Inorganic Chemistry
Welcome to Module 4 of the chemistry course (Module 5, Experiments & Applications, follows it). Inorganic chemistry covers bonding (how atoms stick together), states of matter (solids, liquids, gases and how they change), separation techniques (how to get pure substances), chemical tests (how to identify unknown substances), and quantitative chemistry (calculating masses, volumes and concentrations), as well as redox, electrolysis and electrochemical cells, thermochemistry, acids, bases and salts, and the elements and their compounds (Na, Al, Fe, Cl, S, N, Si). These topics tie together everything from Modules 1–3. Bonding and mole calculations underpin much of the rest of the syllabus. By the end of this module you will be able to explain why diamond is hard but graphite is slippery, calculate the mass of product from any reaction, identify mystery gases with simple bench tests, and tackle titration calculations with confidence.
What the Exam Tests
This module spans three syllabus modules and touches the fourth. The mole, concentration and gas calculations, states of matter, and Acids & Bases with pH are Basic Concepts and Calculations. Bonding and structure are Chemical Theories and Laws (chemical bonds and intermolecular forces). Separation Techniques and the gas tests belong to the fourth module, Chemical Experiments and Applications, which the Experiments & Applications module covers in full. Salts & Solubility, Redox Reactions and Elements & Their Compounds are Properties and Reactions of Substances: common elements, oxides, acids, bases and salts, redox reactions, and ionic reactions and tests. Electrochemistry and Thermochemistry go with Chemical Theories and Laws: the syllabus names the theories of electrolyte solutions but not cells or enthalpy, which build on its redox and equilibrium topics.
Questions can ask you to calculate a pH, judge an ionic equation, identify the oxidising agent, recall the colour change of a test, predict the products at each electrode, or work out an enthalpy change.
In our mock papers the three syllabus areas this module feeds, Basic Concepts and Calculations, Properties and Reactions of Substances, and Chemical Theories and Laws, get 12, 12 and 14 of the 48 questions; those counts are shared with the other chemistry modules, and Experiments and Applications gets the other 10. The CSCA publishes the topics but not how many questions each gets, so this split is our own reading of the syllabus. Standard electrode potentials, pKa calculations and the Born–Haber cycle go beyond the syllabus: they are marked in practice and left out of mocks.
Covalent bonding occurs when two non-metal atoms share one or more pairs of electrons so that each atom achieves a full outer shell. The shared pair of electrons is attracted to both nuclei, holding the atoms together.
Single, Double & Triple Bonds
A single bond shares one pair (2 electrons) — for example H–H in hydrogen (H₂). A double bond shares two pairs (4 electrons) — for example O=O in oxygen (O₂). A triple bond shares three pairs (6 electrons) — for example N≡N in nitrogen (N₂). More shared pairs = shorter, stronger bond.
Dot-and-Cross Diagrams
To draw a dot-and-cross diagram: (1) Work out how many outer-shell electrons each atom has. (2) Put the atoms next to each other. (3) Share electrons in the overlap region so every atom gets a full outer shell (2 for hydrogen, 8 for everything else). Use dots for one atom's electrons and crosses for the other's.
In Chinese papers this is the 电子式 (electron formula), drawn with · or × for the outer electrons; every lone pair must be shown (e.g. the 电子式 of N₂ shows three shared pairs plus one lone pair on each N; N≡N is its 结构式). A covalent bond = 共价键; a bond between identical atoms is non-polar (非极性键), between different atoms polar (极性键).
Key Examples
| Molecule | Shared pairs | Bond type | Shape |
|---|---|---|---|
| H₂ | 1 | Single | Linear |
| H₂O | 2 (+ 2 lone pairs on O) | Single ×2 | Bent |
| CO₂ | 4 (2 double bonds) | Double ×2 | Linear |
| CH₄ | 4 (single ×4) | Single ×4 | Tetrahedral |
| N₂ | 3 | Triple | Linear |
Properties of Simple Covalent Molecules
Simple covalent substances have low melting/boiling points because the intermolecular forces (forces between molecules) are weak — even though the covalent bonds within each molecule are strong. They do not conduct electricity because there are no free ions or delocalised electrons.
Intermolecular Forces vs Covalent Bonds
This distinction trips many students up. When ice melts, you are not breaking O–H bonds — you are overcoming the weak forces between H₂O molecules. The covalent bonds inside each water molecule stay intact. Water's boiling point (100 °C) is far below what breaking O–H bonds would need, yet it is unusually high for so small a molecule (H₂S boils at about −60 °C), because water molecules are held by hydrogen bonds (氢键), which are much stronger than ordinary van der Waals forces (范德华力). Hydrogen bonds form between H on N, O or F and a lone pair on another N, O or F; they explain the high boiling points of H₂O, NH₃ and HF and why ice is less dense than water.
Key takeaway: Covalent bonds form when non-metals share electron pairs. Simple covalent molecules have low melting points because weak intermolecular forces (not the strong covalent bonds) break when the substance melts or boils.
Key Formulas
Single bond = 1 shared pair (2 e⁻) | Double bond = 2 shared pairs (4 e⁻) | Triple bond = 3 shared pairs (6 e⁻) | Full shell: H = 2, others = 8 | Simple covalent → low m.p./b.p., no conduction
Worked Example 1
Example 1: Draw the dot-and-cross diagram for water (H₂O).
Step 1: Oxygen has 6 outer electrons; hydrogen has 1 outer electron each.
Step 2: O needs 2 more electrons → shares 1 electron with each H.
Step 3: Each H has 2 electrons (full), O has 8 electrons (full). O also has 2 lone pairs not involved in bonding.
Answer: The diagram shows O in the centre with 2 shared pairs (one with each H) and 2 lone pairs.
Worked Example 2
Example 2: Explain why methane (CH₄) has a low boiling point (−162 °C).
Step 1: CH₄ is a simple covalent molecule.
Step 2: The intermolecular forces between CH₄ molecules are very weak.
Step 3: Only a small amount of energy is needed to overcome these weak forces, so the boiling point is low.
Key: Do NOT say "the covalent bonds are weak" — the C–H bonds are strong; it's the forces between molecules that are weak.
Worked Example 3
Example 3: Why does CO₂ not conduct electricity?
Step 1: CO₂ is a simple covalent molecule.
Step 2: It has no free electrons or ions.
Step 3: Therefore it cannot carry an electric charge, so it does not conduct.
Exam tip: When asked "why does [molecule] have a low boiling point?", always mention: (1) simple molecular structure, (2) weak intermolecular forces, (3) little energy needed to overcome them. Never say "weak covalent bonds"!
Exam tip: Dot-and-cross diagrams: count total outer electrons first. For almost every molecule you will draw the total is even (NO and NO₂, with 11 and 17, are the odd-electron exceptions); an odd total usually means a counting error. Show every lone pair, on the central atom and on outer atoms such as Cl in CCl₄: a 电子式 that leaves any out is wrong.
Common mistake: Saying "the bonds are weak": The covalent bonds WITHIN molecules are strong. It's the INTERMOLECULAR forces between molecules that are weak. Keep the two apart whenever melting or boiling is explained.
Common mistake: Forgetting lone pairs: In H₂O, oxygen has 2 lone pairs. In NH₃, nitrogen has 1 lone pair. Always show them.
Try This Yourself
Nitrogen (N₂) has a boiling point of −196 °C. Explain why, in terms of bonding and structure.
Module Summary
Open and read all sections to complete this module