s-Block Elements — Alkali & Alkaline Earth Metals

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Groups 1 and 2 are the most predictable groups in the periodic table: one or two loosely held s electrons, large atoms, strong reducing power. The marks come from the few places where the trend bends — lithium and beryllium, hydration enthalpy, and the solubility and stability orders of their salts.

What you will be able to do

• Explain trends in size, ionisation enthalpy, hydration and reducing power down groups 1 and 2
• Predict the oxide formed by each alkali metal with oxygen and explain solutions in liquid ammonia
• Explain the anomalous behaviour of Li and Be and their diagonal relationships
• Order thermal stability and solubility of carbonates, sulphates and hydroxides, with reasons
• Write the preparation and uses of NaOH, Na₂CO₃, NaHCO₃, CaO and plaster of Paris

The big picture

Easy-to-lose s electrons drive everything.

In one line: Group 1 (ns¹) and group 2 (ns²) metals lose their s electrons easily to form M⁺ and M²⁺ ions; they are strong reducing agents whose reactivity rises down the group, while the small first members Li and Be behave differently.

Property	Group 1 (Li → Cs)	Group 2 (Be → Ba)
Configuration	[noble gas] ns¹	[noble gas] ns²
Common ion	M⁺	M²⁺
Atomic size	Largest in each period; increases down	Smaller than group 1; increases down
Ionisation enthalpy	Lowest in each period; decreases down	Higher than group 1; decreases down
Hydration enthalpy	Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺	Be²⁺ > Mg²⁺ > Ca²⁺ > Sr²⁺ > Ba²⁺
Metallic character	Soft; low m.p.	Harder; higher m.p.

Think of it like this: The single outer electron of an alkali metal is like a loose coin in a large pocket: the bigger the pocket (atom), the easier it falls out. Group 2 has two coins in a slightly smaller, tighter pocket — still easy to lose, but it takes more effort.

Physical trends

Size, colour and motion in water.

Group 1	Li	Na	K	Rb	Cs
Metallic radius / pm	152	186	227	248	265
Ionic radius M⁺ / pm	76	102	138	152	167
ΔiH₁ / kJ mol⁻¹	520	496	419	403	376
Hydration enthalpy / kJ mol⁻¹	−506	−406	−330	−310	−276
E° (M⁺/M) / V	−3.04	−2.71	−2.93	−2.93	−2.93
m.p. / K	454	371	336	312	302

Group 2	Be	Mg	Ca	Sr	Ba
Metallic radius / pm	112	160	197	215	222
Ionic radius M²⁺ / pm	31	72	100	118	135
ΔiH₁ / kJ mol⁻¹	899	737	590	549	503
Hydration enthalpy / kJ mol⁻¹	−2494	−1921	−1577	−1443	−1305
E° (M²⁺/M) / V	−1.97	−2.36	−2.84	−2.89	−2.92

Key idea: Read the tables together. Group 2 ions have hydration enthalpies four to five times those of group 1, because they carry twice the charge on a smaller ion. That is why group 2 compounds are more often hydrated (MgCl₂·6H₂O, CaCl₂·6H₂O) while group 1 compounds mostly are not (except Li).

• Melting points fall down group 1 (Li 454 K to Cs 302 K) because metallic bonding weakens as atoms grow.
• Flame colours: the energy from the flame excites the outer electron, which emits visible light as it falls back.
• Hydrated radius is the reverse of ionic radius: small Li⁺ holds the most water, so in water it is the largest hydrated ion and the least mobile. Ionic mobility and conductance in water: Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺.
• Cs and K emit electrons when light falls on them, so they are used in photoelectric cells.

Element	Li	Na	K	Rb	Cs	Ca	Sr	Ba
Flame colour	Crimson red	Golden yellow	Violet (lilac)	Red-violet	Blue	Brick red	Crimson	Apple green

Key idea: Be and Mg give no flame colour. Their electrons are held too tightly for the flame to excite them into the visible range.

Chemical reactions

Strong reducing agents meet oxygen, water and ammonia.

