Chemical Bonding¶
Syllabus mapping
CH-11-01 — Properties and Structure of Matter. Covers: IUPAC naming of inorganic substances; metallic bonding as electrostatic attraction, modelled as a lattice, and its link to melting point, boiling point and conductivity; using electronegativity to decide whether a bond is ionic or covalent; predicting ions from the periodic table; the arrangement of ions in ionic lattices; a secondary-source investigation into melting point trends in group 1 and group 2 chlorides; covalent bonding in discrete molecules and covalent networks; electron dot diagrams for molecules and simple polyatomic ions; bond polarity; predicting molecular shape and polarity with VSEPR theory; dispersion forces, dipole–dipole forces and hydrogen bonding; comparing the strength of intermolecular forces with ionic, covalent and metallic bonds; explaining physical properties from bonding type (including a practical investigation); evaluating models of ionic, covalent and metallic substances; a secondary-source investigation of carbon allotropes.
Wording above is paraphrased for teaching use — check the official syllabus before treating any line as verbatim NESA text.
You need to know: how to name and write formulas for simple compounds, the three types of chemical bond (metallic, ionic, covalent) and how to tell which one is present, how to draw electron dot diagrams and predict molecular shapes, the three types of intermolecular force, and how bonding explains a substance's physical properties.
This is the biggest page in Year 11 Chemistry. It's split into four parts, all of them Core content — take them one at a time:
- Naming compounds
- The three types of bond
- Molecules: dot diagrams, shapes and polarity
- Intermolecular forces and properties
Assumed knowledge check¶
Before this page makes sense, you should be comfortable with:
- Working out an ion's charge from the periodic table
- Electronegativity and how it changes across the periodic table
- Counting valence (outer) electrons, ions and the three bond types from Year 10
Part 1: Naming compounds and writing formulas¶
Ionic compounds (metal + non-metal)¶
- Name the metal first, unchanged.
- Name the non-metal second, with its ending changed to -ide (chlorine → chloride, oxygen → oxide, sulfur → sulfide, nitrogen → nitride).
So \(\ce{NaCl}\) is sodium chloride and \(\ce{MgO}\) is magnesium oxide.
Writing the formula: the positive and negative charges must cancel out to zero.
Worked example
Write the formula for (a) magnesium chloride and (b) aluminium oxide.
(a) Magnesium forms \(\ce{Mg^{2+}}\); chlorine forms \(\ce{Cl-}\). You need two \(\ce{Cl-}\) to cancel one \(\ce{Mg^{2+}}\): \(+2 + 2(-1) = 0\). Formula: \(\ce{MgCl2}\).
(b) Aluminium forms \(\ce{Al^{3+}}\); oxygen forms \(\ce{O^{2-}}\). The smallest numbers that cancel are 2 aluminium and 3 oxygen: \(2(+3) + 3(-2) = 0\). Formula: \(\ce{Al2O3}\).
Shortcut ("swap and drop"): the number of each charge becomes the subscript of the other ion — then simplify if you can.
Metals with more than one possible charge (mostly transition metals) show the charge in Roman numerals: \(\ce{FeCl2}\) is iron(II) chloride (\(\ce{Fe^{2+}}\)); \(\ce{FeCl3}\) is iron(III) chloride (\(\ce{Fe^{3+}}\)). Copper(I)/copper(II) works the same way.
Polyatomic ions¶
Some ions are made of several atoms bonded together, carrying a charge as a group. Learn these:
| Ion | Formula | Ion | Formula |
|---|---|---|---|
| Hydroxide | \(\ce{OH-}\) | Ammonium | \(\ce{NH4+}\) |
| Nitrate | \(\ce{NO3-}\) | Carbonate | \(\ce{CO3^{2-}}\) |
| Sulfate | \(\ce{SO4^{2-}}\) | Hydrogen carbonate | \(\ce{HCO3-}\) |
| Phosphate | \(\ce{PO4^{3-}}\) | Acetate (ethanoate) | \(\ce{CH3COO-}\) |
They follow the same "charges cancel" rule. If you need more than one polyatomic ion, put it in brackets: calcium hydroxide is \(\ce{Ca(OH)2}\), not CaOH2.
