2.1 Atomic structure and the Periodic Table
Cambridge (CIE) IGCSE Chemistry Revision Notes
📌 Topic 2: Atoms, Elements and Compounds
2.1 Elements, compounds and mixtures
1 Describe the differences between elements, compounds and mixtures
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| Type | Definition | Example | Separation |
|---|---|---|---|
| Element | Pure substance made of only one type of atom | O₂, Fe, Au | Cannot be broken down |
| Compound | Two or more elements chemically bonded | H₂O, CO₂, NaCl | Only by chemical reaction |
| Mixture | Substances physically combined, not bonded | Air, salt water | Physical methods (filtration, evaporation) |
💡 Key Point: Elements cannot be split into simpler substances. Compounds have fixed ratios. Mixtures have variable composition.
2.2 Atomic structure and the Periodic Table
1 Describe the structure of the atom as a central nucleus containing neutrons and protons surrounded by electrons in shells
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Atomic Structure:
- Nucleus (centre): Contains protons (+) and neutrons (neutral)
- Electron shells: Electrons (-) orbit the nucleus in energy levels (shells)
Electron shells (energy levels)
/ \
/ \
| ⚫ | ← Nucleus (protons + neutrons)
\ /
\ /
-----
2 State the relative charges and relative masses of a proton, a neutron and an electron
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| Particle | Relative Charge | Relative Mass | Location |
|---|---|---|---|
| Proton | +1 | 1 | Nucleus |
| Neutron | 0 | 1 | Nucleus |
| Electron | -1 | 1/1840 (negligible) | Electron shells |
3 Define proton number/atomic number as the number of protons in the nucleus of an atom
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Atomic Number (Z) = Number of protons in the nucleus
In a neutral atom, the number of protons = number of electrons
💡 Example: Carbon has atomic number 6 → 6 protons, 6 electrons
4 Define mass number/nucleon number as the total number of protons and neutrons in the nucleus of an atom
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Mass Number (A) = Number of protons + number of neutrons
Number of neutrons = Mass number - Atomic number
💡 Example: Carbon-12 → 12 - 6 = 6 neutrons
5 Determine the electronic configuration of elements and their ions with proton number 1 to 20
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| Element | Atomic Number | Electronic Configuration |
|---|---|---|
| Hydrogen | 1 | 1 |
| Helium | 2 | 2 |
| Lithium | 3 | 2,1 |
| Beryllium | 4 | 2,2 |
| Boron | 5 | 2,3 |
| Carbon | 6 | 2,4 |
| Nitrogen | 7 | 2,5 |
| Oxygen | 8 | 2,6 |
| Fluorine | 9 | 2,7 |
| Neon | 10 | 2,8 |
| Sodium | 11 | 2,8,1 |
| Magnesium | 12 | 2,8,2 |
| Aluminium | 13 | 2,8,3 |
| Silicon | 14 | 2,8,4 |
| Phosphorus | 15 | 2,8,5 |
| Sulfur | 16 | 2,8,6 |
| Chlorine | 17 | 2,8,7 |
| Argon | 18 | 2,8,8 |
| Potassium | 19 | 2,8,8,1 |
| Calcium | 20 | 2,8,8,2 |
💡 Shell Rules: 1st shell = max 2 electrons, 2nd shell = max 8 electrons, 3rd shell = max 8 electrons (simplified model)
6 State that Group VIII noble gases have a full outer electron shell; the number of outer shell electrons equals the group number; the number of occupied electron shells equals the period number
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Relationship to Periodic Table:
- Period number = Number of occupied electron shells
- Group number (I-VII) = Number of electrons in the outer shell
- Group VIII (Noble gases) = Full outer shell (stable, unreactive)
💡 Example: Sodium (Period 3, Group 1) → 3 shells, 1 outer electron → 2,8,1
2.3 Isotopes
1 Define isotopes as different atoms of the same element that have the same number of protons but different numbers of neutrons
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Definition: Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.
Example: Chlorine
- Chlorine-35: 17 protons, 18 neutrons
- Chlorine-37: 17 protons, 20 neutrons
💡 Isotopes have the same chemical properties (same electrons) but different physical properties (different mass).