With oxygen

Metal	Main product in excess air	Oxide type
Li	Li₂O	Normal oxide (O²⁻)
Na	Na₂O₂	Peroxide (O₂²⁻)
K, Rb, Cs	KO₂, RbO₂, CsO₂	Superoxide (O₂⁻), paramagnetic and coloured

Larger cations stabilise larger anions: the small Li⁺ cannot stabilise the big peroxide or superoxide ion; K⁺ and heavier ions can.

With water

2M + 2H₂O → 2MOH + H₂

Reactivity with water increases down the group: Li slowly, Na vigorously, K catches fire, Cs explodes

Key idea: Li has the most negative E° (−3.04 V), yet reacts most slowly with water. E° reflects the overall energy, which is dominated by Li⁺’s huge hydration enthalpy. The speed depends on kinetics — Li has a high melting point and does not melt on the water surface, so it reacts more gently.

In liquid ammonia

• Alkali metals dissolve to give blue solutions containing ammoniated cations and ammoniated electrons: M + (x + y)NH₃ → [M(NH₃)ₓ]⁺ + [e(NH₃)ᵧ]⁻.
• Dilute solutions are blue, paramagnetic and good conductors (the free electrons). Concentrated solutions turn bronze-coloured and diamagnetic as the electrons pair up.
• The solution is a powerful reducing agent (Birch reduction). On standing it slowly gives the amide and H₂.

Solved example 01 — Superoxide in breathing masks

Why is KO₂ used in space suits and submarine breathing apparatus? Write the reaction.

1. Exhaled CO₂ reacts with KO₂ and releases oxygen.
2. 4KO₂ + 2CO₂ → 2K₂CO₃ + 3O₂.

Answer: It absorbs CO₂ and regenerates O₂ at the same time.

Solved example 02 — Peroxide with water and CO₂

Write the reactions of Na₂O₂ with water and with CO₂.

1. With water: Na₂O₂ + 2H₂O → 2NaOH + H₂O₂ (in cold water; warm water gives O₂).
2. With CO₂: 2Na₂O₂ + 2CO₂ → 2Na₂CO₃ + O₂.

Answer: Na₂O₂ gives H₂O₂ with water and releases O₂ with CO₂ — another way to purify air in closed spaces.

Solved example 03 — Ordering conductance

Arrange LiCl, NaCl, KCl and CsCl (same concentration in water) in order of increasing molar conductance, and explain.

1. Conductance depends on how fast the cation moves through water.
2. Li⁺ is the most hydrated, so it is the largest moving unit and the slowest.
3. Cs⁺ is the least hydrated and the fastest.

Answer: LiCl < NaCl < KCl < CsCl.

Check yourself

• C1. Which alkali metal forms a normal oxide in excess air?
• C2. Why does a concentrated solution of Na in liquid NH₃ become diamagnetic?

Answers at the end.

Hydrides, halides and salts

The compounds group 1 and 2 metals make with H, X, CO₃²⁻ and NO₃⁻.

Hydrides

• Alkali metals heated with H₂ give ionic (saline) hydrides M⁺H⁻. Their stability decreases down the group: LiH > NaH > KH > RbH > CsH (lattice energy falls).
• They react with water to give H₂: NaH + H₂O → NaOH + H₂. LiH reacts with Al₂Cl₆ to give LiAlH₄, the reducing agent of organic chemistry.
• Group 2: all except Be form ionic hydrides directly. BeH₂ is made indirectly (BeCl₂ + LiAlH₄) and is a covalent, polymeric hydride with Be–H–Be bridges; MgH₂ is also partly covalent.

Halides

• Group 1 halides MX are ionic solids with high melting points, except that Li halides are partly covalent (LiI most covalent).
• Melting points fall in the order fluoride > chloride > bromide > iodide for a given metal.
• Group 2 halides are mostly ionic; BeCl₂ is covalent. Group 2 chlorides crystallise as hydrates; heating hydrated MgCl₂·6H₂O does not give anhydrous MgCl₂ — it hydrolyses to MgO (or Mg(OH)Cl) and HCl.
• CaF₂ is almost insoluble: F⁻ is small, so its lattice energy with Ca²⁺ is very high. Other group 2 fluorides are also poorly soluble, while the chlorides are soluble.

