Big cation stabilises big anion’ recurs through the whole chapter.
- Small Li⁺ only supports small O2⁻ → simple oxide
- With H2 they form ionic hydrides: NaH, CaH2 (hydrolith)
- Memorise the ladder: oxide → peroxide → superoxide
Flame Colours
Loosely-held electrons get excited by heat and emit visible light as they relax back.
Be and Mg give NO colour — electrons too tightly held, need UV.
- Energy gap fixes the colour emitted
- View through blue cobalt glass to spot K behind Na’s yellow
- Be/Mg are THE flame-test exceptions examiners reuse
Hydroxides, Carbonates & Sulphates
Solubility goes up or down a group depending on whether the anion is small or large (lattice vs hydration tug-of-war).
Small anion (OH⁻) → solubility ↑ down; large anion (SO42⁻, CO32⁻) → solubility ↓ down.
- Be(OH)2 is amphoteric; the rest are basic
- Sulphates: BeSO4, MgSO4 soluble → SrSO4, BaSO4 insoluble (BaSO4 = barium meal)
- Carbonate thermal stability ↑ down group: BeCO3 < ... < BaCO3
Anomalous Li, Be & Diagonal Twins
Tiny size and high polarising power (charge ÷ size) make Li and Be misbehave and twin diagonally.
Li, Be lack d-orbitals → extra covalent character.
- Li: hardest, highest m.p., least reactive in Group 1; forms only Li2O
- Be & Al: amphoteric oxides, covalent Lewis-acidic chlorides, passivated by conc. HNO3
- Li is the strongest reducing agent in solution due to huge HYDRATION enthalpy, not low IE
Key Compounds of Na, Ca & Mg
A handful of compounds run the kitchen, hospital and factory; tie each to name, prep and one reaction.
Solvay makes Na2CO3 but NOT K2CO3 (KHCO3 too soluble to precipitate).
- NaOH (caustic soda) via Castner-Kellner/chlor-alkali; 2NaHCO3 —Δ→ Na2CO3 + H2O + CO2
- CaCO3 —Δ→ CaO (quicklime) → Ca(OH)2 (slaked lime/lime water turns milky with CO2)
- Biology: Mg2⁺ → chlorophyll, Ca2⁺ → bones/clotting, Na⁺/K⁺ → nerve impulse
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Frequently Asked Questions
The s-block elements are those in which the last electron enters an s-orbital, comprising Group 1 (alkali metals: Li, Na, K, Rb, Cs, Fr) with an ns1 configuration and Group 2 (alkaline earth metals: Be, Mg, Ca, Sr, Ba, Ra) with an ns2 configuration. They are soft, low-melting, highly reactive metals that readily lose their outer electrons to form positive ions.
The product depends on cation size, because a large cation can stabilise a large anion. Lithium forms only the oxide (Li2O), sodium forms the peroxide (Na2O2), and the larger potassium, rubidium and caesium form superoxides (such as KO2).
Alkali metals react vigorously with water to give the hydroxide and hydrogen gas, for example 2Na plus 2H2O gives 2NaOH plus H2, with reactivity rising down the group from lithium (calm) to caesium (explosive). Group 2 metals are milder: beryllium does not react, magnesium reacts only with steam, and calcium, strontium and barium react with cold water.
Lithium and beryllium are anomalous because of their very small size, high polarising power (charge to size ratio) and absence of d-orbitals, which gives their compounds extra covalent character. This makes lithium resemble magnesium and beryllium resemble aluminium, a pattern called the diagonal relationship.
Yes, it is a high-scoring part of NEET inorganic chemistry that usually contributes direct fact-based questions. The most frequently tested points are characteristic flame colours, the oxide-peroxide-superoxide rule, anomalous behaviour of Li and Be, and named compounds like Plaster of Paris and washing soda.