2 Interpret and use symbols for atoms and ions
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Atomic Notation: ¹²₆C
- Top number (12) = Mass number (protons + neutrons)
- Bottom number (6) = Atomic number (protons)
💡 Example: ³⁵₁₇Cl⁻ means a chloride ion with 17 protons, 18 neutrons, 18 electrons
3 State that isotopes have the same chemical properties because they have the same number of electrons Supplement
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Why Same Chemical Properties? Chemical properties depend on electron configuration. All isotopes of an element have the same number of electrons, so they react the same way chemically.
Why Different Physical Properties? Different mass affects physical properties like density, boiling point, and rate of diffusion.
4 Calculate the relative atomic mass from isotopic abundances Supplement
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Formula: Aᵣ = ( % of isotope A × mass A ) + ( % of isotope B × mass B ) / 100
Worked Example: Chlorine
- Chlorine-35: 75% abundance
- Chlorine-37: 25% abundance
Aᵣ = (75 × 35) + (25 × 37) / 100 = (2625 + 925) / 100 = 3550 / 100 = 35.5
2.4 Ions and ionic bonds
1 Describe the formation of positive ions (cations) and negative ions (anions)
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Cations (positive ions): Metals lose electrons to achieve a full outer shell
- Example: Na → Na⁺ + e⁻ (sodium loses 1 electron)
Anions (negative ions): Non-metals gain electrons to achieve a full outer shell
- Example: Cl + e⁻ → Cl⁻ (chlorine gains 1 electron)
💡 Group 1 metals form 1⁺ ions, Group 2 form 2⁺, Group 6 form 2⁻, Group 7 form 1⁻
2 State that an ionic bond is a strong electrostatic attraction between oppositely charged ions
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An ionic bond is the strong electrostatic attraction between oppositely charged ions.
Forms when a metal reacts with a non-metal.
3 Describe the formation of ionic bonds between Group I and Group VII using dot-and-cross diagrams
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Sodium Chloride (NaCl)
Na atom: 2,8,1 Cl atom: 2,8,7
↓ ↓
Na⁺ ion: 2,8 Cl⁻ ion: 2,8,8
[Na]⁺ [ Cl ]⁻
••
••
💡 Group I metals lose 1 electron → +1 charge. Group VII non-metals gain 1 electron → -1 charge.
4 Describe the properties of ionic compounds: high melting/boiling points; good conductivity when aqueous or molten, poor when solid
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| Property | Explanation |
|---|---|
| High melting/boiling points | Strong electrostatic forces between ions require lots of energy to overcome |
| Conducts when molten/dissolved | Ions are free to move and carry charge |
| Poor conductor when solid | Ions are fixed in position in the lattice |
| Hard but brittle | Layers shift, like charges align and repel |
5 Describe the giant lattice structure of ionic compounds Supplement
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Giant Lattice Structure: Ions are arranged in a regular, repeating 3D pattern with alternating positive and negative ions.
+ - + -
- + - +
+ - + -
- + - +
(Alternating positive and negative ions)
6 Explain the properties of ionic compounds in terms of structure and bonding Supplement
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- High melting/boiling points: Strong electrostatic forces between oppositely charged ions in all directions require large amounts of energy to overcome.
- Conductivity when molten/dissolved: Ions become free to move and can carry electric charge.
- No conductivity when solid: Ions are fixed in lattice positions and cannot move.
- Brittleness: When layers shift, like charges align and repel, causing the crystal to split.
2.5 Simple molecules and covalent bonds
1 State that a covalent bond is formed when a pair of electrons is shared between two atoms
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Covalent Bond: Formed when a pair of electrons is shared between two non-metal atoms.
Each atom contributes one electron to the shared pair, and both atoms achieve a full outer shell (noble gas configuration).
2 Describe the formation of covalent bonds in simple molecules using dot-and-cross diagrams
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H₂: H• + •H → H:H Cl₂: Cl•• + ••Cl → Cl:Cl (with 6 other electrons each) H₂O: H:O:H (O shares 2 pairs) CH₄: H:C:H with H above and below (C shares 4 pairs) NH₃: H:N:H with lone pair on N (N shares 3 pairs) HCl: H:Cl (H shares 1 pair)
💡 Only outer shell electrons are shown in dot-and-cross diagrams.