Nitrates and carbonates on heating

Salt	On heating
Group 1 nitrates (Na → Cs)	2MNO₃ → 2MNO₂ + O₂
LiNO₃ and group 2 nitrates	4LiNO₃ → 2Li₂O + 4NO₂ + O₂; 2M(NO₃)₂ → 2MO + 4NO₂ + O₂
Group 1 carbonates (Na → Cs)	Stable — melt without decomposing
Li₂CO₃ and group 2 carbonates	MCO₃ → MO + CO₂ (easier for smaller cations)
Group 1 bicarbonates	2MHCO₃ → M₂CO₃ + H₂O + CO₂

Solved example 04 — Hydride with water

Write the reactions of CaH₂ and LiH with water, and explain why CaH₂ is used to make H₂ in the field.

1. CaH₂ + 2H₂O → Ca(OH)₂ + 2H₂.
2. LiH + H₂O → LiOH + H₂.
3. CaH₂ is a stable solid that releases two moles of H₂ per mole and is easy to carry (“hydrolith”).

Answer: Ionic hydrides give H₂ and the hydroxide with water; CaH₂ is a portable source of hydrogen.

Check yourself

• C3. Arrange LiH, NaH and KH in order of decreasing thermal stability.
• C4. What does NaNO₃ give on heating? And LiNO₃?

Answers at the end.

Lithium: the odd one out

Small size and high polarising power.

Li⁺ is very small with high charge density, so it polarises anions strongly and its compounds have covalent character. Its chemistry resembles magnesium’s (diagonal relationship).

Property	Lithium (and Mg)	Other alkali metals
Reaction with N₂	Forms Li₃N (Mg₃N₂ too)	No nitride
Carbonate on heating	Li₂CO₃ → Li₂O + CO₂ (MgCO₃ too)	Stable to heat
Nitrate on heating	4LiNO₃ → 2Li₂O + 4NO₂ + O₂ (like Mg)	2NaNO₃ → 2NaNO₂ + O₂
Hydrogencarbonate	LiHCO₃ not obtained as a solid	NaHCO₃, KHCO₃ are solids
Halides	LiCl is deliquescent, soluble in ethanol (covalent)	Ionic
Oxide in air	Li₂O only	Peroxides, superoxides

Solved example 05 — Why lithium salts are hydrated

LiCl crystallises as LiCl·2H₂O but NaCl is anhydrous. Explain.

1. Li⁺ is the smallest alkali metal ion (76 pm), with the highest charge density.
2. It attracts water molecules strongly (hydration enthalpy −506 kJ mol⁻¹ against −406 for Na⁺).
3. So water of crystallisation is held in the solid for Li salts but not for Na salts.

Answer: Small, highly hydrated Li⁺ holds water of crystallisation; Na⁺ does not.

Group 2 and beryllium

Two electrons, higher charge, smaller ions.

• Group 2 metals are harder, denser and higher-melting than group 1, because two electrons take part in metallic bonding.
• They are less reactive: Be and Mg are protected by an oxide film; Mg reacts with hot water or steam; Ca, Sr and Ba react with cold water.
• All burn in air to give MO (and nitride M₃N₂); Mg burns with a dazzling white flame, used in flash photography and fireworks.

Reactions of group 2 metals

Reagent	Product	Note
O₂ (burning)	MO	Ba also forms BaO₂ (peroxide)
N₂ (heating)	M₃N₂	M₃N₂ + 6H₂O → 3M(OH)₂ + 2NH₃
H₂O	M(OH)₂ + H₂	Be none; Mg with hot water; Ca, Sr, Ba with cold water
Dilute acids	M²⁺ salt + H₂	Be dissolves slowly; Be is passive to conc. HNO₃
Halogens	MX₂	BeCl₂ also made from BeO + C + Cl₂ at 600–800 K
Liquid NH₃	Blue solutions, like group 1	Ammoniates [M(NH₃)₆]²⁺ form

Beryllium is anomalous

• Be²⁺ is tiny and highly polarising, so Be compounds are largely covalent.
• BeCl₂ is a polymeric chain in the solid (each Be bridged by two Cl), a dimer in the vapour, and a linear monomer above about 1200 K.
• Be(OH)₂ and BeO are amphoteric: Be(OH)₂ + 2OH⁻ → [Be(OH)₄]²⁻.
• Maximum covalency of Be is 4 (no d orbitals), e.g. [Be(OH)₄]²⁻; Mg can reach 6.
• Diagonal relationship with Al: both are rendered passive by conc. HNO₃, form amphoteric oxides and hydroxides, form covalent chlorides that are Lewis acids and fume in moist air, and give complexes such as [BeF₄]²⁻ and [AlF₆]³⁻.

