Describe the change from metallic to non-metallic character across a period

IGCSE Chemistry (0620) – Complete Revision Notes

Learning Objective

Describe the change from metallic to non‑metallic character across a period and understand how this trend integrates with the wider Cambridge IGCSE Chemistry syllabus.

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1. States of Matter

• Solid, liquid, gas – particle arrangement, kinetic energy and intermolecular forces.
• Changes of state – melting, boiling, sublimation; always endothermic (energy absorbed).
• Gas laws (core) – pV = nRT. Qualitative trends:
  • Increasing temperature → volume ↑ (at constant pressure).
  • Increasing pressure → volume ↓ (at constant temperature).
• Diffusion (core & supplement) – spontaneous mixing of gases or vapours; rate increases with temperature and decreases with molecular mass.
• Key formulae – p = F/A, V = nRT/p, ρ = m/V.

2. Atoms, Elements & Compounds

2.1 Atomic Structure (core)

• Protons (+), neutrons (0), electrons (–). Atomic number (Z) = number of protons.
• Isotopes – same Z, different mass number (A). Example: ¹²C, ¹³C.
• Effective nuclear charge (Z_eff) = Z – S (shielding). Increases across a period.

2.2 Electronic Configuration (core, 1–20)

Element	Z	Configuration
H	1	1s¹
He	2	1s²
Li	3	1s² 2s¹
Be	4	1s² 2s²
B	5	1s² 2s² 2p¹
C	6	1s² 2s² 2p²
N	7	1s² 2s² 2p³
O	8	1s² 2s² 2p⁴
F	9	1s² 2s² 2p⁵
Ne	10	1s² 2s² 2p⁶

2.3 Mixtures, Elements & Compounds (core)

• Elements – made of one type of atom; cannot be broken down chemically.
• Compounds – two or more elements chemically combined in fixed ratios; can be broken down by chemical reactions.
• Mixtures – physical combinations of two or more substances; components retain their own properties and can be separated physically.

2.4 Ions and Bonding (core)

• Ion formation – metals lose electrons → cations; non‑metals gain electrons → anions.
• Ionic bonding – electrostatic attraction between oppositely charged ions (e.g., NaCl).
• Covalent bonding – sharing of electron pairs; represented by dot‑and‑cross diagrams (e.g., H₂O).
• Metallic bonding – delocalised “sea of electrons” surrounding a lattice of metal cations (giant‑metallic structure). Explains conductivity, malleability and high melting points.
• Alloys (core) – mixtures of two or more metals (or a metal and a non‑metal) that form a single phase, e.g., brass (Cu + Zn), steel (Fe + C).
• Formula writing – empirical vs. molecular formulas; e.g., glucose: empirical C₆H₁₂O₆, molecular C₆H₁₂O₆ (same in this case).

3. Stoichiometry

• Relative atomic mass (Ar) and relative molecular mass (Mr).
• Mole concept (core) – 1 mol = 6.02 × 10²³ particles; n = m/Mr.
• Mass‑mass, mass‑mole, mole‑mole calculations; limiting‑reactant and percentage‑yield (supplement).
• Empirical & molecular formulas (supplement) – derived from percent composition or combustion analysis.
• “Mole‑free” shortcuts for common compounds (e.g., 1 g H ≈ 8 g O in water).

4. Electrochemistry

• Electrolytic cells (core) – external power drives a non‑spontaneous reaction.
  • Molten NaCl → Na (cathode) + Cl₂ (anode).
  • Aqueous CuSO₄ → Cu (cathode) + O₂/H₂ (anode, depending on electrode material).
• Half‑equations (supplement) – essential for balancing electrolysis.

  Cathode:  Na⁺ + e⁻ → Na
  Anode:   2Cl⁻ → Cl₂ + 2e⁻
          

• Product prediction for binary molten compounds (core) – metal cation reduced, non‑metal anion oxidised.
• Fuel cells (core) – spontaneous redox reaction produces electricity (e.g., H₂ + ½ O₂ → H₂O).
• Electroplating (core) – deposition of a metal onto a surface using an electrolytic cell.

