Chemistry of selected elements: properties and reactions

Inorganic Chemistry – Chemistry of Selected Elements (Cambridge AS & A Level 9701)

1. Syllabus Scope, Learning Outcomes & Assessment Objectives

Scope covered in these notes (required AS topics):

• Group 2 – Alkaline‑Earth Metals
• Group 17 – Halogens
• Group 15 – Nitrogen (selected species)
• Group 16 – Sulphur (selected species)
• Periodicity of chemical properties (trends across the above groups)
• Optional/extension material: Group 1 – Alkali Metals and Transition Metals (useful for A‑Level questions)

Learning outcomes (mapped to Cambridge Assessment Objectives)

1. AO1 – Knowledge and understanding
   • Recall the electronic configurations, oxidation states and characteristic properties of the selected elements.
   • Write balanced equations for the reactions required by the syllabus.
2. AO2 – Application of knowledge
   • Explain trends in physical and chemical properties using atomic structure, bond type and lattice energy.
   • Predict the products of reactions involving the selected elements and justify the prediction.
   • Calculate enthalpy changes for selected reactions (e.g., formation of metal oxides, decomposition of carbonates).
3. AO3 – Practical skills and data handling
   • Interpret experimental data (e.g., gas evolution, mass‑loss, conductivity) to confirm reaction mechanisms.
   • Evaluate sources of error in common laboratory investigations (e.g., thermal decomposition of carbonates, preparation of hydrogen halides).
   • Design simple experiments to test a hypothesis about reactivity trends.

2. Key Concepts & Patterns (Cambridge “Key Concepts”)

1. Atoms, ions and forces – how charge, size and lattice energy influence solubility and thermal stability.
2. Experiments and evidence – linking observed colour, gas evolution, conductivity, and mass change to reaction pathways.
3. Patterns in chemical behaviour – periodic trends (ionisation energy, electronegativity, metallic/ non‑metallic character) that explain reactivity series.
4. Bonds, structures and properties – covalent vs ionic bonding in oxides, halides, and interhalogen compounds.
5. Energy changes – exothermic formation of oxides, endothermic decomposition of carbonates/nitrates, and the role of enthalpy in industrial processes.

3. Group 1 – Alkali Metals (optional/extension)

General Characteristics

• One valence electron; very low first ionisation energies (≈ 500 kJ mol⁻¹ for Li, ↓ to ≈ 380 kJ mol⁻¹ for Cs).
• Form +1 cations (M⁺) and highly basic oxides/hydroxides.
• Reactivity increases down the group (Li < Na < K < Rb < Cs).

Typical Reactions (syllabus‑relevant families)

1. Reaction with water

   General: M + H₂O → MOH + H₂↑

   Examples:

   • 2 Na + 2 H₂O → 2 NaOH + H₂↑
   • 2 K + 2 H₂O → 2 KOH + H₂↑ (vigorous, ignites H₂)
2. Formation of oxides, peroxides & superoxides
   • 2 Na + O₂ → Na₂O₂ (peroxide)
   • 2 K + O₂ → 2 KO₂ (superoxide)
3. Reaction with halogens

   2 M + X₂ → 2 MX (ionic halide)

4. Displacement reactions (metal‑metal)

   2 Na + CaCl₂ → 2 NaCl + Ca (Na displaces Ca²⁺)

Experimental note (AO3)

When Na is added to cold water, the temperature rise can be measured with a thermistor. Plot ΔT vs mass of Na to illustrate the quantitative relationship between reactivity and enthalpy released.

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4. Group 2 – Alkaline‑Earth Metals (required)

4.1 General Characteristics & Periodic Trends

Element	Electron config.	Atomic radius (pm)	First IE (kJ mol⁻¹)	Melting point (°C)
Be	[He] 2s²	112	899	1287
Mg	[Ne] 3s²	160	738	650
Ca	[Ar] 4s²	197	590	842
Sr	[Kr] 5s²	215	550	777
Ba	[Xe] 6s²	222	503	727

• Metallic character increases down the group – lower ionisation energy, larger atomic radius.
• Oxidation state is invariably +2 (exception: Be forms covalent compounds).
• Oxides are basic; they react with water to give hydroxides (except BeO, which is amphoteric).
• Solubility trends – hydroxides become more soluble down the group; sulfates become less soluble down the group.