Stability and solubility

Two competing energies decide every order.

Key idea: Solubility = hydration enthalpy versus lattice enthalpy. Down a group both fall. If the anion is large (sulphate, carbonate), lattice enthalpy hardly changes, so falling hydration enthalpy wins → solubility decreases. If the anion is small (OH⁻, F⁻), lattice enthalpy falls faster → solubility increases.

Compound class (group 2)	Trend down the group	Example
Sulphates	Solubility decreases	BeSO₄, MgSO₄ soluble; BaSO₄ insoluble
Hydroxides	Solubility and basic strength increase	Be(OH)₂ amphoteric; Ba(OH)₂ soluble, strong base
Carbonates	Thermal stability increases	BeCO₃ < MgCO₃ < CaCO₃ < SrCO₃ < BaCO₃
Nitrates	All decompose to MO + NO₂ + O₂	Easier for smaller cations

• Thermal stability of carbonates increases down both groups: a small, highly polarising cation distorts the large carbonate ion and pulls it apart into CO₂ + O²⁻. Group 1 carbonates (except Li₂CO₃) are stable to heat.
• Group 1 hydroxides are all strong bases, strength increasing down: LiOH < NaOH < KOH < RbOH < CsOH.
• Group 1 carbonates and bicarbonates are soluble (Li₂CO₃ only slightly); solubility increases down the group.

Solved example 06 — Explaining a solubility order

Explain why BaSO₄ is insoluble while MgSO₄ is soluble in water.

1. SO₄²⁻ is a large anion, so the lattice enthalpy changes little from MgSO₄ to BaSO₄.
2. The hydration enthalpy falls sharply as the cation grows (Mg²⁺ → Ba²⁺).
3. For BaSO₄ the hydration enthalpy can no longer overcome the lattice enthalpy.

Answer: Hydration enthalpy drops faster than lattice enthalpy down the group.

Solved example 07 — Explaining a stability order

Why does BeCO₃ decompose at a much lower temperature than BaCO₃?

1. Be²⁺ is tiny (31 pm) with a very high charge density; Ba²⁺ is large (135 pm).
2. Be²⁺ strongly polarises the carbonate ion, weakening a C–O bond and favouring loss of CO₂.
3. The product BeO also has a very high lattice energy, which helps the decomposition.

Answer: Small, highly polarising Be²⁺ destabilises the carbonate; thermal stability increases down the group.

Solved example 08 — Hydroxide solubility

Mg(OH)₂ is sparingly soluble but Ba(OH)₂ dissolves readily. Explain.

1. OH⁻ is small, so lattice enthalpy depends strongly on cation size.
2. From Mg²⁺ to Ba²⁺ the lattice enthalpy falls faster than the hydration enthalpy.

Answer: Solubility of group 2 hydroxides increases down the group; Ba(OH)₂ is soluble and a strong base.

Check yourself

• C5. Arrange BeCO₃, CaCO₃ and BaCO₃ in increasing thermal stability.
• C6. Which is more basic: Mg(OH)₂ or Ba(OH)₂?

Answers at the end.

Sodium compounds

The industrial chemicals of this chapter.

