Describe and draw the structure of proteins as

IGCSE Chemistry (0620) – Complete Revision Notes

Learning Objective

Describe and draw the structure of proteins as polymers of amino‑acids and understand how this topic fits into the wider Cambridge IGCSE Chemistry syllabus.

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1. States of Matter & Particle Theory

• Particle model – Solids: particles vibrate in fixed positions.
  Liquids: particles move past one another.
  Gases: particles move freely and are far apart.
• Kinetic particle theory – Temperature = average kinetic energy; pressure = collisions of particles with container walls.
• Heating / cooling curves – Show energy absorbed or released during phase changes (solid → liquid → gas).
• Diffusion & effusion – Rate ↑ with temperature, ↓ with particle mass; for gases, rate ↓ with higher pressure.
• Ideal‑gas equation – \(PV = nRT\) (use \(R = 0.0821\; \text{L·atm·mol}^{-1}\text{K}^{-1}\)).
• Real‑gas behaviour – Deviations at high pressure / low temperature due to intermolecular forces.

Example (AO2)

Calculate the volume occupied by 0.50 mol of an ideal gas at 298 K and 1 atm.

\(V = \dfrac{nRT}{P}= \dfrac{0.50\times 0.0821\times 298}{1}=12.2\text{ L}\)

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2. Atomic Structure, the Periodic Table & Electron Configuration

• Atoms consist of protons, neutrons and electrons; atomic number (Z) = number of protons.
• Isotopes: same Z, different mass number (A).
• Electron configuration (Z = 1‑20)

Z	Element	Electron configuration
1	H	1s¹
2	He	1s²
3	Li	1s² 2s¹
4	Be	1s² 2s²
5	B	1s² 2s² 2p¹
6	C	1s² 2s² 2p²
7	N	1s² 2s² 2p³
8	O	1s² 2s² 2p⁴
9	F	1s² 2s² 2p⁵
10	Ne	1s² 2s² 2p⁶
11	Na	1s² 2s² 2p⁶ 3s¹
12	Mg	1s² 2s² 2p⁶ 3s²
13	Al	1s² 2s² 2p⁶ 3s² 3p¹
14	Si	1s² 2s² 2p⁶ 3s² 3p²
15	P	1s² 2s² 2p⁶ 3s² 3p³
16	S	1s² 2s² 2p⁶ 3s² 3p⁴
17	Cl	1s² 2s² 2p⁶ 3s² 3p⁵
18	Ar	1s² 2s² 2p⁶ 3s² 3p⁶
19	K	1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹
20	Ca	1s² 2s² 2p⁶ 3s² 3p⁶ 4s²

• Group‑number → ion charge rule (for Groups 1‑7):
  Group 1 → +1, Group 2 → +2, Group 13 → +3, Group 15 → –3, Group 16 → –2, Group 17 → –1.
• Key groups for the syllabus: Alkali metals (1), Alkaline earth metals (2), Halogens (17), Noble gases (18).

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3. Bonding

Bond type	Formation	Key properties
Ionic	Transfer of electrons from metal → non‑metal, forming cations + anions.	High melting/boiling points, soluble in water, conduct electricity when molten or aqueous.
Covalent (molecular)	Sharing of electrons between non‑metals.	Low melting/boiling points, poor conductors, often soluble in non‑polar solvents.
Metallic	Delocalised “sea of electrons” among metal cations.	Conduct electricity & heat, malleable, ductile, usually high melting points.
Giant covalent (e.g., SiO₂, diamond)	Each atom covalently bonded to many neighbours in a 3‑D network.	Very high melting points, hard, poor conductors (except graphite).

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4. Chemical Energetics

• Exothermic reactions – Release heat (ΔH < 0). Example: combustion of methane.
• Endothermic reactions – Absorb heat (ΔH > 0). Example: thermal decomposition of calcium carbonate.
• Energy‑profile diagrams show activation energy and overall enthalpy change.
• Heating curves illustrate the energy required for each phase change.

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5. Acids, Bases & Salts

• Acid: produces H⁺ in water (e.g., HCl → H⁺ + Cl⁻).
• Base: produces OH⁻ in water (e.g., NaOH → Na⁺ + OH⁻).
• Neutralisation: acid + base → salt + water (e.g., H₂SO₄ + 2 NaOH → Na₂SO₄ + 2 H₂O).
• pH scale: 0–14; strong acids pH ≤ 3, strong bases pH ≥ 11.

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6. The Mole Concept & Calculations

• Mole – 1 mol = \(6.02\times10^{23}\) particles (Avogadro’s constant).
• Molar mass (g mol⁻¹) = sum of atomic masses of the formula.
• Conversions –
  • mass (g) = moles × molar mass
  • moles = mass ÷ molar mass
  • particles = moles × \(6.02\times10^{23}\)
• Gas volume at r.t.p. – 1 mol of any gas occupies 24 L (298 K, 1 atm).
• Limiting‑reactant & % yield – Identify the reactant that produces the smallest amount of product, then: \[ \%\text{ yield}= \frac{\text{actual mass}}{\text{theoretical mass}}\times100 \]

Example (Limiting reactant)

2.00 g Na reacts with excess Cl₂: Na + ½ Cl₂ → NaCl.
M(Na)=23 g mol⁻¹ → 0.087 mol Na.
M(NaCl)=58.5 g mol⁻¹ → mass = 0.087 × 58.5 = 5.09 g NaCl.