Covalent compounds (two non-metals)¶
Use prefixes to show how many of each atom: mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6). "Mono" is left off the first element.
- \(\ce{CO}\) — carbon monoxide
- \(\ce{CO2}\) — carbon dioxide
- \(\ce{N2O4}\) — dinitrogen tetroxide
- \(\ce{SF6}\) — sulfur hexafluoride
Part 2: The three types of chemical bond¶
All three types of bond are electrostatic attractions — positive charges attracting negative charges. What's different is what is attracting what.
Which bond? Use electronegativity¶
| Atoms involved | What happens to electrons | Bond type |
|---|---|---|
| Metal + metal | Outer electrons are shared by all atoms in a "sea" | Metallic |
| Metal + non-metal | Big difference in electronegativity → electrons are transferred from metal to non-metal | Ionic |
| Non-metal + non-metal | Similar electronegativity → electrons are shared in pairs | Covalent |
A rough guide using the difference in electronegativity between the two atoms: a difference above about 1.8 is mostly ionic; between about 0.4 and 1.8 is polar covalent; below about 0.4 is non-polar covalent. (There's no sharp cut-off — bonding is a sliding scale.)
Metallic bonding¶
Metal atoms let go of their outer electrons. This leaves a lattice (a regular 3D arrangement) of positive metal ions surrounded by a "sea" of delocalised electrons — electrons that aren't stuck to any one atom and can move freely. The metallic bond is the electrostatic attraction between the positive ions and the sea of electrons.
🖼️ Diagram needed — a 2D grid of circles labelled "+" (metal cations) surrounded by small "e⁻" dots scattered between them, with a caption "sea of delocalised electrons". A second panel shows a row of ions sliding past another row while the electron sea stays around them (explains malleability). TODO — placeholder until sourced/drawn.
This model explains metals' properties:
| Property | Why |
|---|---|
| Conduct electricity (solid and liquid) | The delocalised electrons can move through the lattice, carrying charge |
| Conduct heat | Moving electrons pass kinetic energy along quickly |
| Malleable and ductile (bend and stretch without breaking) | Layers of ions can slide over each other — the electron sea moves with them, so the attraction isn't broken |
| Usually high melting and boiling points | The attraction between ions and electrons is strong and acts in all directions, so it takes a lot of energy to break up the lattice |
Melting point and bonding strength. More delocalised electrons per atom, and smaller ions, mean a stronger metallic bond:
| Metal | Electrons per atom in the "sea" | Melting point (°C) |
|---|---|---|
| Sodium | 1 | 98 |
| Magnesium | 2 | 650 |
| Aluminium | 3 | 660 |
Down group 1 (lithium 181 °C, sodium 98 °C, potassium 63 °C), the ions get bigger, so the sea of electrons is further from the centre of each ion. The attraction is weaker and the melting point drops.
Ionic bonding¶
A metal atom transfers its outer electrons to a non-metal atom. Both end up with noble-gas configurations: the metal becomes a positive ion, the non-metal a negative ion. The ionic bond is the electrostatic attraction between oppositely charged ions.
The ions don't pair off. They pack into a huge 3D ionic lattice where every ion is surrounded by ions of the opposite charge. In sodium chloride, each \(\ce{Na+}\) is surrounded by 6 \(\ce{Cl-}\) ions, and each \(\ce{Cl-}\) by 6 \(\ce{Na+}\) ions. That's why the formula \(\ce{NaCl}\) is a ratio (1 : 1), not a molecule.
