Microscopic World I (Part 2): Covalent Bonding and Molecular Structures
FoundationA covalent bond forms when non-metal atoms share electron pairs. Simple molecular substances melt easily because only weak van der Waals' forces hold molecules together, while giant covalent substances such as diamond and silicon dioxide must break strong covalent bonds to melt — very hard with extremely high melting points.
- Covalent bond = non-metal atoms sharing electron pairs so outer shells reach an octet
- Electron diagrams: H₂ single bond, O₂ double, N₂ triple, H₂O two singles plus two lone pairs, CH₄ four singles, CO₂ two doubles
- Simple molecular substances: strong bonds within molecules, weak van der Waals' forces between them — low melting points, non-conducting
文體小知識：科學知識說明文——以電子圖講解共價鍵如何由共用電子對形成，並比較簡單分子物質與巨型共價物質的結構和性質差異。
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1Ionic bonding joins a metal to a non-metal by electron transfer. When two non-metal atoms meet, neither will give up electrons, so they share instead: each atom contributes one electron to a shared pair, attracted by both nuclei, which holds the atoms together. This is a covalent bond. After sharing, the outermost shell of each atom usually reaches an octet, as stable as a noble gas.
2Electron diagrams show covalent bonds clearly: in a hydrogen molecule one shared pair forms a single bond; in O₂ two oxygen atoms share two pairs, forming a double bond; in N₂ three pairs form a triple bond. In H₂O the oxygen bonds to each hydrogen with one pair; in CH₄ carbon forms four single bonds; in CO₂ carbon forms a double bond with each of the two oxygens.
3Carbon dioxide and iodine consist of separate molecules — simple molecular substances. The covalent bonds inside each molecule are strong, but between molecules there are only weak van der Waals' forces, so a little heat pulls the molecules apart. Such substances have low melting and boiling points, are often gases or volatile liquids, and do not conduct electricity when molten, since there are no mobile charges anywhere.
4Diamond, graphite and quartz (silicon dioxide) are giant covalent substances: countless atoms are joined into one whole network by covalent bonds. Melting requires breaking many strong bonds, so they have extremely high melting points and are very hard. In diamond each carbon is bonded to four others, making it one of the hardest natural materials; quartz is hard and heat-resistant, used for heatproof ware. The properties of a substance follow from its structure and bonding.
1Summary 重點摘要
A covalent bond forms when non-metal atoms share electron pairs. Simple molecular substances melt easily because only weak van der Waals' forces hold molecules together, while giant covalent substances such as diamond and silicon dioxide must break strong covalent bonds to melt — very hard with extremely high melting points.
2Vocabulary cards Vocabulary
A chemical bond formed when two non-metal atoms share a pair of electrons.
A pair formed by one electron from each atom, attracted to both nuclei.
A diagram using dots and crosses for outer-shell electrons to show bonding in a molecule.
A covalent bond formed by one shared pair, e.g. the O—H bonds in water.
A covalent bond of two shared pairs, as in O₂ and CO₂.
A stable outermost shell of eight electrons, like a noble gas.
A substance made of separate molecules held together only by van der Waals' forces, e.g. CO₂ and iodine.
Weak attractions between molecules, far weaker than covalent bonds, responsible for low melting points.
A structure in which countless atoms are joined by covalent bonds throughout, e.g. diamond and quartz.
The mineral form of silicon dioxide, a giant covalent solid, hard and heat-resistant.
3Step by step Steps
- Step 1: How non-metal atoms shake handsTwo non-metal atoms both need more electrons to complete their outer shells, and neither will hand electrons over. The only way is to share: each contributes one electron. For example, two hydrogen atoms each have one electron; after sharing, each hydrogen effectively has two, stable like helium, and a hydrogen molecule forms. The shared pair is attracted to both positive nuclei at once, like two hands gripping the atoms together — this shared pair is the covalent bond.
- Step 2: Drawing single bonds with electron diagramsElectron diagrams use dots and crosses for the outer-shell electrons of different atoms; start by counting the valence electrons. Take water: oxygen has six outer electrons, each hydrogen one; oxygen shares one pair with each hydrogen, giving two single bonds, while oxygen keeps two lone pairs that do not bond. Take methane: carbon has four valence electrons and forms four single bonds, completing its octet. Always draw each shared pair between the two atoms, and never leave out the lone pairs.
- Step 3: Double and triple bondsSome atoms must share several pairs to fill their outer shells. An oxygen atom has six valence electrons, so two oxygen atoms share two pairs — a double bond — forming O₂. In carbon dioxide, carbon shares two pairs with each oxygen, two double bonds, and all three atoms reach an octet. A nitrogen atom has five valence electrons, three short, so two nitrogens share three pairs in a triple bond; this makes the N₂ bond very strong and nitrogen gas chemically unreactive.