3 Describe the properties of simple molecular compounds: low melting/boiling points; poor electrical conductivity
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| Property | Explanation |
|---|---|
| Low melting/boiling points | Weak intermolecular forces between molecules (not the strong covalent bonds inside molecules) |
| Poor electrical conductivity | No free electrons or ions to carry charge |
| Often gases or liquids at room temperature | Weak forces mean they evaporate easily |
4 Describe the formation of covalent bonds in additional molecules (O₂, N₂, CO₂, C₂H₄, CH₃OH) Supplement
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O₂: O::O (double bond)
N₂: N:::N (triple bond)
CO₂: O::C::O (double bonds)
C₂H₄: H₂C=CH₂ (double bond between C)
CH₃OH: H
|
H—C—O—H
|
H
5 Explain the properties of simple molecular compounds in terms of structure and bonding Supplement
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Key Concept:
- Intramolecular forces (covalent bonds) = strong, hold atoms together within a molecule
- Intermolecular forces = weak, hold molecules together
- Low melting/boiling points because only weak intermolecular forces need to be overcome
- Poor conductivity because no free electrons or ions are present
2.6 Giant covalent structures
1 Describe the giant covalent structures of graphite and diamond
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Diamond: Each carbon atom bonded to 4 others (tetrahedral) - all strong covalent bonds
Graphite: Each carbon bonded to 3 others, forming layers of hexagons. Strong covalent bonds within layers; weak intermolecular forces between layers. One free electron per carbon atom becomes delocalised.
2 Relate the structures and bonding of graphite and diamond to their uses
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| Property | Diamond | Graphite |
|---|---|---|
| Bonding | 4 covalent bonds per carbon | 3 covalent bonds per carbon |
| Structure | 3D tetrahedral | Layered hexagonal |
| Hardness | Very hard (hardest natural substance) | Soft, slippery |
| Conductivity | Non-conductor (no free electrons) | Conductor (delocalised electrons) |
| Uses | Cutting tools, drills, jewellery | Pencil lead, lubricant, electrodes |
2.7 Metallic bonding
1 Describe metallic bonding as the electrostatic attraction between positive ions and a 'sea' of delocalised electrons Supplement
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Metallic Bonding:
- Metal atoms lose their outer electrons to become positive ions
- These electrons become delocalised (free to move) forming a "sea of electrons"
- The metallic bond is the electrostatic attraction between positive metal ions and the delocalised electrons
Metal lattice: + + + +
+ + + + (with sea of electrons around)
+ + + +
2 Explain the properties of metals: good electrical conductivity; malleability and ductility Supplement
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| Property | Explanation |
|---|---|
| Good electrical conductivity | Delocalised electrons are free to move and carry charge |
| Good thermal conductivity | Delocalised electrons transfer kinetic energy quickly |
| High melting/boiling points | Strong electrostatic forces between ions and electrons require lots of energy to overcome |
| Malleable (can be hammered into shape) | Layers of positive ions can slide over each other without breaking the metallic bond |
| Ductile (can be drawn into wires) | Same reason as malleability - layers can slide |
⚠️ Common Mistakes to Avoid:
• ❌ "Atoms expand when heated" → ✅ Particles gain energy and move more, they don't expand
• ❌ "Isotopes have different chemical properties" → ✅ Same chemical properties, different physical properties
• ❌ "Ionic compounds conduct electricity when solid" → ✅ Only when molten or dissolved
• ❌ "Covalent bonds are weak" → ✅ Covalent bonds are strong; intermolecular forces are weak
• ❌ "Graphite is hard like diamond" → ✅ Graphite is soft and slippery due to layered structure
• ❌ "Atoms expand when heated" → ✅ Particles gain energy and move more, they don't expand
• ❌ "Isotopes have different chemical properties" → ✅ Same chemical properties, different physical properties
• ❌ "Ionic compounds conduct electricity when solid" → ✅ Only when molten or dissolved
• ❌ "Covalent bonds are weak" → ✅ Covalent bonds are strong; intermolecular forces are weak
• ❌ "Graphite is hard like diamond" → ✅ Graphite is soft and slippery due to layered structure
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