Compound	Preparation	Key facts and uses
NaOH (caustic soda)	Electrolysis of brine in the Castner–Kellner cell (mercury cathode): Na amalgam + water → NaOH + H₂	Deliquescent; soap, paper, petroleum refining
Na₂CO₃·10H₂O (washing soda)	Solvay (ammonia–soda) process	Effloresces to Na₂CO₃·H₂O; above 373 K → anhydrous soda ash; water softening, glass
NaHCO₃ (baking soda)	Intermediate of the Solvay process	On heating: 2NaHCO₃ → Na₂CO₃ + H₂O + CO₂; antacid, fire extinguishers, baking
NaCl	Evaporation of sea water; purified with HCl gas (common ion)	Impure salt is deliquescent because of MgCl₂ and CaCl₂

Solvay process

1. 2NH₃ + H₂O + CO₂ → (NH₄)₂CO₃
2. (NH₄)₂CO₃ + H₂O + CO₂ → 2NH₄HCO₃
3. NH₄HCO₃ + NaCl → NH₄Cl + NaHCO₃↓ (sparingly soluble, crystallises)
4. 2NaHCO₃ → Na₂CO₃ + CO₂ + H₂O (on heating); NH₃ recovered from NH₄Cl with Ca(OH)₂

Key idea: K₂CO₃ cannot be made by the Solvay process. KHCO₃ is too soluble to crystallise out in step 3.

Calcium compounds

From limestone to the plaster on your wall.

Compound	Preparation	Key facts
CaO (quicklime)	Heat limestone at 1070–1270 K: CaCO₃ → CaO + CO₂	CO₂ removed to push equilibrium forward; basic oxide
Ca(OH)₂ (slaked lime)	CaO + H₂O → Ca(OH)₂	Clear solution = lime water; suspension = milk of lime; turns milky with CO₂ (CaCO₃), clear again with excess CO₂ (Ca(HCO₃)₂)
Bleaching powder	Cl₂ on slaked lime: 2Ca(OH)₂ + 2Cl₂ → CaCl₂ + Ca(OCl)₂ + 2H₂O	Disinfectant and bleach (via OCl⁻)
CaSO₄·½H₂O (plaster of Paris)	Heat gypsum CaSO₄·2H₂O at 393 K	Sets with water back to gypsum; above 393 K gives anhydrous “dead burnt plaster” that does not set

• Cement is a mixture of calcium silicates and aluminates (CaO about 50–60%); gypsum is added to slow setting.

Solved example 09 — Plaster of Paris

Why does plaster of Paris harden when mixed with water, and why must gypsum not be overheated when making it?

1. Setting: 2CaSO₄·½H₂O + 3H₂O → 2CaSO₄·2H₂O — gypsum crystals interlock as a hard mass, with slight expansion.
2. Above 393 K all water of crystallisation is lost: anhydrous CaSO₄ (dead burnt plaster) cannot rehydrate quickly.

Answer: Rehydration to gypsum makes it set; overheating destroys this ability.

More compounds you should know

Magnesium salts, calcium carbonate and cement in detail.

Compound	Formula	Use or fact
Epsom salt	MgSO₄·7H₂O	Laxative; sulphates of Be and Mg are soluble
Magnesia	MgO	Refractory lining of furnaces (very high m.p.)
Milk of magnesia	Mg(OH)₂ suspension	Antacid
Dolomite	CaCO₃·MgCO₃	Source of Mg; used in construction
Calcium carbonate	CaCO₃	Limestone, chalk, marble; Ca(OH)₂ + CO₂ → CaCO₃ + H₂O, or CaCl₂ + Na₂CO₃ → CaCO₃ + 2NaCl
Calcium chloride (anhydrous)	CaCl₂	Drying agent (not for NH₃ or alcohols, with which it forms adducts)

Cement

Component	Approximate %
CaO	50–60
SiO₂	20–25
Al₂O₃	5–10
MgO	2–3
Fe₂O₃	1–2
SO₃	1–2

• For good cement, the ratio SiO₂ : Al₂O₃ is between 2.5 and 4, and the ratio of CaO to the total of SiO₂, Al₂O₃ and Fe₂O₃ is close to 2.
• Raw materials: limestone and clay, heated in a rotary kiln to give clinker; about 2–3% gypsum is ground in to slow down setting.
• Main constituents of clinker: dicalcium silicate (Ca₂SiO₄, 26%), tricalcium silicate (Ca₃SiO₅, 51%) and tricalcium aluminate (Ca₃Al₂O₆, 11%).

