7.1
Occurrence & Physical Properties of Group 2 Elements
Occurrence (Alkaline Earth Metals)
Group 2 elements (Be, Mg, Ca, Sr, Ba, Ra) are the alkaline earth metals. Like Group 1, they are too reactive to occur as free elements in nature — found as ionic compounds in minerals:
- Beryllium: beryl (Be₃Al₂Si⁶O)
- Magnesium: magnesite (MgCO₃), dolomite (MgCO₃·CaCO₃), seawater
- Calcium: calcite/limestone/chalk (CaCO₃), gypsum (CaSO₄·2H₂O), fluorite (CaF₂)
- Strontium: strontianite (SrCO₃), celestine (SrSO₄)
- Barium: barite (BaSO₄), witherite (BaCO₃)
- Radium: radioactive; trace amounts in uranium ores
| Element | Z | Config | M.p. (°C) | Density (g/cm³) | Flame colour |
|---|---|---|---|---|---|
| Beryllium (Be) | 4 | [He]2s² | 1287 | 1.85 | No characteristic colour |
| Magnesium (Mg) | 12 | [Ne]3s² | 650 | 1.74 | Brilliant white |
| Calcium (Ca) | 20 | [Ar]4s² | 842 | 1.55 | Brick red/orange-red |
| Strontium (Sr) | 38 | [Kr]5s² | 777 | 2.64 | Crimson red |
| Barium (Ba) | 56 | [Xe]6s² | 727 | 3.51 | Apple green |
| Radium (Ra) | 88 | [Rn]7s² | 700 | 5.50 | Crimson (radioactive) |
Trends in Physical Properties Down Group 2
- Melting point: No simple trend (Be highest; irregular due to changing structure/bonding). Generally lower than expected from metallic bond argument alone — Be anomalously high.
- Density: increases — molar mass increases faster than atomic volume.
- Atomic/ionic radius: increases — new shell added each period.
- First ionisation energy: decreases — outer electron in higher shell, more shielded.
- Reactivity: increases — easier to lose both outer s electrons.
- Electronegativity: decreases
Comparison with Group 1: Group 2 metals have higher melting points, harder, and higher densities than corresponding Group 1 metals because they contribute 2 delocalised electrons per atom → stronger metallic bonding.
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Section 7.1 — Group 2 Properties
10 QuestionsQ1 of 10
Down Group 2, reactivity with water:
Be: no reaction with water. Mg: reacts slowly with hot water/steam. Ca: reacts with cold water (Ca(OH)₂). Sr, Ba: react vigorously. Reactivity increases down group.
Q2 of 10
The equation for Ca reacting with cold water:
Ca + 2H₂O → Ca(OH)₂ + H₂. Ca(OH)₂ is slightly soluble (limewater). The solution turns milky with CO₂.
Q3 of 10
Group 2 oxides react with water to give:
MO + H₂O → M(OH)₂. Example: CaO + H₂O → Ca(OH)₂ (exothermic, quicklime slaking). Products are alkaline.
Q4 of 10
The trend in solubility of Group 2 hydroxides going DOWN the group:
Hydroxide solubility increases: Be(OH)₂ (insoluble) → Mg(OH)₂ (sparingly) → Ca(OH)₂ (slightly) → Ba(OH)₂ (soluble). Lattice energy decreases faster than hydration energy.
Q5 of 10
The trend in solubility of Group 2 sulfates going DOWN the group:
Sulfate solubility DECREASES: MgSO₄ (soluble) → CaSO₄ (slightly) → BaSO₄ (insoluble). Opposite trend to hydroxides. BaSO₄ insolubility used in Ba²⁺ test.
Q6 of 10
Mg(OH)₂ is used in:
Mg(OH)₂ (milk of magnesia): antacid, reacts with HCl in stomach. Also used as laxative. Ca(OH)₂ used in agriculture (neutralise acid soils).
Q7 of 10
Thermal stability of Group 2 carbonates DOWN the group:
Thermal stability increases down Group 2. Larger M²⁺ → lower charge density → less polarisation of CO₃²⁻ → more stable. MgCO₃ decomposes ~300°C; BaCO₃ ~1400°C.
Q8 of 10
BaSO₄ is used as a contrast agent in X-ray diagnosis because:
BaSO₄: highly insoluble (Ksp very small) → not absorbed into bloodstream → non-toxic. Dense (absorbs X-rays) → shows stomach/intestine outline. 'Barium meal'.
Q9 of 10
Ca(OH)₂ (limewater) is used to test for CO₂. The reaction is:
CO₂ + Ca(OH)₂ → CaCO₃↓ (white ppt, turns limewater milky). Excess CO₂: CaCO₃ + CO₂ + H₂O → Ca(HCO₃)₂ (soluble, clears again).
Q10 of 10
Magnesium burns in CO₂ to give:
2Mg + CO₂ → 2MgO + C. Mg is such a strong reducing agent it reduces CO₂ → white MgO + black carbon. This is why Mg fires cannot be extinguished with CO₂ extinguishers.
7.2
Reactivity of Group 2 Elements
Reactions with Oxygen
General: 2M + O₂ → 2MO (metal oxide, normal oxide for all Group 2)
2Mg + O₂ → 2MgO (bright white flame)
2Ca + O₂ → 2CaO
2Ba + O₂ → 2BaO (Ba may also form BaO₂ peroxide in excess O₂)
All Group 2 metals burn brilliantly in oxygen. Mg burns with a very bright white flame (used in flares). BaO₂ can form with excess oxygen for Ba.
Reactions with Water
General: M + 2H₂O → M(OH)₂ + H₂
Be: No reaction with water (even steam) at room temp — protected by oxide layer
Mg: Very slow reaction with cold water; reacts with steam:
Mg + H₂O(g) → MgO + H₂
Ca: Reacts slowly with cold water:
Ca + 2H₂O → Ca(OH)₂ + H₂ (milky suspension)
Sr: Reacts more vigorously than Ca
Ba: Most vigorous of the common Group 2 metals with water
Reactivity increases Ca < Sr < Ba. The solution formed is alkaline (M(OH)₂ is a base).
Reactions with Dilute Acids
All Group 2 metals react readily with dilute acids:
Mg + 2HCl → MgCl₂ + H₂ (vigorous)
Ca + 2HCl → CaCl₂ + H₂ (vigorous)
Mg + H₂SO₄ → MgSO₄ + H₂ (initially fast, then slows —
MgSO₄ forms a protective layer)
Note: Ca + H₂SO₄ → CaSO₄ (insoluble layer) + H₂. CaSO₄ is sparingly soluble and coats Ca surface → reaction slows.