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7. Metals – Reactivity & Extraction

• Reactivity series (most → least reactive): K > Na > Ca > Mg > Al > Zn > Fe > Sn > Cu > Ag > Au.
• More reactive metal displaces a less reactive one from its compound (e.g., Zn + CuSO₄ → ZnSO₄ + Cu).
• Extraction methods
  • Highly reactive metals (Na, K) – electrolysis of molten salts.
  • Less reactive metals (Fe, Cu) – reduction with carbon or CO.
  • Very unreactive metals (Au) – cyanide leaching.
• Corrosion – Oxidation of iron → Fe₂O₃·nH₂O (rust).

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8. Chemistry of the Environment

• Water‑testing parameters: pH, turbidity, dissolved oxygen, conductivity, nitrates, phosphates.
• Major pollutants
  • Acid rain – SO₂, NOₓ.
  • Greenhouse gases – CO₂, CH₄.
  • Ozone depletion – CFCs.
• Mitigation – Scrubbers, catalytic converters, renewable energy, recycling.
• Climate change – Global warming potential, carbon‑footprint calculations.

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9. Organic Chemistry – Core Topics

9.1 Naming & Functional Groups

• Alkanes (‑ane), alkenes (‑ene), alcohols (‑ol), carboxylic acids (‑anoic acid), esters (‑oate).
• Functional‑group symbols: –OH, –COOH, –C=O, –COO‑.

9.2 Isomerism

• Chain isomers – different carbon skeletons.
• Position isomers – functional group attached at different positions.
• Geometric (cis‑trans) isomers – applicable to alkenes with restricted rotation.

9.3 Hydrocarbons

• Alkanes – saturated, single bonds, formula CₙH₂ₙ₊₂.
• Alkenes – at least one C=C double bond, formula CₙH₂ₙ.
• Key reactions: combustion, substitution (alkanes), addition (alkenes).

9.4 Alcohols & Carboxylic Acids

• Alcohols – –OH attached to an sp³ carbon; combustion gives CO₂ + H₂O.
• Carboxylic acids – –COOH; neutralise with bases to give salts + water.
• Esters – formed by esterification (acid + alcohol ⇌ ester + water).

9.5 Fuels & Energy

• General combustion: \(C_xH_y + O_2 \rightarrow CO_2 + H_2O + \text{heat}\).
• Energy density: gasoline ≈ 44 MJ kg⁻¹, diesel ≈ 45 MJ kg⁻¹.

9.6 Polymers – Overview

• Monomer → polymer via condensation (loss of H₂O) or addition (radical) reactions.
• Examples: polyethene, polypropene, PVC, nylon, polyester, proteins.

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10. Polymers – Proteins (Detailed)

10.1 Amino‑Acid Monomer

General structure:

\[\mathrm{H_2N\!-\!CH(R)\!-\!COOH}\]

• R‑group – side chain that distinguishes the 20 common amino‑acids.
• Contains both an amine (–NH₂) and a carboxyl (–COOH) group → zwitterion at physiological pH.

10.2 Formation of the Peptide (Amide) Bond

Condensation reaction:

\[\mathrm{-COOH + H_2N- \;\longrightarrow\; -CO\!-\!NH- + H_2O}\]

• Repeats to give a long chain called a polypeptide.
• Each peptide bond is planar and rigid, restricting rotation about the C‑N axis.

10.3 Levels of Protein Structure

Level	Definition	Stabilising interactions / features
Primary	Linear sequence of amino‑acids linked by peptide bonds.	Covalent peptide bonds; determines all higher‑order structure.
Secondary	Regular folding of the backbone into α‑helices or β‑pleated sheets.	Hydrogen bonds between backbone –NH and –C=O groups.
Tertiary	Three‑dimensional shape of a single polypeptide.	Side‑chain interactions: hydrogen bonds, ionic bonds, hydrophobic packing, Van der Waals forces, and disulfide bridges (–S–S–).
Quaternary	Assembly of two or more polypeptide subunits into a functional protein.	Same non‑covalent forces as tertiary; additional disulfide bridges may link subunits.

10.4 Drawing a Short Protein Segment

Use the schematic below as a guide. Label the backbone atoms (N, Cα, C) and each side‑chain (R). Indicate peptide bonds and, if desired, hydrogen‑bond arrows for an α‑helix.

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11. Electrochemistry

• Electrolysis – Use of electricity to drive a non‑spontaneous redox reaction in the molten or aqueous state.
• Electrode identification
  • Anode – site of oxidation, negative in electrolytic cells.
  • Cathode – site of reduction, positive in electrolytic cells.
• Typical electrolytic cells

  Electrolyte	Anode (oxidation)	Cathode (reduction)	Products
  Molten NaCl	Cl⁻ → Cl₂ (g) + 2e⁻	Na⁺ + e⁻ → Na (l)	Cl₂(g) and Na(l)
  Aqueous CuSO₄	Cu(s) → Cu²⁺ + 2e⁻	2H₂O + 2e⁻ → H₂(g) + 2OH⁻	Cu(s) at anode, H₂(g) at cathode
  Molten CaCl₂	2Cl⁻ → Cl₂(g) + 2e⁻	Ca²⁺ + 2e⁻ → Ca(l)	Cl₂(g) and Ca(l)

• Calculations – Use Faraday’s laws: 1 F = 96 485 C = 1 mol e⁻. Mass of product = (Q ÷ F) × (atomic mass ÷ n), where n is electrons transferred.

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These notes cover every core topic of the Cambridge IGCSE Chemistry (0620) syllabus, include clear examples and tables, and give a detailed, exam‑focused description of protein structure.