4.2 Key Reaction Families (required by the syllabus)

1. Reaction with water (trend ↓ reactivity)
   • Be – no observable reaction.
   • Mg – very slow at room temperature; reaction accelerates when heated.
   • Ca, Sr, Ba – vigorous: \[ \text{M} + 2\text{H}_2\text{O} \rightarrow \text{M(OH)}_2 + \text{H}_2\uparrow \]
2. Reaction with oxygen (formation of oxides) \[ 2\text{M} + \text{O}_2 \rightarrow 2\text{MO} \] (e.g., 2 Mg + O₂ → 2 MgO, a white, high‑melting basic oxide).
3. Reaction with dilute acids \[ \text{M} + 2\text{HX} \rightarrow \text{MX}_2 + \text{H}_2\uparrow \] (HX = HCl, H₂SO₄ etc.)
4. Thermal decomposition of carbonates

   Trend: ↓ solubility → ↓ decomposition temperature.

   Carbonate	Decomposition temp. (°C)	Products
   MgCO₃	≈ 350	MgO + CO₂
   CaCO₃	≈ 825	CaO + CO₂
   SrCO₃	≈ 770	SrO + CO₂
   BaCO₃	≈ 850	BaO + CO₂

5. Thermal decomposition of nitrates \[ \text{M(NO}_3)_2 \xrightarrow{\Delta} \text{MO} + 2\text{NO}_2 + \tfrac12\text{O}_2 \] (Ca, Sr, Ba decompose at ≈ 550 °C; Mg(NO₃)₂ decomposes at ≈ 350 °C producing MgO.)
6. Solubility trends (selected salts)
   • Hydroxides:  Be(OH)₂ < Mg(OH)₂ < Ca(OH)₂ < Sr(OH)₂ ≈ Ba(OH)₂ (increasing solubility).
   • Sulfates:  BeSO₄ ≈ MgSO₄ > CaSO₄ ≈ SrSO₄ ≈ BaSO₄ (decreasing solubility).

4.3 Energy Changes (AO2)

Enthalpy of formation (ΔH_f°) of the oxides becomes more exothermic down the group (e.g., ΔH_f° MgO ≈ ‑601 kJ mol⁻¹, ΔH_f° CaO ≈ ‑635 kJ mol⁻¹). This explains the increasing vigor of the water reaction.

4.4 Practical Skills (AO3)

• Thermal decomposition of carbonates – set up a crucible‑tube, heat a known mass of carbonate, collect evolved CO₂ over water, and calculate % mass loss. Compare experimental decomposition temperature with literature values.
• Reactivity with water – measure the rate of H₂ evolution using a gas syringe for Ca, Sr and Ba; plot volume of H₂ vs time to illustrate the trend.
• Solubility determination – prepare saturated solutions of Mg(OH)₂, Ca(OH)₂, Sr(OH)₂ at 25 °C, filter, and titrate the filtrate with HCl to determine the concentration of dissolved OH⁻.

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5. Group 17 – Halogens (required)

5.1 General Characteristics & Periodic Trends

Element	State (25 °C)	Boiling point (°C)	Electronegativity (Pauling)	Bond dissociation energy, X–X (kJ mol⁻¹)
F	Gas	-188	3.98	158
Cl	Gas	-34	3.16	242
Br	Liquid	59	2.96	193
I	Solid	184	2.66	151

• All exist as diatomic molecules (X₂) and are strong oxidising agents.
• Reactivity decreases down the group (F > Cl > Br > I) because bond dissociation energy falls.
• Form –1 anions (X⁻) and a wide range of interhalogen compounds (X–X′).