- Step 4: Why simple molecular substances melt easilyCarbon dioxide, iodine and dry ice consist of separate molecules. The covalent bonds within each molecule are strong, but between molecules there are only weak van der Waals' forces. Melting does not break the molecules apart; it merely pulls them further from one another, so only the weak intermolecular attractions must be overcome. Picture a pack of cards: a single card is hard to tear (strong covalent bond) but the whole pile scatters with a light push (weak van der Waals' forces). Hence low melting points — CO₂ is a gas at room temperature.
- Step 5: The steel mesh of giant covalent substancesDiamond, graphite and quartz belong to another class: there are no separate molecules, and countless atoms are linked by covalent bonds into one vast network — a giant covalent structure. In diamond each carbon bonds to four neighbours with single bonds extending in all directions; in quartz each silicon is joined to four oxygens, equally rigid. Melting such a substance requires breaking a huge number of strong covalent bonds, so melting points are extremely high and the solids very hard. For example, a tiny diamond chip on a glass cutter scores glass effortlessly.
- Step 6: CO₂ a gas, SiO₂ a solidCarbon dioxide and silicon dioxide both pair carbon or silicon with oxygen, yet their properties differ enormously — the key is structure. CO₂ is a simple molecular substance: each molecule stands alone, held to others only by van der Waals' forces, so it vaporises with a little heat and is a gas at room temperature. SiO₂ is giant covalent: Si—O covalent bonds run through the whole crystal, and melting means breaking these strong bonds, so it melts above 1500°C and is a hard solid. For example, soft drinks dissolve CO₂ gas, while sand and quartz are chiefly SiO₂.
- Step 7: Why simple molecular substances do not conductConduction needs mobile charges, and simple molecular substances have neither kind: electrons are locked into covalent bonds and cannot flow, and the molecules do not split into ions. So they do not conduct as solids, when molten, or usually in solution. For example, distilled water and sugar solution conduct no electricity, in contrast to salt water — salt is an ionic compound with free ions once dissolved. In answers, write the full phrase 'no mobile charges present'; merely writing 'no ions' is incomplete and loses marks.
4Diagrams Diagrams
5Worked example Example
Data question: three students draw electron diagrams for water, carbon dioxide and methane, then discuss why carbon dioxide is a gas at room temperature while quartz (silicon dioxide), of similar composition, is a hard solid. State the bond types in the three molecules and explain the difference in properties.
- Count valence electrons: H has one, O six, C four. In H₂O oxygen shares one pair with each hydrogen — two single bonds — and keeps two lone pairs
- In CO₂ carbon shares two pairs with each oxygen: two double bonds, four shared pairs in all, and all atoms reach an octet
- In CH₄ carbon forms four single C—H bonds; all three molecules satisfy outer shells by sharing electron pairs
- CO₂ is a simple molecular substance: melting or boiling only overcomes weak van der Waals' forces between molecules, so it is a gas at room temperature
- SiO₂ is giant covalent: Si—O bonds run through the whole crystal, so melting requires breaking many strong bonds — extremely high melting point and hardness
- Answer: Water contains two O—H single bonds (with two lone pairs on oxygen); carbon dioxide contains two C=O double bonds, four shared pairs in total; methane contains four C—H single bonds. Although both CO₂ and SiO₂ contain covalent bonds, their structures differ: CO₂ consists of separate molecules held only by weak van der Waals' forces, so a little heat separates the molecules and it is a gas at room temperature. SiO₂ is a giant covalent network whose bonds run throughout the crystal; melting must break many strong covalent bonds, so it has an extremely high melting point and is a hard solid.
6Common mistakes Mistakes
7Quick check Quick check
Q1What is a covalent bond?
Q2How many shared pairs of electrons are there in one CO₂ molecule?
Q3Why is carbon dioxide a gas at room temperature while silicon dioxide is a high-melting solid?
Q4Name two giant covalent substances.
Q5Why do simple molecular substances not conduct electricity?
8Recap Recap
- Covalent bond = non-metal atoms sharing electron pairs so outer shells reach an octet
- Electron diagrams: H₂ single bond, O₂ double, N₂ triple, H₂O two singles plus two lone pairs, CH₄ four singles, CO₂ two doubles
- Simple molecular substances: strong bonds within molecules, weak van der Waals' forces between them — low melting points, non-conducting
- Giant covalent substances (diamond, graphite, quartz): extremely hard, extremely high melting points — melting breaks covalent bonds
- CO₂ a gas but SiO₂ a hard solid: the whole difference lies between simple molecular and giant covalent structures