Reactions with Halogens
M + X₂ → MX₂ (metal dihalide)
Mg + Cl₂ → MgCl₂
Ca + Br₂ → CaBr₂
Ba + F₂ → BaF₂
7.3
Properties of Group 2 Compounds
Oxides
All Group 2 oxides are basic (ionic, with O²⁻). They react with water to form hydroxides, and with acids to form salts:
MgO + H₂O → Mg(OH)₂ (slightly exothermic; Mg(OH)₂ slightly soluble)
CaO + H₂O → Ca(OH)₂ (very exothermic — "slaking" of quicklime; produces
slaked lime)
MgO + 2HCl → MgCl₂ + H₂O
CaO + H₂SO₄ → CaSO₄ + H₂O
BeO is amphoteric (reacts with both acids AND bases) — anomalous (see 7.4).
Hydroxides
Group 2 hydroxides M(OH)₂ are basic. Solubility and basicity increase down the group:
| Hydroxide | Solubility | Notes |
|---|---|---|
| Be(OH)₂ | Insoluble; amphoteric | Reacts with both HCl and NaOH; anomalous |
| Mg(OH)₂ | Sparingly soluble | Weakly alkaline suspension; antacid (milk of magnesia) |
| Ca(OH)₂ | Slightly soluble (limewater) | pH ~12 in solution; used to test for CO₂ |
| Sr(OH)₂ | Moderately soluble | Stronger base than Ca(OH)₂ |
| Ba(OH)₂ | Soluble | Most soluble and most alkaline Group 2 hydroxide |
The increase in solubility down the group is due to decreasing lattice energy (larger M²⁺) outpacing decreasing hydration energy → solubility increases.
Carbonates
All Group 2 carbonates are thermally unstable — they decompose on heating to give metal oxide + CO₂:
MCO₃ → MO + CO₂
MgCO₃ → MgO + CO₂ (decomposes easily)
CaCO₃ → CaO + CO₂ (limestone → quicklime; needs ~900°C)
SrCO₃ → SrO + CO₂
BaCO₃ → BaO + CO₂ (requires higher temperature)
Thermal stability increases down the group: larger M²⁺ ions have lower charge density → less polarisation of CO²⁻ → harder to break → higher decomposition temperature required.
| Carbonate | Approx. decomp. temp (°C) |
|---|---|
| MgCO₃ | ~540 |
| CaCO₃ | ~900 |
| SrCO₃ | ~1290 |
| BaCO₃ | ~1360 |
Sulfates
Solubility of Group 2 sulfates decreases down the group — the opposite trend to hydroxides:
| Sulfate | Solubility | Notes |
|---|---|---|
| BeSO₄ | Soluble | — |
| MgSO₄ | Soluble | Epsom salt (MgSO₄·7H₂O) |
| CaSO₄ | Slightly soluble (sparingly) | Gypsum; causes "permanent" hard water |
| SrSO₄ | Insoluble | — |
| BaSO₄ | Virtually insoluble | Used as barium meal (X-ray contrast); confirms SO²⁻ |
The decrease in solubility is because the lattice energy decreases only slightly (the SO²⁻ anion is large), while hydration energy decreases significantly as M²⁺ gets larger → less energy released on dissolving → lower solubility.
Nitrates
All Group 2 nitrates decompose on heating to give metal oxide + NO₂ + O₂:
2M(NO₃)₂ → 2MO + 4NO₂ + O₂
2Mg(NO₃)₂ → 2MgO + 4NO₂ + O₂
2Ca(NO₃)₂ → 2CaO + 4NO₂ + O₂
Unlike Group 1 nitrates (which give nitrite + O₂), Group 2 nitrates decompose more completely to give oxide + NO₂ (brown gas) + O₂. This is because M²⁺ cations have higher charge density than M⁺ → more strongly polarise NO⁻ → break N–O bonds more fully. Thermal stability of nitrates increases down Group 2 (same reasoning as carbonates).
Summary: Solubility Trends
| Compound | Solubility trend down Group 2 |
|---|---|
| Hydroxides M(OH)₂ | Increases: Be(OH)₂ insoluble → Ba(OH)₂ soluble |
| Sulfates MSO₄ | Decreases: MgSO₄ soluble → BaSO₄ insoluble |
| Carbonates MCO₃ | Decreases: MgCO₃ slightly soluble → BaCO₃ insoluble |
| Fluorides MF₂ | Decreases: MgF₂ slightly soluble → BaF₂ more soluble (complex trend) |
7.4
Anomalous Properties of Beryllium
Why Beryllium is Anomalous
Be is the first member of Group 2 and has a very small ionic radius (Be²⁺: 31 pm) and high charge density. This leads to properties much more similar to aluminium (Al) than to Mg — a diagonal relationship (Be and Al are diagonal neighbours in the periodic table).
| Property | Typical Group 2 behaviour | Be (anomalous) | Resembles Al |
|---|---|---|---|
| Oxide character | Basic (ionic) | BeO is amphoteric (reacts with both acids and bases) | Al₂O₃ is amphoteric |
| Hydroxide character | Basic | Be(OH)₂ is amphoteric | Al(OH)₃ is amphoteric |
| Bonding in compounds | Ionic | BeCl₂ is covalent (high charge density polarises Cl⁻) | AlCl₃ is covalent |
| Reaction with NaOH | No reaction | Be + 2NaOH + 2H₂O → Na₂[Be(OH)₄] + H₂ | Al reacts with NaOH |
| Chloride structure | Ionic MX₂ | BeCl₂ is a polymer chain (linear); Lewis acid | AlCl₃ is dimeric/covalent |
| Reaction with water | React to give M(OH)₂ + H₂ | Be does not react with water | Al does not react with water readily (oxide layer) |
| Complexes | Simple ionic compounds | Be forms complex ions [Be(OH)₄]²⁻ | Al forms [Al(OH)₄]⁻ |
Amphoteric reactions of Be(OH)₂:
With acid: Be(OH)₂ + 2HCl → BeCl₂ + 2H₂O (acts as base)
With base: Be(OH)₂ + 2NaOH → Na₂[Be(OH)₄] (acts as acid)
Beryllium toxicityBeryllium and its compounds are extremely toxic. Inhalation of Be dust causes berylliosis (a chronic lung disease). BeCl₂ and other Be compounds must be handled with great care. This limits its industrial use despite its excellent mechanical properties.
