Topic 2: Atoms, elements and compounds

📝 Study Notes

<div class="main-heading">📌 Topic 2: Atoms, Elements and Compounds</div>

<div class="sub-heading">2.1 Elements, compounds and mixtures</div>

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<span><span class="point-number">1</span> Describe the differences between elements, compounds and mixtures</span>
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<tr style="background:#1E3A8A; color:white;"><th>Type</th><th>Definition</th><th>Example</th><th>Separation</th></tr>
<tr><td><strong>Element</strong></td><td>Pure substance made of only one type of atom</td><td>O₂, Fe, Au</td><td>Cannot be broken down</td></tr>
<tr style="background:#F9FAFB;"><td><strong>Compound</strong></td><td>Two or more elements chemically bonded</td><td>H₂O, CO₂, NaCl</td><td>Only by chemical reaction</td></tr>
<tr><td><strong>Mixture</strong></td><td>Substances physically combined, not bonded</td><td>Air, salt water</td><td>Physical methods (filtration, evaporation)</td></tr>
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<div class="key-definition">💡 <strong>Key Point:</strong> Elements cannot be split into simpler substances. Compounds have fixed ratios. Mixtures have variable composition.</div>
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<div class="sub-heading">2.2 Atomic structure and the Periodic Table</div>

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<span><span class="point-number">1</span> Describe the structure of the atom as a central nucleus containing neutrons and protons surrounded by electrons in shells</span>
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<p><strong>Atomic Structure:</strong></p>
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<li><strong>Nucleus</strong> (centre): Contains <strong>protons</strong> (+) and <strong>neutrons</strong> (neutral)</li>
<li><strong>Electron shells</strong>: Electrons (-) orbit the nucleus in energy levels (shells)</li>
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Electron shells (energy levels)
/ \
/ \
| ⚫ | ← Nucleus (protons + neutrons)
\ /
\ /
-----</pre>
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<span><span class="point-number">2</span> State the relative charges and relative masses of a proton, a neutron and an electron</span>
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<tr style="background:#1E3A8A; color:white;"><th>Particle</th><th>Relative Charge</th><th>Relative Mass</th><th>Location</th></tr>
<tr><td>Proton</td><td>+1</td><td>1</td><td>Nucleus</td></tr>
<tr style="background:#F9FAFB;"><td>Neutron</td><td>0</td><td>1</td><td>Nucleus</td></tr>
<tr><td>Electron</td><td>-1</td><td>1/1840 (negligible)</td><td>Electron shells</td></tr>
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<span><span class="point-number">3</span> Define proton number/atomic number as the number of protons in the nucleus of an atom</span>
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<p><strong>Atomic Number (Z)</strong> = Number of protons in the nucleus</p>
<p>In a neutral atom, the number of protons = number of electrons</p>
<div class="key-definition">💡 Example: Carbon has atomic number 6 → 6 protons, 6 electrons</div>
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<span><span class="point-number">4</span> Define mass number/nucleon number as the total number of protons and neutrons in the nucleus of an atom</span>
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<p><strong>Mass Number (A)</strong> = Number of protons + number of neutrons</p>
<p><strong>Number of neutrons</strong> = Mass number - Atomic number</p>
<div class="key-definition">💡 Example: Carbon-12 → 12 - 6 = 6 neutrons</div>
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<span><span class="point-number">5</span> Determine the electronic configuration of elements and their ions with proton number 1 to 20</span>
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<tr style="background:#1E3A8A; color:white;"><th>Element</th><th>Atomic Number</th><th>Electronic Configuration</th></tr>
<tr><td>Hydrogen</td><td>1</td><td>1</td></tr>
<tr style="background:#F9FAFB;"><td>Helium</td><td>2</td><td>2</td></tr>
<tr><td>Lithium</td><td>3</td><td>2,1</td></tr>
<tr style="background:#F9FAFB;"><td>Beryllium</td><td>4</td><td>2,2</td></tr>
<tr><td>Boron</td><td>5</td><td>2,3</td></tr>
<tr style="background:#F9FAFB;"><td>Carbon</td><td>6</td><td>2,4</td></tr>
<tr><td>Nitrogen</td><td>7</td><td>2,5</td></tr>
<tr style="background:#F9FAFB;"><td>Oxygen</td><td>8</td><td>2,6</td></tr>
<tr><td>Fluorine</td><td>9</td><td>2,7</td></tr>
<tr style="background:#F9FAFB;"><td>Neon</td><td>10</td><td>2,8</td></tr>
<tr><td>Sodium</td><td>11</td><td>2,8,1</td></tr>
<tr style="background:#F9FAFB;"><td>Magnesium</td><td>12</td><td>2,8,2</td></tr>
<tr><td>Aluminium</td><td>13</td><td>2,8,3</td></tr>
<tr style="background:#F9FAFB;"><td>Silicon</td><td>14</td><td>2,8,4</td></tr>
<tr><td>Phosphorus</td><td>15</td><td>2,8,5</td></tr>
<tr style="background:#F9FAFB;"><td>Sulfur</td><td>16</td><td>2,8,6</td></tr>
<tr><td>Chlorine</td><td>17</td><td>2,8,7</td></tr>
<tr style="background:#F9FAFB;"><td>Argon</td><td>18</td><td>2,8,8</td></tr>
<tr><td>Potassium</td><td>19</td><td>2,8,8,1</td></tr>
<tr style="background:#F9FAFB;"><td>Calcium</td><td>20</td><td>2,8,8,2</td></tr>
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<div class="key-definition">💡 <strong>Shell Rules:</strong> 1st shell = max 2 electrons, 2nd shell = max 8 electrons, 3rd shell = max 8 electrons (simplified model)</div>
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<span><span class="point-number">6</span> 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</span>
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<p><strong>Relationship to Periodic Table:</strong></p>
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<li><strong>Period number</strong> = Number of occupied electron shells</li>
<li><strong>Group number (I-VII)</strong> = Number of electrons in the outer shell</li>
<li><strong>Group VIII (Noble gases)</strong> = Full outer shell (stable, unreactive)</li>
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<div class="key-definition">💡 Example: Sodium (Period 3, Group 1) → 3 shells, 1 outer electron → 2,8,1</div>
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<div class="sub-heading">2.3 Isotopes</div>