🖼️ Diagram needed — (1) dot-and-cross diagram of sodium transferring one electron to chlorine, forming Na⁺ and Cl⁻ with full outer shells shown in square brackets; (2) a 3D sodium chloride lattice (alternating small Na⁺ and large Cl⁻ spheres in a cube), with one Na⁺ highlighted to show its 6 Cl⁻ neighbours. TODO — placeholder until sourced/drawn.
| Property | Why |
|---|---|
| High melting and boiling points | Many strong attractions in every direction must be broken |
| Hard but brittle | A sharp knock shifts a layer so that ions of the same charge line up — they repel, and the crystal shatters |
| Don't conduct as a solid | The ions are locked in place |
| Conduct when molten or dissolved in water | The ions are free to move and carry charge |
| Often dissolve in water | Water molecules pull the ions out of the lattice |
Secondary-source investigation: group 1 and group 2 chlorides¶
Melting points (°C):
| Group 1 | LiCl | NaCl | KCl | RbCl | CsCl |
|---|---|---|---|---|---|
| Melting point | 605 | 801 | 770 | 718 | 645 |
| Group 2 | BeCl₂ | MgCl₂ | CaCl₂ | SrCl₂ | BaCl₂ |
|---|---|---|---|---|---|
| Melting point | 399 | 714 | 772 | 874 | 962 |
Your task: graph each group's melting points, describe the trend within each group and between the groups, and try to explain it.
A starting point for the explanation: from NaCl down to CsCl, the metal ion gets bigger, so the positive and negative ions sit further apart. The attraction is weaker, so less energy is needed to melt the lattice — the melting point falls.
But notice LiCl and the whole of group 2 don't fit that simple pattern. The simple "ions further apart → weaker" idea isn't the whole story. The Advanced box explains why — and spotting where a simple model breaks down is exactly the kind of thing an "evaluate" question rewards.
Covalent bonding¶
Two non-metal atoms share a pair of electrons so that each gets a full outer energy level. A covalent bond is the electrostatic attraction between the shared pair of electrons and the nuclei of both atoms.
- A single bond shares one pair (e.g. \(\ce{H-H}\)).
- A double bond shares two pairs (e.g. \(\ce{O=O}\)).
- A triple bond shares three pairs (e.g. N≡N).
Covalent substances come in two very different kinds:
| Discrete (simple) molecules | Covalent network (lattice) | |
|---|---|---|
| Structure | Small, separate molecules — strong covalent bonds inside each molecule, weak forces between molecules | One giant structure — every atom covalently bonded to its neighbours, all the way through |
| Examples | \(\ce{H2O}\), \(\ce{CO2}\), \(\ce{CH4}\), \(\ce{O2}\) | Diamond (C), silicon dioxide (\(\ce{SiO2}\), sand/quartz) |
| Melting point | Low (only the weak forces between molecules need breaking) | Very high (strong covalent bonds need breaking) |
Part 3: Molecules — electron dot diagrams, shape and polarity¶
Drawing electron dot diagrams¶
An electron dot diagram (Lewis diagram) shows the valence electrons only, as dots (or dots and crosses). Shared pairs sit between the atoms; lone pairs (non-bonding pairs) sit on one atom.
Step by step:
- Count the total valence electrons (group 1 → 1, group 15 → 5, group 16 → 6, group 17 → 7, etc.). For a negative ion, add one per negative charge; for a positive ion, subtract one per positive charge.
- Place the atom that needs the most bonds in the centre (never hydrogen). Join each outer atom to it with one shared pair.
- Fill the outer atoms' octets with lone pairs (hydrogen only ever needs 2 electrons).
- Put any leftover electrons on the central atom.
- If the central atom still has fewer than 8, move a lone pair from an outer atom to make a double (or triple) bond.
Worked example — water, \(\ce{H2O}\)
1. Count: O has 6, each H has 1 → \(6 + 1 + 1 = 8\) electrons (4 pairs).
2. Place: O in the centre, one shared pair to each H — that uses 2 pairs.
3–4. Each H now has its 2 electrons. The 2 remaining pairs go on the O as lone pairs.