7.5
Identification Test for Ba²⁺ Ions in Aqueous Solution
Test for Barium Ions
Barium ions are identified by two key tests:
- Flame test: Ba²⁺ gives a characteristic apple green flame.
- Sulfate precipitation test: Add dilute H₂SO₄ (or dilute HCl then BaCl₂ solution, or dilute HNO₃ then Ba(NO₃)₂):
Ba²⁺(aq) + SO²⁻(aq) → BaSO(s) (dense white precipitate, insoluble in dilute HCl).
Test for Ba²⁺:
Add dilute HCl to acidify, then add dilute H₂SO₄.
Observation: dense white precipitate of BaSO₄ forms.
BaSO₄ is insoluble in dilute HCl — this distinguishes it from BaSO₃
(which would dissolve in HCl).
Ionic equation: Ba²⁺(aq) + SO²⁻(aq) → BaSO(s)
BaSO₄ as a barium mealBaSO₄ is used as a contrast agent in X-ray examination of the gastrointestinal tract ("barium meal" or "barium enema"). It is opaque to X-rays (shows white on the image) and is safe because it is essentially insoluble in water and digestive fluids — it cannot be absorbed into the bloodstream (unlike soluble Ba²⁺ which is highly toxic).
Additional Group 2 Ion Tests
| Ion | Test | Observation |
|---|---|---|
| Ca²⁺ | Add Na₂CO₃(aq) or (NH₄)₂CO₃(aq) | White precipitate of CaCO₃ (soluble in dilute HCl) |
| Ca²⁺ | Add dilute H₂SO₄ | White precipitate of CaSO₄ (sparingly soluble) |
| Ba²⁺ | Add dilute H₂SO₄ | Dense white precipitate of BaSO₄ (insoluble in HCl) |
| Mg²⁺ | Add NaOH(aq) | White gelatinous precipitate of Mg(OH)₂ (insoluble in excess NaOH) |
| Be²⁺ | Add NaOH(aq) | White precipitate initially; dissolves in excess NaOH (amphoteric) |
7.6
Uses of Group 2 Elements and Their Compounds
| Substance | Uses |
|---|---|
| Beryllium (Be) | Aerospace alloys (Be-Cu very hard & strong); X-ray windows (low atomic number → transparent to X-rays); nuclear reactor components; gyroscopes |
| Magnesium (Mg) | Lightweight alloys (Al-Mg for aircraft, cars); Grignard reagents in organic synthesis; reducing agent in metallurgy (Ti, Zr extraction); flares and fireworks (burns brightly) |
| MgO | Refractory lining for furnaces (very high m.p.); antacid (milk of magnesia, Mg(OH)₂); cement; insulation |
| Mg(OH)₂ | Antacid ("milk of magnesia" — neutralises excess stomach acid); flame retardant |
| MgSO₄·7H₂O | Epsom salts: laxative, bath salts, muscle relaxant; agriculture (Mg fertiliser) |
| Calcium (Ca) | Reducing agent in metallurgy (production of rare metals like U, Th); calcium supplements |
| CaCO₃ (limestone) | Building material; manufacture of CaO (quicklime); glass making; iron production (flux); paper; soil neutralisation |
| CaO (quicklime) | Steel making (removes silica impurities as slag); manufacture of Ca(OH)₂; soil treatment; drying agent |
| Ca(OH)₂ (slaked lime) | Cement and mortar; soil treatment (reduces acidity, raises pH for agriculture); water treatment (softening); flue gas desulfurisation |
| CaSO₄·2H₂O (gypsum) | Plaster of Paris (CaSO·½H₂O) — casts and moulds; blackboard chalk; building material; soil conditioner |
| CaF₂ (fluorite/fluorspar) | Manufacture of HF; optical components; steel making (flux) |
| BaSO₄ | Barium meal X-ray contrast agent; white pigment (lithopone); drilling mud in oil wells |
| BaCO₃ | Rat poison (toxic Ba compound); glass and ceramics; CRT screens |
| Sr compounds | Strontium carbonate/nitrate in red fireworks and signal flares (crimson flame); SrTiO₃ as diamond substitute |
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Section 7.2 — Uses of Group 2 Compounds
10 QuestionsQ1 of 10
CaO (quicklime) is made by:
CaCO₃ → CaO + CO₂ (calcination). Industrial kilns. CaO used in cement, steel making, neutralising acidic soils.
Q2 of 10
Ca(OH)₂ is used in agriculture to:
Ca(OH)₂ (slaked lime, agricultural lime): neutralises acidic soil (increases pH), improves soil structure. Cheaper and more available than other alkalis.
Q3 of 10
MgO is used as a refractory material because:
MgO: ionic compound with 2+/2− charges → enormous lattice energy → extremely high m.p. → stable at furnace temperatures → used to line furnaces, kilns.
Q4 of 10
CaCO₃ (limestone) is used in:
CaCO₃: glass (with SiO₂ + Na₂CO₃), cement (with clay, heated → calcium silicates), neutralising acid in lakes (liming), blast furnace flux.
Q5 of 10
Which Group 2 compound is used in fireworks to give a red colour?
Sr²⁺: red/crimson flame. Ba²⁺: green. Ca²⁺: orange-red. Mg: white. Different emission spectra from different metals used for firework colours.
Q6 of 10
Ba²⁺ ions are toxic. Why is BaSO₄ safe to use medicinally?
Ksp(BaSO₄) = 1×10⁻¹⁰ mol²/dm⁶ — extremely insoluble. Free Ba²⁺ concentration negligible → cannot be absorbed → non-toxic. Soluble Ba²⁺ compounds (BaCl₂) ARE toxic.
Q7 of 10
Dolomite (MgCO₃·CaCO₃) is heated to give:
Dolomite calcination: MgCO₃·CaCO₃ → MgO + CaO + 2CO₂. Dolime (MgO+CaO) used in steel making.
Q8 of 10
Which Group 2 hydroxide is used in making cement?
Ca(OH)₂ (slaked lime) + SiO₂/Al₂O₃ → calcium silicates/aluminates (cement minerals). Cement hardens when mixed with water.
Q9 of 10
CaSO₄·½H₂O (plaster of Paris) sets hard because:
Plaster of Paris + water → gypsum. Slight expansion on setting → fills moulds precisely. Used in casts, moulding, building.
Q10 of 10
Which Group 2 element is essential for plant chlorophyll?
Magnesium (Mg) is the central atom in the chlorophyll molecule (porphyrin ring). Mg²⁺ deficiency → yellowing of leaves (chlorosis). Ca²⁺ also essential (cell wall structure).