5. Chemical Energetics

• Exothermic vs. endothermic – energy released (ΔH < 0) or absorbed (ΔH > 0).
• Enthalpy change (ΔH) – measured by calorimetry; sign convention must be stated explicitly.
• Bond‑energy diagram – breaking bonds requires energy; forming bonds releases energy.
• Approximate calculation:
  ΔH ≈ Σ(bond energies broken) – Σ(bond energies formed).
• Activation energy (supplement) – minimum energy required for a reaction to proceed; explains effect of temperature on rate.

6. Chemical Reactions

6.1 Physical vs. Chemical Change (core)

• Physical change – state or form changes, composition unchanged (e.g., melting ice).
• Chemical change – new substances formed, composition altered (e.g., rusting of iron).

6.2 Rate of Reaction (core)

• Factors: concentration, temperature, surface area, catalyst.
• Collision theory – effective collisions required; activation energy is the barrier.

6.3 Equilibrium (core)

• Reversible reactions: A + B ⇌ C + D.
• Dynamic equilibrium – forward and reverse rates equal.
• Le Chatelier’s principle – predicts shift when concentration, temperature, pressure or catalyst changes.
• Typical reversible examples: CuSO₄·5H₂O ⇌ CuSO₄ + 5H₂O, NH₄Cl(s) ⇌ NH₃(g) + HCl(g).

6.4 Redox (core & supplement)

• Oxidation – loss of electrons, increase in oxidation number.
• Reduction – gain of electrons, decrease in oxidation number.
• Redox can be identified by:
  • Change in oxidation numbers.
  • Transfer of electrons in half‑equations.
• Example:
  2 Mg + O₂ → 2 MgO
  Mg: 0 → +2 (oxidised); O: 0 → –2 (reduced).

7. Acids, Bases & Salts

• Acids – sour taste, turn blue litmus red, produce H⁺ in water (e.g., HCl, H₂SO₄).
• Strong acids – dissociate completely (HCl, H₂SO₄, HNO₃, HBr, HI).
• Weak acids – only partially dissociate (CH₃COOH, H₂CO₃).
• Bases – slippery feel, turn red litmus blue, produce OH⁻ (e.g., NaOH, Ca(OH)₂).
• Strong bases – soluble metal hydroxides (NaOH, KOH, Ca(OH)₂).
• Weak bases – ammonia (NH₃) and amines.
• Neutralisation – acid + base → salt + water (e.g., H₂SO₄ + 2 NaOH → Na₂SO₄ + 2 H₂O).
• pH scale (0–14) – indicator colour chart (core):

  pH	Colour (Universal Indicator)
  0‑1	Red
  2‑3	Orange
  4‑5	Yellow
  6‑7	Green
  8‑9	Blue‑green
  10‑11	Blue
  12‑14	Purple

8. The Periodic Table – Layout & General Trends

• Groups (vertical) – elements in the same group have similar chemical properties.
• Periods (horizontal) – properties change progressively from left to right.

8.1 Trends Across a Period

• Atomic radius ↓ – electrons are pulled closer by increasing Z_eff.
• Ionisation energy ↑ – more energy required to remove the outer electron.
• Electronegativity ↑ – stronger ability to attract bonding electrons.
• Metallic character ↓ – elements become less willing to lose electrons.

8.2 Trends Down a Group

• Atomic radius ↑ – additional electron shells.
• Ionisation energy ↓ – outer electrons are farther from the nucleus.
• Electronegativity ↓.
• Metallic character ↑ – greater tendency to lose electrons.

8.3 Group‑Specific Patterns (core)

• Group 1 (alkali metals) – always form +1 ions; very reactive, low IE and EN.
• Group 2 (alkaline earth metals) – form +2 ions; slightly less reactive than Group 1.
• Group 7 (halogens) – form –1 ions; high EN, strong oxidising agents.
• Group 18 (noble gases) – full valence shells, essentially inert.
• Transition metals (Groups 3‑12) – variable oxidation states, form coloured compounds, often act as catalysts.