Castner–Kellner cell in equations

Cathode (Hg): Na⁺ + e⁻ + Hg → Na–amalgam
Anode (graphite): Cl⁻ → ½Cl₂ + e⁻
Denuder: 2Na–amalgam + 2H₂O → 2NaOH + 2Hg + H₂

The mercury cathode makes Na discharge (not H₂) because of hydrogen’s high overvoltage on Hg

• NaOH absorbs CO₂ from air: 2NaOH + CO₂ → Na₂CO₃ + H₂O.
• Na₂CO₃ solution is alkaline by hydrolysis: CO₃²⁻ + H₂O ⇌ HCO₃⁻ + OH⁻.
• NaHCO₃ is only mildly alkaline, which is why it can be used as an antacid.

Solved example 10 — Lime water and CO₂

CO₂ is bubbled through lime water. Explain why it first turns milky and then clear again.

1. Ca(OH)₂ + CO₂ → CaCO₃↓ + H₂O — insoluble CaCO₃ makes it milky.
2. Excess CO₂: CaCO₃ + CO₂ + H₂O → Ca(HCO₃)₂, which is soluble.

Answer: Milky from CaCO₃; clear again as soluble calcium bicarbonate forms. (This is also how temporary hardness arises in nature.)

Check yourself

• C7. Why can anhydrous MgCl₂ not be made simply by heating MgCl₂·6H₂O?
• C8. Name the three main constituents of cement clinker.

Answers at the end.

Biological role

Four ions you depend on every second.

• Na⁺ mainly outside cells and K⁺ mainly inside; the sodium–potassium pump (driven by ATP) keeps these gradients for nerve impulses.
• Mg²⁺ is at the centre of chlorophyll and activates many enzymes that use ATP.
• Ca²⁺ builds bones and teeth (as hydroxyapatite) and is needed for muscle contraction and blood clotting.

Exam trap gallery

Where marks are lost in the s-block.

#	What students write	What is actually correct	Why
1	Li reacts most vigorously with water because its E° is most negative.	Li reacts most slowly; E° is thermodynamics, speed is kinetics.	High hydration enthalpy of Li⁺ explains E°, not the rate.
2	Li⁺ is the most mobile ion in water because it is smallest.	Li⁺ is the least mobile: it is the most hydrated.	Hydrated radius: Li⁺ largest.
3	Na burns in air to Na₂O.	Mainly Na₂O₂ (peroxide); K gives KO₂.	Larger cations stabilise larger anions.
4	Solubility of group 2 sulphates increases down the group.	It decreases: BaSO₄ is insoluble.	Hydration enthalpy falls faster than lattice enthalpy.
5	Solubility of group 2 hydroxides decreases down the group.	It increases: Ba(OH)₂ is soluble.	Small OH⁻: lattice enthalpy falls faster.
6	K₂CO₃ can be made by the Solvay process.	No: KHCO₃ is too soluble to precipitate.	The process depends on NaHCO₃ crystallising.
7	BeCl₂ is ionic like MgCl₂.	BeCl₂ is covalent: polymeric solid, dimer in vapour.	Tiny Be²⁺ is very polarising.
8	Heating gypsum strongly makes better plaster.	Above 393 K it gives dead burnt plaster that will not set.	All water of crystallisation is lost.
9	Thermal stability of alkali metal hydrides increases down the group.	It decreases: LiH is the most stable.	Lattice energy falls as the cation grows.
10	Heating MgCl₂·6H₂O gives anhydrous MgCl₂.	It hydrolyses to MgO (or Mg(OH)Cl) and HCl.	Mg²⁺ is small and strongly hydrated.
11	All group 1 nitrates give the nitrite on heating.	LiNO₃ gives Li₂O + NO₂ + O₂, like Mg(NO₃)₂.	Diagonal relationship of Li with Mg.
12	CaF₂ is soluble like CaCl₂.	CaF₂ is almost insoluble.	Very high lattice energy with the small F⁻ ion.

Memory hook: “Li-Na-K: oxide, peroxide, superoxide.” · “Big anion — down goes solubility; small anion — up it goes.” · “Small cation cracks carbonate” — less stable to heat. · “Blue dilute, bronze concentrated” — metals in liquid NH₃.

Practice MCQs

Levels: Easy, Moderate, Tough.