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Exercises
- State and explain the trend in thermal stability of Group 2 carbonates (MgCO₃ to BaCO₃). Write equations for the decomposition of MgCO₃ and CaCO₃.
Thermal stability increases down the group: MgCO₃ decomposes most easily (~540°C); BaCO₃ requires the highest temperature (~1360°C).
Reason: Smaller M²⁺ cations (Mg²⁺) have higher charge density → strongly polarise the CO²⁻ ion → distort electron density, weakening C–O bonds → easier decomposition. Larger cations (Ba²⁺) have lower charge density → polarise CO²⁻ less → more stable carbonate → higher temperature needed.
MgCO₃ → MgO + CO₂
CaCO₃ → CaO + CO₂ - Explain why the solubility of Group 2 hydroxides increases down the group but the solubility of Group 2 sulfates decreases. Relate both trends to lattice energy and hydration energy.
Hydroxides (increasing solubility): As M²⁺ gets larger, lattice energy decreases significantly (because lattice energy ∝ 1/r) since OH⁻ is small → the lattice weakens noticeably. Hydration energy also decreases but less steeply. Net result: the lattice energy decrease dominates → dissolving becomes more energetically favourable → solubility increases (Mg(OH)₂ sparingly soluble; Ba(OH)₂ soluble).
Sulfates (decreasing solubility): SO²⁻ is a large anion → lattice energy changes only slightly as M²⁺ increases (because the lattice energy is already dominated by the large SO²⁻). However, hydration energy of M²⁺ decreases significantly as radius increases. Net result: the decrease in hydration energy dominates → dissolving becomes less favourable → solubility decreases (MgSO₄ soluble; BaSO₄ insoluble). - State the reagent and observation for the identification of Ba²⁺ ions. Write an ionic equation for the reaction.
Add dilute HCl (to acidify and remove interfering ions), then add dilute H₂SO₄ (or dilute HNO₃ then Na₂SO₄ solution).
Observation: dense white precipitate of BaSO₄ forms. The precipitate is insoluble in dilute HCl (this distinguishes it from other white precipitates like BaCO₃ which dissolves in HCl).
Ionic equation: Ba²⁺(aq) + SO²⁻(aq) → BaSO(s) - Describe THREE ways in which beryllium behaves anomalously in Group 2, and explain what element it resembles (diagonal relationship).
Be resembles aluminium (Al) — its diagonal neighbour (Period 3, Group 13). Both have small ionic radii and high charge density.
(1) Amphoteric oxide and hydroxide: BeO and Be(OH)₂ react with both acids (BeO + 2HCl → BeCl₂ + H₂O) and bases (Be(OH)₂ + 2NaOH → Na₂[Be(OH)₄]). Other Group 2 oxides/hydroxides are purely basic. Al₂O₃ and Al(OH)₃ are also amphoteric.
(2) Covalent bonding in halides: BeCl₂ is covalent (polymeric chain), not ionic. Other Group 2 chlorides are ionic. AlCl₃ is also covalent.
(3) No reaction with water: Be does not react with water or steam. Other Group 2 metals (Ca, Sr, Ba) react with water. Al also does not react with water readily (oxide layer). (Also acceptable: Be reacts with NaOH; other Group 2 metals don't.) - Write balanced equations for: (a) calcium reacting with water; (b) CaO reacting with water; (c) Ca(OH)₂ reacting with CO₂; (d) CaCO₃ reacting with dilute HCl.
(a) Ca + 2H₂O → Ca(OH)₂ + H₂
(b) CaO + H₂O → Ca(OH)₂ (exothermic; slaking of quicklime)
(c) Ca(OH)₂ + CO₂ → CaCO₃ + H₂O (limewater turns milky)
(d) CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂ - Explain why Ca(OH)₂ is used to treat acidic soil, while CaCO₃ is also used. Which is more effective per mole and why?
Acidic soil has a low pH that inhibits plant growth. Both Ca(OH)₂ and CaCO₃ are bases that neutralise acid:
Ca(OH)₂ + 2H⁺ → Ca²⁺ + 2H₂O
CaCO₃ + 2H⁺ → Ca²⁺ + H₂O + CO₂
Ca(OH)₂ is more effective per mole because it provides 2 OH⁻ ions per formula unit → neutralises 2 H⁺ ions. It is also more soluble than CaCO₃. However, CaCO₃ is cheaper, abundant (limestone), and safer to handle (Ca(OH)₂ is more caustic). CaCO₃ also releases CO₂ which acidifies slightly, but the net effect is still pH increase. In practice, powdered limestone (CaCO₃) is widely used because of cost and availability.
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Multiple Choice Quiz — 25 Questions
Unit 7 Quiz
Select one answer per questionQ1
The thermal stability of Group 2 carbonates increases down the group because:
Smaller M²⁺ (e.g. Mg²⁺) has high charge density → strongly polarises CO²⁻ → decomposes easily. Larger M²⁺ (e.g. Ba²⁺) has lower charge density → polarises CO²⁻ less → more stable → higher decomposition temperature needed.
Q2
The solubility of Group 2 sulfates:
Sulfate solubility decreases: MgSO₄ (soluble) → BaSO₄ (insoluble). Large SO²⁻ means lattice energy changes little, but hydration energy of M²⁺ falls steeply → less soluble going down.
Q3
The reagent used to test for Ba²⁺ ions is:
Adding dilute H₂SO₄ to a solution containing Ba²⁺ gives a dense white precipitate of BaSO₄: Ba²⁺ + SO²⁻ → BaSO(s). BaSO is insoluble in dilute HCl.
Q4
Beryllium is anomalous in Group 2 because it:
Be is anomalous: BeO/Be(OH)₂ are amphoteric (react with both acids and bases); BeCl₂ is covalent; Be does not react with water; Be resembles Al (diagonal relationship).
Q5
The flame colour of barium is:
Ba²⁺ gives an apple green flame. Ca = brick red; Sr = crimson; Mg = brilliant white (not characteristic for flame test purposes).
Q6
When CaO reacts with water, the product is:
CaO + H₂O → Ca(OH)₂ (slaked lime). This highly exothermic reaction is "slaking" of quicklime (CaO).
Q7
Why is BaSO₄ used as a contrast agent in X-ray examination?
BaSO is virtually insoluble → cannot be absorbed into blood (soluble Ba²⁺ is toxic). Its high atomic number makes it opaque to X-rays → shows up white on X-ray images → outlines the GI tract.