Check: O has 2 shared pairs + 2 lone pairs = 8 electrons ✓
Result: O in the centre, 2 bonding pairs, 2 lone pairs.
Worked example — carbon dioxide, \(\ce{CO2}\)
1. Count: C has 4, each O has 6 → \(4 + 6 + 6 = 16\) electrons (8 pairs).
2. C in the centre, one pair to each O (2 pairs used).
3. Fill each O with 3 lone pairs (6 pairs used — 8 total). Nothing left over.
5. C only has 4 electrons (2 shared pairs). Move one lone pair from each O into a bond → two double bonds: \(\ce{O=C=O}\).
Check: every atom has 8 ✓
🖼️ Diagram needed — electron dot diagrams in one consistent style for H₂O, NH₃, CH₄, CO₂, N₂, and two polyatomic ions (NH₄⁺ and OH⁻, each in square brackets with the charge outside). TODO — placeholder until sourced/drawn.
Summary of common examples:
| Molecule | Bonding pairs on central atom | Lone pairs on central atom | Structural formula |
|---|---|---|---|
| \(\ce{CH4}\) | 4 | 0 | C with 4 single bonds to H |
| \(\ce{NH3}\) | 3 | 1 | N with 3 single bonds to H |
| \(\ce{H2O}\) | 2 | 2 | H–O–H |
| \(\ce{CO2}\) | 2 double bonds | 0 | O=C=O |
| \(\ce{N2}\) | 1 triple bond | 1 on each N | N≡N |
Polar bonds¶
When two atoms with different electronegativities share electrons, the more electronegative atom pulls the shared pair closer to itself. It gets a small negative charge (\(\delta^-\), "delta negative") and the other atom a small positive charge (\(\delta^+\)). This is a polar covalent bond.
In \(\ce{H-Cl}\), chlorine is more electronegative, so it's \(\delta^-\) and hydrogen is \(\delta^+\). In \(\ce{H-H}\) or \(\ce{Cl-Cl}\), both atoms are the same, so the bond is non-polar.
Molecular shape: VSEPR theory¶
VSEPR (valence shell electron pair repulsion) theory: electron pairs around the central atom repel each other, so they spread out to be as far apart as possible. Lone pairs repel a little more strongly than bonding pairs. A double or triple bond counts as one region.
| Bonding regions | Lone pairs | Shape | Bond angle | Example |
|---|---|---|---|---|
| 2 | 0 | Linear | 180° | \(\ce{CO2}\) |
| 3 | 0 | Trigonal planar | 120° | \(\ce{BF3}\) |
| 4 | 0 | Tetrahedral | 109.5° | \(\ce{CH4}\) |
| 3 | 1 | Trigonal pyramidal | about 107° | \(\ce{NH3}\) |
| 2 | 2 | Bent (V-shaped) | about 104.5° | \(\ce{H2O}\) |
🖼️ Diagram needed — 3D ball-and-stick sketches of the five shapes in the table above, with lone pairs drawn as lobes and bond angles labelled. Wedge/dash notation for the tetrahedral and pyramidal shapes. TODO — placeholder until sourced/drawn.
Is the whole molecule polar?¶
A molecule is polar if it has polar bonds and an uneven (asymmetric) shape, so the charges don't cancel out.
- \(\ce{CO2}\): both C=O bonds are polar, but the molecule is linear and symmetrical, so they pull equally in opposite directions and cancel → non-polar molecule.
- \(\ce{H2O}\): both O–H bonds are polar, and the molecule is bent, so they don't cancel → polar molecule (O end \(\delta^-\), H end \(\delta^+\)).
- \(\ce{CH4}\): symmetrical tetrahedral → non-polar. \(\ce{NH3}\): pyramidal (lone pair on one side) → polar.
A quick test: if the central atom has lone pairs, or the outer atoms aren't all the same, the molecule is usually polar.