Q01. The alkali metal giving a crimson red flame is: (Easy)

• (a) Na
• (b) K
• (c) Li
• (d) Cs

Q02. Potassium burns in excess oxygen to give: (Easy)

• (a) K₂O
• (b) K₂O₂
• (c) KO₂
• (d) KOH

Q03. Which group 2 element gives no flame colour? (Easy)

• (a) Ca
• (b) Sr
• (c) Ba
• (d) Mg

Q04. The chemical formula of plaster of Paris is: (Easy)

• (a) CaSO₄·2H₂O
• (b) CaSO₄·½H₂O
• (c) CaSO₄
• (d) CaCO₃

Q05. Washing soda is: (Easy)

• (a) NaHCO₃
• (b) Na₂CO₃·10H₂O
• (c) NaOH
• (d) Na₂SO₄·10H₂O

Q06. The least mobile ion in aqueous solution is: (Moderate)

• (a) Li⁺
• (b) Na⁺
• (c) K⁺
• (d) Cs⁺

Q07. Which carbonate decomposes most easily on heating? (Moderate)

• (a) Na₂CO₃
• (b) K₂CO₃
• (c) Li₂CO₃
• (d) Cs₂CO₃

Q08. Least soluble in water: (Moderate)

• (a) MgSO₄
• (b) CaSO₄
• (c) SrSO₄
• (d) BaSO₄

Q09. Most basic hydroxide: (Moderate)

• (a) Be(OH)₂
• (b) Mg(OH)₂
• (c) Ca(OH)₂
• (d) Ba(OH)₂

Q10. Blue colour of alkali metals in liquid ammonia is due to: (Moderate)

• (a) metal ions
• (b) ammoniated electrons
• (c) NH₂⁻ ions
• (d) metal atoms

Q11. Diagonal relationship is shown by: (Moderate)

• (a) Li and Mg
• (b) Na and Mg
• (c) Be and Mg
• (d) Li and Be

Q12. The Solvay process cannot make K₂CO₃ because: (Moderate)

• (a) K is too reactive
• (b) KHCO₃ is too soluble
• (c) KCl is insoluble
• (d) NH₃ reacts with K

Q13. Be(OH)₂ dissolves in excess NaOH because it is: (Tough)

• (a) basic
• (b) acidic only
• (c) amphoteric
• (d) neutral

Q14. A concentrated solution of Na in liquid NH₃ is: (Tough)

• (a) blue, paramagnetic
• (b) bronze, diamagnetic
• (c) colourless, diamagnetic
• (d) blue, diamagnetic

Q15. Li has the most negative E° among alkali metals mainly because of: (Tough)

• (a) lowest ionisation enthalpy
• (b) highest hydration enthalpy of Li⁺
• (c) lowest sublimation enthalpy
• (d) largest atomic size

Q16. The alkali metal with the highest melting point is: (Easy)

• (a) Li
• (b) Na
• (c) K
• (d) Cs

Q17. Epsom salt is: (Easy)

• (a) MgSO₄·7H₂O
• (b) CaSO₄·2H₂O
• (c) Na₂SO₄·10H₂O
• (d) BaSO₄

Q18. Most stable alkali metal hydride: (Easy)

• (a) LiH
• (b) NaH
• (c) KH
• (d) CsH

Q19. NaNO₃ on heating gives: (Moderate)

• (a) Na₂O + NO₂ + O₂
• (b) NaNO₂ + O₂
• (c) Na + NO₂
• (d) Na₂O₂ + N₂

Q20. Mg₃N₂ with water gives: (Moderate)

• (a) Mg(OH)₂ + N₂
• (b) Mg(OH)₂ + NH₃
• (c) MgO + NO₂
• (d) Mg + HNO₃

Q21. BeH₂ is: (Moderate)

• (a) ionic
• (b) covalent and polymeric
• (c) a gas at room temperature
• (d) unknown

Q22. Gypsum is added to cement to: (Moderate)

• (a) speed up setting
• (b) slow down setting
• (c) increase strength only
• (d) change its colour

Q23. In the Castner–Kellner cell, the cathode is: (Moderate)