9. Metallic → Non‑Metallic Character Across a Period

9.1 Why the Trend Occurs

1. Electrons are added to the same principal energy level; Z_eff felt by valence electrons increases.
2. Higher Z_eff pulls the electron cloud closer → atomic radius decreases.
3. Smaller atoms hold outer electrons more tightly → first ionisation energy rises.
4. Greater ability to attract electrons in a bond → electronegativity rises.
5. Result: elements on the left readily lose electrons (metallic, reducing agents); those on the right readily gain electrons (non‑metallic, oxidising agents).

9.2 Illustrative Data – Period 2

Group	Element	Metallic Character	Electronegativity (Pauling)	First IE (kJ mol⁻¹)	Typical Oxidation State(s)
1	Li	High	0.98	520	+1
2	Be	Moderate	1.57	900	+2
13	B	Low	2.04	800	+3, –3
14	C	Very Low	2.55	1086	+4, –4
15	N	Negligible	3.04	1402	–3, +5
16	O	None	3.44	1314	–2
17	F	None	3.98	1681	–1

9.3 Illustrative Data – Period 3 (selected)

Group	Element	Metallic Character	Electronegativity	First IE (kJ mol⁻¹)	Typical Oxidation State(s)
1	Na	High	0.93	496	+1
2	Mg	Moderate	1.31	738	+2
13	Al	Low	1.61	578	+3
14	Si	Very Low	1.90	787	+4, –4
15	P	Negligible	2.19	1012	–3, +5
16	S	None	2.58	1000	–2, +6
17	Cl	None	3.16	1251	–1

9.4 Summary of the Trend

• Left‑hand elements (e.g., Li, Na) are strong reducing agents, form basic oxides and ionic compounds.
• Right‑hand elements (e.g., F, Cl) are strong oxidising agents, form acidic oxides and covalent compounds.
• Both ionisation energy and electronegativity increase together, reflecting the shift from electron‑donating to electron‑accepting behaviour.

9.5 Chemical Behaviour Implications

• Metals → metallic bonds, malleable, conductive, form cations; typical reactions:
  • Metal + acid → salt + H₂ (e.g., Zn + 2 HCl → ZnCl₂ + H₂).
  • Metal + non‑metal oxide → salt (e.g., 2 Mg + O₂ → 2 MgO).
• Non‑metals → covalent bonds, poor conductors, form anions or share electrons; typical reactions:
  • Non‑metal + metal oxide → acid (e.g., C + 2 MgO → 2 MgCO₃ → CO₂ + H₂O).
  • Halogen + metal → ionic halide (e.g., 2 Na + Cl₂ → 2 NaCl).

10. Metals – Properties, Extraction & Reactivity

• Physical properties – shiny, ductile, malleable, good conductors of heat and electricity.
• Reactivity series (core) – K > Na > Ca > Mg > Al > Zn > Fe > Sn > Pb > (less reactive) Cu > Ag > Au.
• Extraction methods
  • Electrolysis – e.g., Al from Al₂O₃ (Hall‑Héroult process).
  • Reduction with carbon – e.g., Fe from Fe₂O₃ in a blast furnace.
  • Distillation – e.g., Zn from ZnO (volatilisation and condensation).
• Corrosion – oxidation of metals, especially iron; prevention by coating, alloying, cathodic protection.
• Alloys (core) – single‑phase mixtures of metals (or metal + non‑metal) with improved properties; examples: brass (Cu + Zn), stainless steel (Fe + Cr + Ni).

11. The Environment

• Green chemistry principles – waste minimisation, atom economy, safer solvents, energy efficiency.
• Key pollutants – CO₂ (global warming), NOₓ (acid rain), SO₂ (acid rain), particulate matter.
• Water‑testing methods (core)
  • Dry‑cobalt(II) chloride test – colour change (blue → pink) indicates presence of water vapour.
  • Copper(II) sulphate test – blue crystals dissolve in water, indicating moisture.
  • Conductivity test – pure water is a poor conductor; presence of ions (from dissolved salts) increases conductivity.
• Acid rain – formation of H₂SO₄ and HNO₃ from SO₂ and NOₓ; impacts on buildings, soils and aquatic life.
• Ozone depletion – catalytic destruction by CFCs; relevance to the chemistry of halogens.