Q8
Which Group 2 hydroxide is most soluble?
Solubility of Group 2 hydroxides increases down the group. Ba(OH)₂ is the most soluble and most strongly alkaline.
Q9
Group 2 nitrates decompose to give:
2M(NO₃)₂ → 2MO + 4NO₂ + O₂. Group 2 nitrates (unlike Group 1 except Li) decompose all the way to oxide + brown NO₂ gas + O₂.
Q10
Which element does beryllium most resemble due to the diagonal relationship?
Be (Period 2, Group 2) has a diagonal relationship with Al (Period 3, Group 13). Both have small, highly polarising cations; both have amphoteric oxides/hydroxides; both form covalent chlorides.
Q11
Limewater (Ca(OH)₂) turns milky when CO₂ is bubbled through because:
Ca(OH)₂ + CO₂ → CaCO₃(s) + H₂O. Insoluble white CaCO₃ precipitates → milky appearance. With excess CO₂: CaCO₃ + CO₂ + H₂O → Ca(HCO₃)₂ (soluble) → clears.
Q12
Magnesium does not react significantly with cold water but reacts with steam because:
Mg has a thin protective MgO/Mg(OH)₂ surface layer that slows cold water reaction. High temperature of steam overcomes this barrier. Mg + H₂O(g) → MgO + H₂.
Q13
Ca(OH)₂ is added to acidic soil to:
Ca(OH)₂ is a base. Acidic soil has low pH → inhibits plant growth. Ca(OH)₂ neutralises excess H⁺: Ca(OH)₂ + 2H⁺ → Ca²⁺ + 2H₂O → raises pH to optimum for plants.
Q14
Which of these Group 2 compounds is used as an antacid?
Mg(OH)₂ (milk of magnesia) is used as an antacid. It neutralises excess HCl in the stomach: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O. CaCO₃ is also used in some antacids (e.g. Tums).
Q15
The reactivity of Group 2 metals increases down the group because:
Both IE₁ and IE₂ decrease down Group 2 (outer 2 electrons are in higher shells, more shielded) → both electrons are easier to lose → greater reactivity going down the group.
Q16
CaSO₄ is only slightly soluble in water. This explains why Ca reacts with H₂SO₄ only initially because:
Ca + H₂SO → CaSO₄(s) + H₂. CaSO₄ is sparingly soluble → deposits on Ca surface as a protective layer → blocks access of H₂SO₄ to Ca → reaction effectively stops (passivation).
Q17
Which carbonate requires the highest temperature to decompose?
BaCO₃ (~1360°C) has the highest decomposition temperature. Ba²⁺ is the largest cation → lowest charge density → least polarisation of CO²⁻ → most stable carbonate.
Q18
When excess CO₂ is bubbled through limewater (initially turned milky), the solution clears because:
Excess CO₂ converts insoluble CaCO to soluble calcium hydrogen carbonate Ca(HCO)₂ → precipitate dissolves → solution clears. CaCO₃ + CO₂ + H₂O → Ca(HCO₃)₂(aq). This is reversible (heating reverses it → CaCO precipitates again — basis of stalactite/stalagmite formation).
Q19
The ionic equation for the reaction of Be(OH)₂ with NaOH is best described as:
Be(OH)₂ is amphoteric. With NaOH (a base): Be(OH)₂ + 2OH⁻ → [Be(OH)₄]²⁻. Be(OH)₂ accepts OH⁻ ions — it acts as a Lewis acid (electron pair acceptor), behaving like an acid in this context.
Q20
Epsom salt (MgSO₄·7H₂O) is used as:
MgSO·7H₂O (Epsom salt): used as a laxative, bath salts (muscle relaxant), and as a magnesium fertiliser for plants (Mg is essential for chlorophyll).
Q21
CaO is called "quicklime" and is produced by:
CaCO₃ → CaO + CO₂ at ~900°C (industrial lime kiln). This is the thermal decomposition of limestone to produce quicklime (CaO), a major industrial process.
Q22
Group 2 metals have higher melting points than Group 1 metals in the same period because:
Group 2: ns² → 2 delocalised electrons per atom (vs 1 for Group 1). Greater electron density + higher M²⁺ charge + smaller ionic radius → stronger metallic bonding → higher m.p.
Q23
Which test would distinguish Ca²⁺ from Ba²⁺ in solution?
Both Ca²⁺ and Ba²⁺ give white precipitates with H₂SO₄. The distinction: BaSO₄ is insoluble in dilute HCl; CaSO₄ is sparingly soluble and may dissolve/not precipitate cleanly. The flame test also helps: Ca = brick red; Ba = apple green.
Q24
In the Solvay (ammonia-soda) process, Ca(OH)₂ is used to:
In the Solvay process, Ca(OH)₂ reacts with NH₄Cl (a by-product) to regenerate NH for recycling: Ca(OH)₂ + 2NH₄Cl → CaCl₂ + 2NH₃ + 2H₂O. This makes the process economical.
Q25
Magnesium is used in flares and fireworks because it:
2Mg + O₂ → 2MgO. Mg burns with an intense white light (temperature ~3100°C) that is very bright. The white MgO smoke is also visible. Used in flares, incendiary devices, and flash photography.