• (a) graphite
• (b) iron
• (c) mercury
• (d) platinum

Q24. The least soluble fluoride is: (Tough)

• (a) LiF
• (b) NaF
• (c) CaF₂
• (d) KF

Q25. Hydration enthalpy of Mg²⁺ is much larger than that of Na⁺ because: (Tough)

• (a) Mg²⁺ is larger
• (b) Mg²⁺ has higher charge and smaller size
• (c) Na⁺ is more electronegative
• (d) Mg is in group 2 only

Answer key with explanations

1. (c) Li — Li → crimson.
2. (c) KO₂ — Superoxide.
3. (d) Mg — Electrons held too tightly.
4. (b) CaSO₄·½H₂O — Hemihydrate.
5. (b) Na₂CO₃·10H₂O — Decahydrate of sodium carbonate.
6. (a) Li⁺ — Largest hydrated radius.
7. (c) Li₂CO₃ — Small polarising Li⁺.
8. (d) BaSO₄ — Solubility falls down the group.
9. (d) Ba(OH)₂ — Basic strength increases down.
10. (b) ammoniated electrons — Solvated electrons absorb red light.
11. (a) Li and Mg — Similar charge/size ratio.
12. (b) KHCO₃ is too soluble — No precipitate forms.
13. (c) amphoteric — Forms [Be(OH)₄]²⁻.
14. (b) bronze, diamagnetic — Electrons pair up.
15. (b) highest hydration enthalpy of Li⁺ — Hydration energy outweighs the higher ionisation enthalpy.
16. (a) Li — Strongest metallic bonding.
17. (a) MgSO₄·7H₂O — Magnesium sulphate heptahydrate.
18. (a) LiH — Highest lattice energy.
19. (b) NaNO₂ + O₂ — Group 1 nitrate (not Li).
20. (b) Mg(OH)₂ + NH₃ — Nitride hydrolysis.
21. (b) covalent and polymeric — Bridging Be–H–Be bonds.
22. (b) slow down setting — Retards setting.
23. (c) mercury — Forms sodium amalgam.
24. (c) CaF₂ — Very high lattice energy.
25. (b) Mg²⁺ has higher charge and smaller size — Charge density.

Check-yourself answers

• C1. Lithium (Li₂O).
• C2. The ammoniated electrons pair up at high concentration, so there are no unpaired electrons.
• C3. LiH > NaH > KH.
• C4. NaNO₃ → NaNO₂ + O₂; LiNO₃ → Li₂O + NO₂ + O₂.
• C5. BeCO₃ < CaCO₃ < BaCO₃.
• C6. Ba(OH)₂.
• C7. The hydrated salt hydrolyses on heating, giving MgO (or Mg(OH)Cl) and HCl; it must be heated in a stream of HCl gas.
• C8. Dicalcium silicate, tricalcium silicate and tricalcium aluminate.

One-page revision

Read this page the night before the exam.

Trends

• ns¹ / ns²; size ↑, IE ↓ down
• Hydration: Li⁺ largest → least mobile
• Mobility Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺
• m.p. falls down group 1

Flames and oxygen

• Li crimson, Na yellow, K lilac, Cs blue
• Ca brick red, Sr crimson, Ba apple green
• Be, Mg: no colour
• Li₂O, Na₂O₂, KO₂

Water and ammonia

• Li most negative E° but slowest with water
• 2M + 2H₂O → 2MOH + H₂
• Liquid NH₃: blue, paramagnetic (dilute)
• Concentrated: bronze, diamagnetic

Li and Be

• Li–Mg: nitride, carbonate and nitrate decompose
• Be–Al: amphoteric, covalent chlorides
• BeCl₂ polymer (solid), dimer (vapour)
• Be max covalency 4

Solubility and stability

• Sulphates: solubility ↓ down group 2
• Hydroxides: solubility, basicity ↑
• Carbonates: stability ↑ down
• Li₂CO₃ decomposes; Na₂CO₃ stable

Compounds

• NaOH: Castner–Kellner
• Solvay: NaHCO₃ precipitates; not K₂CO₃
• Gypsum → PoP at 393 K; above: dead burnt
• 4KO₂ + 2CO₂ → 2K₂CO₃ + 3O₂