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<span><span class="point-number">1</span> Define isotopes as different atoms of the same element that have the same number of protons but different numbers of neutrons</span>
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<p><strong>Definition:</strong> Isotopes are atoms of the same element with the <strong>same number of protons</strong> but <strong>different numbers of neutrons</strong>.</p>
<p><strong>Example: Chlorine</strong></p>
<ul><li>Chlorine-35: 17 protons, 18 neutrons</li><li>Chlorine-37: 17 protons, 20 neutrons</li></ul>
<div class="key-definition">💡 Isotopes have the same chemical properties (same electrons) but different physical properties (different mass).</div>
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<span><span class="point-number">2</span> Interpret and use symbols for atoms and ions</span>
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<p><strong>Atomic Notation:</strong> ¹²₆C</p>
<ul><li>Top number (12) = Mass number (protons + neutrons)</li><li>Bottom number (6) = Atomic number (protons)</li></ul>
<div class="key-definition">💡 Example: ³⁵₁₇Cl⁻ means a chloride ion with 17 protons, 18 neutrons, 18 electrons</div>
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<span><span class="point-number">3</span> State that isotopes have the same chemical properties because they have the same number of electrons <span class="supplement-badge">Supplement</span></span>
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<p><strong>Why Same Chemical Properties?</strong> Chemical properties depend on <strong>electron configuration</strong>. All isotopes of an element have the same number of electrons, so they react the same way chemically.</p>
<p><strong>Why Different Physical Properties?</strong> Different mass affects physical properties like density, boiling point, and rate of diffusion.</p>
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<span><span class="point-number">4</span> Calculate the relative atomic mass from isotopic abundances <span class="supplement-badge">Supplement</span></span>
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<p><strong>Formula:</strong> Aᵣ = ( % of isotope A × mass A ) + ( % of isotope B × mass B ) / 100</p>
<p><strong>Worked Example: Chlorine</strong></p>
<ul><li>Chlorine-35: 75% abundance</li><li>Chlorine-37: 25% abundance</li></ul>
<p>Aᵣ = (75 × 35) + (25 × 37) / 100 = (2625 + 925) / 100 = 3550 / 100 = <strong>35.5</strong></p>
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<div class="sub-heading">2.4 Ions and ionic bonds</div>