📝
Unit Test — 50 Marks
Section A — Short Answer
20 marksQ1 [4 marks]
Explain the trend in thermal stability of Group 2 carbonates from MgCO₃ to BaCO₃. Write equations for the decomposition of CaCO₃ and BaCO₃ and compare the temperatures required. [4]
Thermal stability increases from MgCO₃ to BaCO₃. [1]
Reason: The ability of M²⁺ to polarise (distort) the CO²⁻ ion weakens down the group. Small cations (Mg²⁺) have high charge density → strongly polarise CO²⁻ → distort electron distribution in C–O bonds → CO₂ released easily at low temperature. Large cations (Ba²⁺) have low charge density → barely polarise CO²⁻ → more energy (higher temperature) needed to break the carbonate. [1]
CaCO₃ → CaO + CO₂ (decomposes at ~900°C — used in lime kilns) [1]
BaCO₃ → BaO + CO₂ (requires ~1360°C — much more stable; cannot be decomposed with a Bunsen burner) [1]
Reason: The ability of M²⁺ to polarise (distort) the CO²⁻ ion weakens down the group. Small cations (Mg²⁺) have high charge density → strongly polarise CO²⁻ → distort electron distribution in C–O bonds → CO₂ released easily at low temperature. Large cations (Ba²⁺) have low charge density → barely polarise CO²⁻ → more energy (higher temperature) needed to break the carbonate. [1]
CaCO₃ → CaO + CO₂ (decomposes at ~900°C — used in lime kilns) [1]
BaCO₃ → BaO + CO₂ (requires ~1360°C — much more stable; cannot be decomposed with a Bunsen burner) [1]
Q2 [4 marks]
Describe the test for Ba²⁺ ions in solution. Write the ionic equation, state the observation, and explain why BaSO₄ is safe to use as a barium meal whereas BaCl₂ would be dangerous. [4]
Test: Acidify with dilute HCl (removes interfering CO²⁻ or SO²⁻ ions), then add dilute H₂SO₄ (or Na₂SO₄(aq)). [1]
Observation: Dense white precipitate forms immediately. The precipitate does not dissolve in dilute HCl. [1]
Ionic equation: Ba²⁺(aq) + SO²⁻(aq) → BaSO(s) [1]
BaSO₄ is safe because it is virtually insoluble in water and digestive fluids → cannot be absorbed through the gut wall into the bloodstream → does not reach toxic concentrations in the body. BaCl₂ is soluble → Ba²⁺ ions would be absorbed into the blood → highly toxic (interferes with K⁺ channels in cardiac muscle → cardiac arrest at sufficient doses). [1]
Observation: Dense white precipitate forms immediately. The precipitate does not dissolve in dilute HCl. [1]
Ionic equation: Ba²⁺(aq) + SO²⁻(aq) → BaSO(s) [1]
BaSO₄ is safe because it is virtually insoluble in water and digestive fluids → cannot be absorbed through the gut wall into the bloodstream → does not reach toxic concentrations in the body. BaCl₂ is soluble → Ba²⁺ ions would be absorbed into the blood → highly toxic (interferes with K⁺ channels in cardiac muscle → cardiac arrest at sufficient doses). [1]
Q3 [4 marks]
Explain the opposite solubility trends for Group 2 hydroxides and Group 2 sulfates as you go down the group. [4]
For any compound to dissolve, hydration energy (released when ions are surrounded by water) must compensate for lattice energy (required to separate ions). Going down Group 2, M²⁺ gets larger → hydration energy decreases (less attraction for water) and lattice energy decreases (weaker attraction between ions).
Hydroxides (increasing solubility) [2]: OH⁻ is a small anion. As M²⁺ increases in size, lattice energy falls steeply (large M²⁺ → more widely spaced ions). Hydration energy falls less sharply. Net result: more energy available after dissolving → solubility increases. Ba(OH)₂ is soluble; Mg(OH)₂ is sparingly soluble.
Sulfates (decreasing solubility) [2]: SO²⁻ is a large anion. Lattice energy doesn’t change much down the group (because the large SO²⁻ dominates the lattice size). However, hydration energy of M²⁺ falls significantly as it gets larger. Net result: insufficient hydration energy to compensate → becomes harder to dissolve → solubility decreases. BaSO₄ insoluble; MgSO₄ soluble.
Hydroxides (increasing solubility) [2]: OH⁻ is a small anion. As M²⁺ increases in size, lattice energy falls steeply (large M²⁺ → more widely spaced ions). Hydration energy falls less sharply. Net result: more energy available after dissolving → solubility increases. Ba(OH)₂ is soluble; Mg(OH)₂ is sparingly soluble.
Sulfates (decreasing solubility) [2]: SO²⁻ is a large anion. Lattice energy doesn’t change much down the group (because the large SO²⁻ dominates the lattice size). However, hydration energy of M²⁺ falls significantly as it gets larger. Net result: insufficient hydration energy to compensate → becomes harder to dissolve → solubility decreases. BaSO₄ insoluble; MgSO₄ soluble.
Q4 [4 marks]
Describe the anomalous properties of beryllium with reference to the diagonal relationship with aluminium. Give THREE specific chemical comparisons between Be and Al. [4]
Be (Group 2, Period 2) and Al (Group 13, Period 3) are diagonal neighbours. They have similar ionic charge densities despite different charges (Be²⁺ radius 31 pm; Al³⁺ radius 53 pm; both small and highly polarising). [1]
Three comparisons: [3]
(1) Amphoteric oxide/hydroxide: BeO and Be(OH)₂ react with both acids (BeO + 2HCl → BeCl₂ + H₂O) and bases (Be(OH)₂ + 2NaOH → Na₂[Be(OH)₄]). Al₂O₃ and Al(OH)₃ are also amphoteric. Other Group 2 oxides/hydroxides are only basic.
(2) Covalent halides: BeCl₂ is covalent (polymeric chain structure; Lewis acid); AlCl₃ is also covalent (dimeric, Al₂Cl⁶). Other Group 2 chlorides are ionic.
(3) Reaction with NaOH: Be reacts with NaOH: Be + 2NaOH + 2H₂O → Na₂[Be(OH)₄] + H₂. Al also dissolves in NaOH: 2Al + 2NaOH + 2H₂O → 2NaAlO₂ + 3H₂. Other Group 2 metals do not react with NaOH.
Three comparisons: [3]
(1) Amphoteric oxide/hydroxide: BeO and Be(OH)₂ react with both acids (BeO + 2HCl → BeCl₂ + H₂O) and bases (Be(OH)₂ + 2NaOH → Na₂[Be(OH)₄]). Al₂O₃ and Al(OH)₃ are also amphoteric. Other Group 2 oxides/hydroxides are only basic.
(2) Covalent halides: BeCl₂ is covalent (polymeric chain structure; Lewis acid); AlCl₃ is also covalent (dimeric, Al₂Cl⁶). Other Group 2 chlorides are ionic.
(3) Reaction with NaOH: Be reacts with NaOH: Be + 2NaOH + 2H₂O → Na₂[Be(OH)₄] + H₂. Al also dissolves in NaOH: 2Al + 2NaOH + 2H₂O → 2NaAlO₂ + 3H₂. Other Group 2 metals do not react with NaOH.