5.2 Syllabus‑specified Reaction Families

1. Direct synthesis of hydrogen halides \[ \text{H}_2 + \text{X}_2 \rightarrow 2\text{HX} \] (Cl, Br, I – F requires special conditions; not examined).
2. Acid–base behaviour of halogen oxides
   • Formation of oxy‑acids: \[ \text{Cl}_2 + \text{H}_2\text{O} \rightarrow \text{HCl} + \text{HOCl} \] \[ \text{ClO}_2 + \text{H}_2\text{O} \rightarrow \text{HClO}_2 + \text{HOCl} \] \[ \text{Cl}_2\text{O}_7 + \text{H}_2\text{O} \rightarrow 2\text{HClO}_4 \] (strong acid)
   • General trend: higher oxidation state oxides give stronger acids.
3. Disproportionation in alkaline solution \[ \text{Cl}_2 + 2\text{OH}^- \rightarrow \text{Cl}^- + \text{ClO}^- + \text{H}_2\text{O} \] (used in water treatment)
   \[ \text{Br}_2 + 2\text{OH}^- \rightarrow \text{Br}^- + \text{BrO}^- + \text{H}_2\text{O} \] (bromine water for bleaching).
4. Reaction with metals – formation of ionic halides \[ 2\text{Na} + \text{Cl}_2 \rightarrow 2\text{NaCl} \] \[ \text{Mg} + \text{Cl}_2 \rightarrow \text{MgCl}_2 \] (Note: with Ag⁺, halides form precipitates – see 5.3).
5. Formation of interhalogen compounds (optional) \[ \text{Cl}_2 + \text{Br}_2 \rightarrow \text{BrCl}_3 \] \[ \text{I}_2 + \text{Cl}_2 \rightarrow \text{ICl}_5 \] (highly reactive, useful in synthesis).
6. Halide ion reactions (AO1/AO2)
   • Precipitation with Ag⁺: \[ \text{X}^- + \text{Ag}^+ \rightarrow \text{AgX(s)} \] (colour varies: AgCl white, AgBr pale yellow, AgI yellow).
   • Complex formation with NH₃: \[ \text{AgX} + 2\text{NH}_3 \rightarrow [\text{Ag(NH}_3)_2]^+ + \text{X}^- \] (solubilises the precipitate).
   • Oxidation by concentrated H₂SO₄ (Cl⁻ → Cl₂, Br⁻ → Br₂, I⁻ → I₂).

5.3 Energy Considerations

ΔH_f° values become less exothermic from F₂ (‑79 kJ mol⁻¹) to I₂ (‑62 kJ mol⁻¹). The decreasing bond energy explains why chlorine is the most commonly used industrial halogen despite fluorine’s higher electronegativity.

5.4 Practical Skills (AO3)

• Preparation of HCl gas – generate HCl by reacting NaCl with conc. H₂SO₄, collect over water, and verify by litmus test and silver nitrate test.
• Disproportionation experiment – add chlorine water to a known excess of NaOH, test the resulting solution with starch‑iodide paper for hypochlorite and with AgNO₃ for chloride.
• Halide identification – set up a series of test tubes containing unknown halide solutions; use AgNO₃ and NH₃ to differentiate Cl⁻, Br⁻ and I⁻ based on precipitate colour and solubility.

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6. Group 15 – Nitrogen (required)

6.1 Key Species & Their Industrial/Environmental Importance

• Ammonia (NH₃) – Haber‑Bosch synthesis: \[ \text{N}_2 + 3\text{H}_2 \xrightleftharpoons[450^\circ\text{C}]{200\ \text{atm}} 2\text{NH}_3 \] (Fe catalyst with K₂O & Al₂O₃ promoters).
• Ammonium ion (NH₄⁺) – formed by acid‑base reaction: \[ \text{NH}_3 + \text{HCl} \rightarrow \text{NH}_4\text{Cl} \] (used in fertilizers).
• Nitrogen oxides
  • NO (nitric oxide) – formed at high temperature: \[ \tfrac12\text{N}_2 + \tfrac12\text{O}_2 \rightarrow \text{NO} \] (ΔH ≈ +90 kJ mol⁻¹).
  • NO₂ (nitrogen dioxide) – oxidation of NO: \[ 2\text{NO} + \text{O}_2 \rightarrow 2\text{NO}_2 \] (ΔH ≈ ‑114 kJ mol⁻¹).
  • Acid rain pathway: \[ 3\text{NO}_2 + \text{H}_2\text{O} \rightarrow 2\text{HNO}_3 + \text{NO} \] (strong acid).
• Nitrous acid (HNO₂) & nitric acid (HNO₃) – formed when NO₂ dissolves in water: \[ \text{NO}_2 + \text{H}_2\text{O} \rightarrow \text{HNO}_3 + \text{HNO}_2 \] (equilibrium mixture).