Part 4: Intermolecular forces and physical properties¶
Intermolecular forces¶
Covalent bonds hold atoms together inside a molecule. Intermolecular forces are the much weaker attractions between separate molecules. When a molecular substance melts or boils, it's the intermolecular forces that break — not the covalent bonds. (Boiling water makes water vapour, not hydrogen and oxygen.)
| Force | Found between | How it works | Relative strength |
|---|---|---|---|
| Dispersion forces | All molecules (the only force between non-polar molecules) | Electrons move around, so at any instant one side of a molecule may have slightly more electrons. This temporary dipole causes one in the next molecule, and they attract | Weakest for small molecules; stronger for bigger molecules with more electrons |
| Dipole–dipole forces | Polar molecules | The \(\delta^+\) end of one molecule attracts the \(\delta^-\) end of the next | Stronger than dispersion for molecules of similar size |
| Hydrogen bonding | Molecules with H bonded directly to N, O or F | That H is very \(\delta^+\) and is attracted to a lone pair on an N, O or F of a neighbouring molecule | Strongest of the three |
🖼️ Diagram needed — three panels: (1) two non-polar molecules with temporary δ+/δ− labels and a dotted line between them (dispersion); (2) a row of HCl molecules lined up δ+ to δ− (dipole–dipole); (3) several water molecules with dotted hydrogen bonds from H on one molecule to a lone pair on O of the next. Dotted lines for intermolecular forces, solid lines for covalent bonds throughout. TODO — placeholder until sourced/drawn.
Evidence — water is unusual:
| Molecule | \(\ce{H2O}\) | \(\ce{H2S}\) | \(\ce{H2Se}\) | \(\ce{H2Te}\) |
|---|---|---|---|---|
| Boiling point (°C) | 100 | −60 | −41 | −2 |
From \(\ce{H2S}\) to \(\ce{H2Te}\) the molecules get bigger, so dispersion forces get stronger and boiling point rises — as expected. Following that trend, water "should" boil at around −80 °C. It boils at 100 °C because of hydrogen bonding. Without it, there'd be no liquid water on Earth.
Comparing the strengths of all the forces¶
From strongest to weakest (in general):
The chemical bonds (ionic, covalent, metallic) are roughly 10 to 100 times stronger than intermolecular forces. That's the single most important reason molecular substances have low melting points and everything else has high ones.
Predicting physical properties from bonding¶
| Ionic | Metallic | Covalent molecular | Covalent network | |
|---|---|---|---|---|
| Example | \(\ce{NaCl}\) | Copper | Water, sugar, \(\ce{CO2}\) | Diamond, \(\ce{SiO2}\) |
| Melting point | High | Usually high | Low | Very high |
| Conducts as a solid? | No | Yes | No | No (graphite is the exception) |
| Conducts when molten or dissolved? | Yes | Yes (molten) | No | — (doesn't melt or dissolve easily) |
| Hard or soft? | Hard, brittle | Malleable | Usually soft | Very hard |
| Dissolves in water? | Often | No | Polar ones often do; non-polar don't | No |
Worked example
Substance X melts at 1085 °C, conducts electricity as a solid, and can be hammered into a sheet. What type of bonding is present?
- High melting point → not covalent molecular.
- Conducts as a solid → must have free-moving charges in the solid. Ionic solids don't; covalent networks (except graphite) don't.
- Malleable → layers slide without shattering.
Metallic bonding. (It's copper.)
Models and their limits¶
Every diagram on this page is a model — a simplified picture that shows some features well and gets others wrong.
| Model | Shows well | Limitations |
|---|---|---|
| Electron dot diagram | Which electrons are shared or transferred; lone pairs | Flat — doesn't show 3D shape; electrons aren't really fixed dots |
| Ball-and-stick | 3D shape and bond angles | Atoms aren't really separated by gaps; "sticks" make bonds look like physical rods |
| Space-filling | Relative sizes; how the molecule takes up space | Hard to see bond angles and the atoms at the back |
| Ionic lattice diagram | The repeating arrangement of ions | Usually only shows a tiny section; ions drawn too far apart |
| Electron sea model | Why metals conduct and bend | Very simplified — doesn't explain differences between metals well |
"Evaluate how models are used" questions want both sides: what the model is good for and what it misses.