Q5 [4 marks]
Write balanced equations for: (a) Mg burning in oxygen [1]; (b) Ca reacting with water [1]; (c) Mg(OH)₂ reacting with HCl [1]; (d) thermal decomposition of Mg(NO₃)₂ [1]. [4]
(a) 2Mg + O₂ → 2MgO [1]
(b) Ca + 2H₂O → Ca(OH)₂ + H₂ [1]
(c) Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O [1]
(d) 2Mg(NO₃)₂ → 2MgO + 4NO₂ + O₂ [1]
(b) Ca + 2H₂O → Ca(OH)₂ + H₂ [1]
(c) Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O [1]
(d) 2Mg(NO₃)₂ → 2MgO + 4NO₂ + O₂ [1]
Section B — Extended Answer
30 marksQ6 [8 marks]
Compare the chemical properties of Group 2 elements (Be to Ba) in terms of: reactions with oxygen, water, and dilute acids. State the trend in reactivity for each, write equations where appropriate, and explain why reactivity increases down the group. Note any exceptions (e.g. Be). [8]
With oxygen [2]: All react to form MO (normal oxide). 2M + O₂ → 2MO. Reactivity increases down the group (Mg burns brightly in O₂ with brilliant white flame; Ba reacts vigorously). Ba may also form BaO₂ in excess O₂. Mg reacts when heated/ignited; Ca reacts more readily; Ba reacts even in air at room temperature.
With water [3]: Be: no reaction with water (even steam) due to protective oxide layer — anomalous. Mg: very slow with cold water; reacts with steam: Mg + H₂O(g) → MgO + H₂. Ca: slow reaction with cold water: Ca + 2H₂O → Ca(OH)₂ + H₂ (milky; effervescence). Sr: more vigorous than Ca. Ba: most vigorous; reacts readily with cold water to give Ba(OH)₂ + H₂. All products (except Be which doesn't react) are alkaline. Solutions turn litmus blue.
With dilute acids [2]: All react vigorously with dilute HCl to give MX₂ + H₂. Mg + 2HCl → MgCl₂ + H₂. Ca + 2HCl → CaCl₂ + H₂. Notable exception: Ca + H₂SO₄: initially reacts but forms insoluble CaSO₄ layer → passivation → reaction slows. Reactivity increases down the group.
Explanation of reactivity trend [1]: Both IE₁ and IE₂ decrease down Group 2: outer 2s electrons are in successively higher shells, farther from nucleus and more shielded → easier to remove → greater reactivity. Be is an exception due to its protective oxide layer.
With water [3]: Be: no reaction with water (even steam) due to protective oxide layer — anomalous. Mg: very slow with cold water; reacts with steam: Mg + H₂O(g) → MgO + H₂. Ca: slow reaction with cold water: Ca + 2H₂O → Ca(OH)₂ + H₂ (milky; effervescence). Sr: more vigorous than Ca. Ba: most vigorous; reacts readily with cold water to give Ba(OH)₂ + H₂. All products (except Be which doesn't react) are alkaline. Solutions turn litmus blue.
With dilute acids [2]: All react vigorously with dilute HCl to give MX₂ + H₂. Mg + 2HCl → MgCl₂ + H₂. Ca + 2HCl → CaCl₂ + H₂. Notable exception: Ca + H₂SO₄: initially reacts but forms insoluble CaSO₄ layer → passivation → reaction slows. Reactivity increases down the group.
Explanation of reactivity trend [1]: Both IE₁ and IE₂ decrease down Group 2: outer 2s electrons are in successively higher shells, farther from nucleus and more shielded → easier to remove → greater reactivity. Be is an exception due to its protective oxide layer.
Q7 [8 marks]
Describe the properties of Group 2 compounds: carbonates, hydroxides, and sulfates. For each compound type: state the trend in solubility and/or thermal stability down the group, explain the trend using lattice energy and polarisation arguments, and give a relevant industrial or biological use. [8]
Carbonates [3]: Thermal stability increases down the group (MgCO₃ ~540°C → BaCO₃ ~1360°C). Smaller M²⁺ (Mg²⁺) has high charge density → polarises CO²⁻ strongly → distorts/weakens C–O bonds → CO₂ released at lower temperature. Larger cations have lower charge density → less polarisation → more stable carbonates. All decompose as: MCO₃ → MO + CO₂. Use: CaCO₃ (limestone) → heated to make CaO (quicklime) for steelmaking and construction. MgCO₃ in antacids.
Hydroxides [3]: Solubility increases down the group: Mg(OH)₂ sparingly soluble → Ba(OH)₂ soluble. As M²⁺ gets larger, lattice energy decreases significantly (OH⁻ is small → lattice spacing dominated by M²⁺ → weaker when M²⁺ is large). Hydration energy decreases less steeply. Net: more favourable dissolution for larger M²⁺ → solubility increases. Also: basicity increases (Ba(OH)₂ > Ca(OH)₂ > Mg(OH)₂). Uses: Ca(OH)₂ for soil treatment and water purification; Mg(OH)₂ as antacid; Ba(OH)₂ in analytical chemistry.
Sulfates [2]: Solubility decreases: MgSO (soluble) → BaSO (insoluble). SO²⁻ is large → lattice energy barely changes down the group (dominated by large anion). Hydration energy of M²⁺ falls significantly as radius increases → insufficient energy to break lattice → solubility decreases. Use: BaSO in barium meal X-ray contrast agent (safe because insoluble); MgSO (Epsom salt) as Mg fertiliser and laxative; CaSO as gypsum/plaster of Paris.
Hydroxides [3]: Solubility increases down the group: Mg(OH)₂ sparingly soluble → Ba(OH)₂ soluble. As M²⁺ gets larger, lattice energy decreases significantly (OH⁻ is small → lattice spacing dominated by M²⁺ → weaker when M²⁺ is large). Hydration energy decreases less steeply. Net: more favourable dissolution for larger M²⁺ → solubility increases. Also: basicity increases (Ba(OH)₂ > Ca(OH)₂ > Mg(OH)₂). Uses: Ca(OH)₂ for soil treatment and water purification; Mg(OH)₂ as antacid; Ba(OH)₂ in analytical chemistry.
Sulfates [2]: Solubility decreases: MgSO (soluble) → BaSO (insoluble). SO²⁻ is large → lattice energy barely changes down the group (dominated by large anion). Hydration energy of M²⁺ falls significantly as radius increases → insufficient energy to break lattice → solubility decreases. Use: BaSO in barium meal X-ray contrast agent (safe because insoluble); MgSO (Epsom salt) as Mg fertiliser and laxative; CaSO as gypsum/plaster of Paris.
Q8 [6 marks]
Calcium is one of the most important Group 2 elements industrially. Describe the industrial chemistry involving calcium compounds: (a) the manufacture of CaO from limestone [2]; (b) the uses of CaO and Ca(OH)₂ [2]; (c) explain the limewater test for CO₂, including what happens with excess CO₂ [2]. [6]
(a) CaO manufacture [2]: CaCO₃ is mined as limestone, chalk, or marble. It is thermally decomposed in a lime kiln at ~900°C: CaCO₃(s) → CaO(s) + CO₂(g). This is an endothermic process requiring sustained high temperatures. CaO (quicklime) is a white, powdery solid. Industrially ~300 million tonnes produced annually worldwide. The CO₂ by-product is a greenhouse gas; modern kilns capture it.