6.2 Physical Data (selected)

Species	State (25 °C)	Boiling point (°C)
N₂	Gas	-196
NH₃	Gas	-33.3
NO	Gas	-152
NO₂	Gas	21.2

6.3 Acid–Base Behaviour of Nitrogen Oxides (AO1/AO2)

• NO is a neutral oxide (does not react with water).
• NO₂ is an acidic oxide; dissolves to give a mixture of HNO₃ (strong) and HNO₂ (weak).
• HNO₃ fully dissociates in water: \[ \text{HNO}_3 \rightarrow \text{H}^+ + \text{NO}_3^- \] (important for nitrate salts).

6.4 Energy & Thermodynamics (AO2)

ΔH_f° values: NH₃ (‑46 kJ mol⁻¹), NO (+90 kJ mol⁻¹), NO₂ (‑57 kJ mol⁻¹). The exothermic formation of NH₃ explains the high temperature required to overcome the activation energy but also the large heat release that can be recovered in the Haber process.

6.5 Practical Skills (AO3)

• Preparation of NO₂ – oxidise copper turnings in concentrated HNO₃, collect the brown gas, and confirm by its brown colour and by its ability to turn moist blue litmus red.
• Quantitative analysis of ammonia – generate NH₃ by adding NaOH to solid NH₄Cl, capture in a known volume of water, and titrate with standard HCl using an indicator.
• Haber‑Bosch simulation – use a sealed glass tube with N₂ and H₂ at 1 atm, heat to 400 °C, and monitor pressure change to infer conversion to NH₃ (demonstrates Le Chatelier’s principle).

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7. Group 16 – Sulphur (required)

7.1 Key Oxides & Their Acid–Base Behaviour

• Sulphur dioxide (SO₂) – produced by combustion of sulphur or metal sulphides. \[ \text{S} + \text{O}_2 \rightarrow \text{SO}_2 \] In water: \[ \text{SO}_2 + \text{H}_2\text{O} \rightleftharpoons \text{H}_2\text{SO}_3 \] (weak diprotic acid).
• Sulphur trioxide (SO₃) – oxidation of SO₂ (contact process): \[ 2\text{SO}_2 + \text{O}_2 \xrightleftharpoons[450^\circ\text{C}]{\text{V}_2\text{O}_5} 2\text{SO}_3 \] Reacts with water to give strong sulphuric acid: \[ \text{SO}_3 + \text{H}_2\text{O} \rightarrow \text{H}_2\text{SO}_4 \]
• Sulphuric acid (H₂SO₄) – strong diprotic acid; reacts with metals, bases and carbonates.

7.2 Industrial Reactions (AO1/AO2)

1. Contact process (manufacture of H₂SO₄)
   • Oxidation of SO₂ to SO₃ over V₂O₅ catalyst (exothermic, ΔH ≈ ‑98 kJ mol⁻¹).
   • Absorption of SO₃ in 98 % H₂SO₄ to form oleum (H₂S₂O₇), then diluted to obtain commercial H₂SO₄.
2. Reaction of sulphuric acid with metals (example: Zn) \[ \text{Zn} + 2\text{H}_2\text{SO}_4 \rightarrow \text{ZnSO}_4 + \text{SO}_2 + 2\text{H}_2\text{O} \] (concentrated acid, red‑brown SO₂ gas).
3. Formation of sulphates – typical solubility trend: \[ \text{BeSO}_4, \text{MgSO}_4 \text{(soluble)} > \text{CaSO}_4 \text{(sparingly soluble)} > \text{SrSO}_4, \text{BaSO}_4 \text{(practically insoluble)} \]

7.3 Energy & Thermodynamics (AO2)

ΔH_f° values: SO₂ (‑296 kJ mol⁻¹), SO₃ (‑396 kJ mol⁻¹), H₂SO₄ (‑814 kJ mol⁻¹). The large exothermicity of SO₃ formation drives the contact process.