Secondary-source investigation: carbon allotropes¶
Allotropes are different structural forms of the same element. Carbon has several, and they have wildly different properties — because of how the atoms are bonded, not what they're made of.
| Diamond | Graphite | Graphene | Fullerenes (e.g. \(\ce{C60}\)) | |
|---|---|---|---|---|
| Structure | Each C bonded to 4 others in a 3D tetrahedral network | Each C bonded to 3 others in flat hexagonal layers; weak dispersion forces between layers | A single layer of graphite, one atom thick | Separate ball-shaped molecules of 60 C atoms |
| Spare electrons? | No — all 4 used in bonds | Yes — 1 per atom, delocalised along the layers | Yes | Shared within each ball |
| Melting point | Extremely high | Extremely high | Extremely high | Much lower (only weak forces between molecules) |
| Conducts electricity? | No | Yes (along the layers) | Yes — very well | Poorly |
| Hardness | Hardest natural substance | Soft and slippery — layers slide | Extremely strong for its thickness | Soft |
| Use | Cutting tools, drill tips | Pencils, lubricants, electrodes | Research into electronics, batteries | Research into medicine delivery, materials |
Your task: use secondary sources to explain each property in the table using the structure (e.g. "graphite conducts because…"). Reference your sources.
Practical¶
Suggested prac: identifying bonding type from properties. For a set of substances (e.g. sodium chloride, copper, sugar, paraffin wax, sand, graphite), test and record:
- Melting point — does it melt easily when gently heated in a test tube? (Low melting point → probably covalent molecular.)
- Conductivity as a solid — test with a simple circuit and globe or multimeter.
- Solubility in water, and conductivity of the solution if it dissolves.
Then use the properties table above to decide the bonding type for each, and justify each decision with your evidence. Heat only small amounts, in a fume cupboard if available — some substances give off irritating fumes.
Advanced
Why the chlorides don't fit the simple trend. A very small, highly charged positive ion (like \(\ce{Li+}\), \(\ce{Be^{2+}}\) or \(\ce{Mg^{2+}}\)) pulls the electron cloud of the nearby \(\ce{Cl-}\) towards itself — the bond becomes partly covalent. That weakens the ionic lattice, so the melting point is lower than the simple model predicts. \(\ce{BeCl2}\) (399 °C) is so covalent it's barely ionic at all. Going down group 2, the metal ion gets bigger and this effect gets weaker, so the bonding becomes more fully ionic and the melting point rises — the opposite direction to group 1.
Dot diagrams for trickier species.
- \(\ce{NH4+}\): \(5 + 4(1) - 1 = 8\) electrons → N with 4 single bonds to H, no lone pairs → tetrahedral. Draw it in square brackets with the + charge outside.
- \(\ce{HCN}\): \(1 + 4 + 5 = 10\) electrons → H–C≡N, one lone pair on N → linear.
- Coordinate (dative) covalent bonds: in \(\ce{NH4+}\), one N–H bond forms when \(\ce{H+}\) attaches to nitrogen's lone pair — both shared electrons come from N. Once formed, it's identical to the other three bonds.
Comparing boiling points where forces compete. A large non-polar molecule can have stronger dispersion forces than a small polar molecule has dipole–dipole forces. Example: iodine, \(\ce{I2}\) (non-polar, 106 electrons) is a solid at room temperature, while hydrogen chloride, \(\ce{HCl}\) (polar, 18 electrons) is a gas. Always consider size as well as type of force.