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<span><span class="point-number">1</span> Describe the formation of positive ions (cations) and negative ions (anions)</span>
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<p><strong>Cations (positive ions):</strong> Metals <strong>lose electrons</strong> to achieve a full outer shell</p>
<ul><li>Example: Na → Na⁺ + e⁻ (sodium loses 1 electron)</li></ul>
<p><strong>Anions (negative ions):</strong> Non-metals <strong>gain electrons</strong> to achieve a full outer shell</p>
<ul><li>Example: Cl + e⁻ → Cl⁻ (chlorine gains 1 electron)</li></ul>
<div class="key-definition">💡 Group 1 metals form 1⁺ ions, Group 2 form 2⁺, Group 6 form 2⁻, Group 7 form 1⁻</div>
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<span><span class="point-number">2</span> State that an ionic bond is a strong electrostatic attraction between oppositely charged ions</span>
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<p>An ionic bond is the <strong>strong electrostatic attraction</strong> between oppositely charged ions.</p>
<p>Forms when a metal reacts with a non-metal.</p>
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<span><span class="point-number">3</span> Describe the formation of ionic bonds between Group I and Group VII using dot-and-cross diagrams</span>
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<p><strong>Sodium Chloride (NaCl)</strong></p>
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Na atom: 2,8,1 Cl atom: 2,8,7
↓ ↓
Na⁺ ion: 2,8 Cl⁻ ion: 2,8,8

[Na]⁺ [ Cl ]⁻
••
••</pre>
<div class="key-definition">💡 Group I metals lose 1 electron → +1 charge. Group VII non-metals gain 1 electron → -1 charge.</div>
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<span><span class="point-number">4</span> Describe the properties of ionic compounds: high melting/boiling points; good conductivity when aqueous or molten, poor when solid</span>
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<tr style="background:#1E3A8A; color:white;"><th>Property</th><th>Explanation</th></tr>
<tr><td>High melting/boiling points</td><td>Strong electrostatic forces between ions require lots of energy to overcome</td></tr>
<tr style="background:#F9FAFB;"><td>Conducts when molten/dissolved</td><td>Ions are free to move and carry charge</td></tr>
<tr><td>Poor conductor when solid</td><td>Ions are fixed in position in the lattice</td></tr>
<tr style="background:#F9FAFB;"><td>Hard but brittle</td><td>Layers shift, like charges align and repel</td></tr>
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<span><span class="point-number">5</span> Describe the giant lattice structure of ionic compounds <span class="supplement-badge">Supplement</span></span>
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<p><strong>Giant Lattice Structure:</strong> Ions are arranged in a regular, repeating 3D pattern with alternating positive and negative ions.</p>
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+ - + -
- + - +
+ - + -
- + - +
(Alternating positive and negative ions)</pre>
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<span><span class="point-number">6</span> Explain the properties of ionic compounds in terms of structure and bonding <span class="supplement-badge">Supplement</span></span>
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<li><strong>High melting/boiling points:</strong> Strong electrostatic forces between oppositely charged ions in all directions require large amounts of energy to overcome.</li>
<li><strong>Conductivity when molten/dissolved:</strong> Ions become free to move and can carry electric charge.</li>
<li><strong>No conductivity when solid:</strong> Ions are fixed in lattice positions and cannot move.</li>
<li><strong>Brittleness:</strong> When layers shift, like charges align and repel, causing the crystal to split.</li>
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<div class="sub-heading">2.5 Simple molecules and covalent bonds</div>

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<span><span class="point-number">1</span> State that a covalent bond is formed when a pair of electrons is shared between two atoms</span>
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<p><strong>Covalent Bond:</strong> Formed when a <strong>pair of electrons is shared</strong> between two non-metal atoms.</p>
<p>Each atom contributes one electron to the shared pair, and both atoms achieve a full outer shell (noble gas configuration).</p>
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<span><span class="point-number">2</span> Describe the formation of covalent bonds in simple molecules using dot-and-cross diagrams</span>
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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)</pre>
<div class="key-definition">💡 Only outer shell electrons are shown in dot-and-cross diagrams.</div>
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<span><span class="point-number">3</span> Describe the properties of simple molecular compounds: low melting/boiling points; poor electrical conductivity</span>
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<tr style="background:#1E3A8A; color:white;"><th>Property</th><th>Explanation</th></tr>
<tr><td>Low melting/boiling points</td><td>Weak intermolecular forces between molecules (not the strong covalent bonds inside molecules)</td></tr>
<tr style="background:#F9FAFB;"><td>Poor electrical conductivity</td><td>No free electrons or ions to carry charge</td></tr>
<tr><td>Often gases or liquids at room temperature</td><td>Weak forces mean they evaporate easily</td></tr>
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<span><span class="point-number">4</span> Describe the formation of covalent bonds in additional molecules (O₂, N₂, CO₂, C₂H₄, CH₃OH) <span class="supplement-badge">Supplement</span></span>
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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
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H—C—O—H
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H</pre>
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<span><span class="point-number">5</span> Explain the properties of simple molecular compounds in terms of structure and bonding <span class="supplement-badge">Supplement</span></span>
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<p><strong>Key Concept:</strong></p>
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<li><strong>Intramolecular forces</strong> (covalent bonds) = strong, hold atoms together within a molecule</li>
<li><strong>Intermolecular forces</strong> = weak, hold molecules together</li>
<li>Low melting/boiling points because only weak intermolecular forces need to be overcome</li>
<li>Poor conductivity because no free electrons or ions are present</li>
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<div class="sub-heading">2.6 Giant covalent structures</div>