7.4 Practical Skills (AO3)

• Preparation of SO₂ – burn elemental sulphur in a test tube, collect the gas over water, and confirm by its characteristic pungent smell and by turning moist blue litmus red.
• Quantitative absorption of SO₂ in NaOH – bubble known volume of SO₂ into excess NaOH, then titrate the resulting Na₂SO₃ solution with standard H₂O₂ or KMnO₄ to determine the amount of SO₂ absorbed.
• Determination of sulphuric acid concentration – use a standardised Na₂CO₃ solution to titrate a sample of H₂SO₄ (phenolphthalein endpoint) and calculate molarity.

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8. Transition Metals (optional/extension – A‑Level)

General Features

• Variable oxidation states (commonly +2, +3, sometimes +1, +4, +6).
• Formation of coloured complex ions; ligand field theory explains d‑orbital splitting.
• Often act as catalysts (e.g., Fe in Haber, V₂O₅ in contact process, Cu/Zn in methanol synthesis).
• Typical reactions: formation of coordination complexes, redox interconversions, precipitation of insoluble hydroxides.

Representative Example – Copper(II) Chemistry

1. Formation of [Cu(H₂O)₆]²⁺ (blue aqueous solution) \[ \text{CuSO}_4·5\text{H}_2\text{O} \xrightarrow{\text{H}_2\text{O}} [\text{Cu(H}_2\text{O})_6]^{2+} + \text{SO}_4^{2-} \]
2. Ammonia complex \[ [\text{Cu(H}_2\text{O})_6]^{2+} + 4\text{NH}_3 \rightarrow [\text{Cu(NH}_3)_4(\text{H}_2\text{O})_2]^{2+} + 4\text{H}_2\text{O} \] (deep blue).
3. Reduction to Cu(I) by sulphite \[ 2\text{Cu}^{2+} + \text{SO}_3^{2-} + \text{H}_2\text{O} \rightarrow 2\text{Cu}^+ + \text{SO}_4^{2-} + 2\text{H}^+ \]

Practical (AO3)

Prepare a series of copper(II) complexes with different ligands (NH₃, H₂O, Cl⁻) and record their visible spectra. Discuss the colour changes in terms of crystal‑field splitting energy (Δ₀).

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9. Periodicity of Chemical Properties (required)

Trends Across the Selected Groups

Property	Group 2 (down)	Group 17 (down)	Group 15 (down)	Group 16 (down)
Atomic radius	↑	↑	↑	↑
First ionisation energy	↓	↓	↓	↓
Electronegativity	↓	↓	↓	↓
Metallic character	↑	↓ (non‑metal to metal)	↑ (from non‑metal to metalloid)	↑ (non‑metal to metalloid)
Basicity of oxides	↓ (more basic down group)	↓ (oxidising power ↓)	↑ (from N₂ to NH₃)	↑ (SO₂ < SO₃ acidity ↑)
Solubility of hydroxides	↑ down group	—	—	—
Solubility of sulfates	↓ down group	—	—	—

Link to Key Concepts

• Decreasing ionisation energy down a group explains the increased reactivity of Ba compared with Be (Group 2) and of I compared with F (Group 17).
• Changes in lattice energy account for the solubility trends of hydroxides and sulfates.
• Bond dissociation energy trends (X–X) rationalise the halogen reactivity series and the ease of forming interhalogen compounds.

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10. Summary of Key Points (required)

• Group 2 metals form +2 cations, basic oxides, and show increasing reactivity with water down the group