(b) Uses of CaO and Ca(OH)₂ [2]: CaO uses: (1) Steel making — added to blast furnace to react with silica impurities: CaO + SiO₂ → CaSiO₃ (slag, removed separately). (2) Manufacture of Ca(OH)₂ by reaction with water (slaking). Ca(OH)₂ uses: (1) Agriculture — added to acidic soil to neutralise acid and raise pH: Ca(OH)₂ + 2H⁺ → Ca²⁺ + 2H₂O. (2) Water treatment — added to drinking water to adjust pH and precipitate Mg²⁺ as Mg(OH)₂. (3) Construction — in mortar/plaster (reacts with CO₂ to form CaCO, hardening). (4) Flue gas desulfurisation — removes SO₂: Ca(OH)₂ + SO₂ → CaSO + H₂O.
(c) Limewater test for CO₂ [2]: Limewater = saturated Ca(OH)₂ solution. Bubbling CO₂ through it forms insoluble white CaCO₃ precipitate: Ca(OH)₂ + CO₂ → CaCO₃(s) + H₂O → solution turns milky/cloudy. With excess CO₂: CaCO dissolves as soluble calcium hydrogen carbonate forms: CaCO₃ + CO₂ + H₂O → Ca(HCO₃)₂(aq) → solution clears. On heating Ca(HCO₃)₂: CaCO reprecipitates (reversal). This sequence (milky then clear with excess CO₂) confirms CO₂.
(b) Uses of CaO and Ca(OH)₂ [2]: CaO uses: (1) Steel making — added to blast furnace to react with silica impurities: CaO + SiO₂ → CaSiO₃ (slag, removed separately). (2) Manufacture of Ca(OH)₂ by reaction with water (slaking). Ca(OH)₂ uses: (1) Agriculture — added to acidic soil to neutralise acid and raise pH: Ca(OH)₂ + 2H⁺ → Ca²⁺ + 2H₂O. (2) Water treatment — added to drinking water to adjust pH and precipitate Mg²⁺ as Mg(OH)₂. (3) Construction — in mortar/plaster (reacts with CO₂ to form CaCO, hardening). (4) Flue gas desulfurisation — removes SO₂: Ca(OH)₂ + SO₂ → CaSO + H₂O.
(c) Limewater test for CO₂ [2]: Limewater = saturated Ca(OH)₂ solution. Bubbling CO₂ through it forms insoluble white CaCO₃ precipitate: Ca(OH)₂ + CO₂ → CaCO₃(s) + H₂O → solution turns milky/cloudy. With excess CO₂: CaCO dissolves as soluble calcium hydrogen carbonate forms: CaCO₃ + CO₂ + H₂O → Ca(HCO₃)₂(aq) → solution clears. On heating Ca(HCO₃)₂: CaCO reprecipitates (reversal). This sequence (milky then clear with excess CO₂) confirms CO₂.
Q9 [8 marks]
An unknown white solid X dissolves in water to give a colourless solution. When excess NaOH solution is added to the solution, a white gelatinous precipitate forms that dissolves in excess NaOH. Flame test on X gives no distinctive colour. (a) Identify X and the precipitate. [2] (b) Write equations for the two reactions described. [2] (c) Predict the products of heating solid X and write the equation. [2] (d) Describe two other tests that would confirm the identity of X, with expected results. [2]
(a) A white precipitate that dissolves in excess NaOH = amphoteric hydroxide → Be(OH)₂. No distinctive flame test colour and Group 2 behaviour → X contains Be²⁺. X is likely BeCl₂ (beryllium chloride) or BeS etc. The precipitate is Be(OH)₂. [2]
(b) With NaOH (precipitate forms): BeCl₂ + 2NaOH → Be(OH)₂(s) + 2NaCl [1]
Dissolving in excess NaOH: Be(OH)₂ + 2NaOH → Na₂[Be(OH)₄] [1]
(c) BeCl₂ is a covalent chloride; heating BeCl₂(s) would not typically give simple thermal decomposition (it is stable). However, if X is BeCO (beryllium carbonate): BeCO₃ → BeO + CO₂. More likely, if X is BeSO: decomposes at high temperature. Accept any sensible answer acknowledging BeCO or similar. If BeCl₂: no thermal decomposition under normal conditions. [2]
(d) Any two of: [1 each] (1) Add AgNO₃(aq) after acidifying with dilute HNO: if X is BeCl₂, a white precipitate of AgCl forms (insoluble in dilute HNO₃, soluble in NH(aq)) → confirms Cl⁻. (2) Add dilute H₂SO: if X is BeCl₂, no precipitate expected (BeSO is soluble). (3) ICP-MS or flame emission spectroscopy on solution would confirm Be (no visible flame, but emission at specific UV wavelengths). (4) Test for the anion: if sulfate, add BaCl₂(aq)/dilute HCl → white precipitate of BaSO confirms SO²⁻. [2]
(b) With NaOH (precipitate forms): BeCl₂ + 2NaOH → Be(OH)₂(s) + 2NaCl [1]
Dissolving in excess NaOH: Be(OH)₂ + 2NaOH → Na₂[Be(OH)₄] [1]
(c) BeCl₂ is a covalent chloride; heating BeCl₂(s) would not typically give simple thermal decomposition (it is stable). However, if X is BeCO (beryllium carbonate): BeCO₃ → BeO + CO₂. More likely, if X is BeSO: decomposes at high temperature. Accept any sensible answer acknowledging BeCO or similar. If BeCl₂: no thermal decomposition under normal conditions. [2]
(d) Any two of: [1 each] (1) Add AgNO₃(aq) after acidifying with dilute HNO: if X is BeCl₂, a white precipitate of AgCl forms (insoluble in dilute HNO₃, soluble in NH(aq)) → confirms Cl⁻. (2) Add dilute H₂SO: if X is BeCl₂, no precipitate expected (BeSO is soluble). (3) ICP-MS or flame emission spectroscopy on solution would confirm Be (no visible flame, but emission at specific UV wavelengths). (4) Test for the anion: if sulfate, add BaCl₂(aq)/dilute HCl → white precipitate of BaSO confirms SO²⁻. [2]