Electronegativity difference in numbers. \(\ce{NaCl}\): \(3.16 - 0.93 = 2.23\) → ionic. \(\ce{HCl}\): \(3.16 - 2.20 = 0.96\) → polar covalent. \(\ce{CH4}\): \(2.55 - 2.20 = 0.35\) → essentially non-polar.
Extension
Beyond the syllabus for this course — won't appear in the HSC, included for interest / depth study inspiration.
Why don't all metals conduct equally well? The "sea of electrons" model can't explain why silver conducts far better than iron, or why some elements (like silicon) are semiconductors. A better model, band theory, treats electrons in a solid as occupying energy "bands". If the top band is only partly filled, electrons move freely (a metal). If there's a large energy gap before the next empty band, nothing conducts (an insulator). If the gap is small, a little heat or light lets some electrons jump it (a semiconductor) — the basis of every computer chip and solar cell.
📎 Depth study idea
Graphene was first isolated in 2004 using sticky tape to peel layers off a pencil "lead" — the work won the 2010 Nobel Prize in Physics. A depth study could investigate how the electrical resistance of pencil lines changes with pencil grade (H → B, which changes the ratio of graphite to clay), or with the length and width of the line, linking the results back to graphite's structure.
Video/visual resources¶
- 🖥️ PhET Simulation — Molecule Shapes — add bonding pairs and lone pairs and watch the shape and bond angles change (VSEPR)
- 🖥️ PhET Simulation — Molecule Polarity — change electronegativity and shape, and see when bond dipoles cancel
- 🖥️ PhET Simulation — States of Matter — watch how the attractions between particles change as a substance is heated and cooled
- 🎥 Khan Academy — TODO: source a VSEPR theory explainer and an intermolecular forces explainer.
- 🎥 TODO: source a short video showing an ionic crystal shattering vs a metal bending — the contrast is far more memorable than a description.
Check yourself¶
-
Write the formula for (a) calcium chloride, (b) sodium sulfate, (c) iron(III) oxide.
Answer
(a) \(\ce{Ca^{2+}}\) and \(\ce{Cl-}\) → \(\ce{CaCl2}\)
(b) \(\ce{Na+}\) and \(\ce{SO4^{2-}}\) → \(\ce{Na2SO4}\)
(c) \(\ce{Fe^{3+}}\) and \(\ce{O^{2-}}\) → \(\ce{Fe2O3}\)
-
Name (a) \(\ce{KNO3}\), (b) \(\ce{PCl3}\), (c) \(\ce{CuO}\).
Answer
(a) Potassium nitrate. (b) Phosphorus trichloride. (c) Copper(II) oxide — oxide is 2−, so copper must be 2+.
-
Explain why solid sodium chloride doesn't conduct electricity, but molten sodium chloride does.
Answer
Conducting needs charged particles that can move. In the solid, the ions are held in fixed positions in the lattice. When it melts, the ions are free to move, so they can carry charge.
-
Predict the shape of (a) \(\ce{CH4}\) and (b) \(\ce{NH3}\), and say whether each molecule is polar.
Answer
(a) 4 bonding pairs, 0 lone pairs → tetrahedral. Symmetrical, so the bond polarities cancel → non-polar.
(b) 3 bonding pairs, 1 lone pair → trigonal pyramidal. Not symmetrical, so the bond polarities don't cancel → polar.
-
Explain why water has a much higher boiling point than hydrogen sulfide, \(\ce{H2S}\), even though \(\ce{H2S}\) is the bigger molecule.
Answer
Water molecules have H bonded directly to O, so they form hydrogen bonds with each other. \(\ce{H2S}\) can't (sulfur isn't N, O or F) — it only has weaker dipole–dipole and dispersion forces. More energy is needed to separate water molecules, so water boils at a much higher temperature.
-
Explain why graphite conducts electricity but diamond does not.
Answer
In graphite each carbon is bonded to only 3 others, leaving one delocalised electron per atom that can move along the layers and carry charge. In diamond each carbon uses all 4 outer electrons in bonds, so there are no free electrons.