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<span><span class="point-number">1</span> Describe the giant covalent structures of graphite and diamond</span>
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<p><strong>Diamond:</strong> Each carbon atom bonded to 4 others (tetrahedral) - all strong covalent bonds</p>
<p><strong>Graphite:</strong> 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.</p>
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<span><span class="point-number">2</span> Relate the structures and bonding of graphite and diamond to their uses</span>
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<tr style="background:#1E3A8A; color:white;"><th>Property</th><th>Diamond</th><th>Graphite</th></tr>
<tr><td>Bonding</td><td>4 covalent bonds per carbon</td><td>3 covalent bonds per carbon</td></tr>
<tr style="background:#F9FAFB;"><td>Structure</td><td>3D tetrahedral</td><td>Layered hexagonal</td></tr>
<tr><td>Hardness</td><td>Very hard (hardest natural substance)</td><td>Soft, slippery</td></tr>
<tr style="background:#F9FAFB;"><td>Conductivity</td><td>Non-conductor (no free electrons)</td><td>Conductor (delocalised electrons)</td></tr>
<tr><td>Uses</td><td>Cutting tools, drills, jewellery</td><td>Pencil lead, lubricant, electrodes</td></tr>
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<div class="sub-heading">2.7 Metallic bonding</div>

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<span><span class="point-number">1</span> Describe metallic bonding as the electrostatic attraction between positive ions and a 'sea' of delocalised electrons <span class="supplement-badge">Supplement</span></span>
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<p><strong>Metallic Bonding:</strong></p>
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<li>Metal atoms lose their outer electrons to become positive ions</li>
<li>These electrons become <strong>delocalised</strong> (free to move) forming a "sea of electrons"</li>
<li>The metallic bond is the <strong>electrostatic attraction</strong> between positive metal ions and the delocalised electrons</li>
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Metal lattice: + + + +
+ + + + (with sea of electrons around)
+ + + +</pre>
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<span><span class="point-number">2</span> Explain the properties of metals: good electrical conductivity; malleability and ductility <span class="supplement-badge">Supplement</span></span>
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<tr style="background:#1E3A8A; color:white;"><th>Property</th><th>Explanation</th></tr>
<tr><td>Good electrical conductivity</td><td>Delocalised electrons are free to move and carry charge</td></tr>
<tr style="background:#F9FAFB;"><td>Good thermal conductivity</td><td>Delocalised electrons transfer kinetic energy quickly</td></tr>
<tr><td>High melting/boiling points</td><td>Strong electrostatic forces between ions and electrons require lots of energy to overcome</td></tr>
<tr style="background:#F9FAFB;"><td>Malleable (can be hammered into shape)</td><td>Layers of positive ions can slide over each other without breaking the metallic bond</td></tr>
<tr><td>Ductile (can be drawn into wires)</td><td>Same reason as malleability - layers can slide</td></tr>
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<div class="warning-box">
<strong>⚠️ Common Mistakes to Avoid:</strong><br><br>
• ❌ "Atoms expand when heated" → ✅ Particles gain energy and move more, they don't expand<br>
• ❌ "Isotopes have different chemical properties" → ✅ Same chemical properties, different physical properties<br>
• ❌ "Ionic compounds conduct electricity when solid" → ✅ Only when molten or dissolved<br>
• ❌ "Covalent bonds are weak" → ✅ Covalent bonds are strong; intermolecular forces are weak<br>
• ❌ "Graphite is hard like diamond" → ✅ Graphite is soft and slippery due to layered structure
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2.1 Atomic structure and the Periodic Table

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2.2 Ions and ionic bonds

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2.3 Molecules and covalent bonds

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2.4 Metallic bonds

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2.5 Giant covalent structures

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