Singapore Junior Chemistry Olympiad
Why does an atom get smaller across a period, even though it gains electrons?
2Why is the first ionisation energy of aluminium lower than that of magnesium?
3Why does sodium oxide make an alkali, but sulfur trioxide make an acid?
4Why is potassium more reactive than sodium, but bromine less reactive than chlorine?
5Which metals can carbon extract from their ores, and which need electrolysis?
6A galvanised bucket is scratched. Why does it still not rust?
1What does the group number of a main-group element tell you?
2Define first ionisation energy.
3Is a substance that loses electrons an oxidising or a reducing agent?
Same group, same outer configuration, similar chemical properties.
Across a period: atomic radius decreases; ionisation energy and electronegativity increase. Metals on the left, non-metals on the right.
Down a group: atomic radius increases; ionisation energy and electronegativity decrease; more metallic.
Across a period: nuclear charge rises, same shell, shielding about the same. Outer electrons are held more strongly.
Down a group: one more filled shell. Outer electrons are further away and more shielded, so held less strongly.
So across a period: radius falls, ionisation energy and electronegativity rise. Down a group: the opposite.
Atomic radius falls across Period 3.
A cation is smaller than its atom: it has lost its outer shell.
An anion is larger: the extra electrons repel each other.
Na+ to Si4+ are isoelectronic: more protons, smaller ion. P3− has one more shell than Si4+.
The energy to remove 1 mole of electrons from 1 mole of gaseous atoms, forming 1 mole of gaseous 1+ ions. Na(g) → Na+(g) + e−
Across a period it increases: more nuclear charge, same shell, same shielding.
Down a group it decreases: the outer electron is further away and more shielded.
Mg …3s2 Al …3s2 3p1
Al loses a 3p electron, higher in energy than 3s.
P …3p3 S …3p4
S loses a paired 3p electron: repulsion in the orbital.
In Chapter 3 we met it briefly: metals attract a shared pair weakly, non-metals strongly. Now, what it really is and why it varies.
The shared pair is pulled by each nucleus, through the inner electrons. The same factors as ionisation energy decide it.
Across a period it increases: more nuclear charge, same shielding, and a smaller atom, so the shared pair is closer and pulled harder.
Down a group it decreases: one more shell, so the shared pair is further away and more shielded.
Fluorine, top right, is the most electronegative element.
Your turn(a) Why is Mg2+ (0.065 nm) much smaller than a Mg atom (0.160 nm)? (b) Why is the first ionisation energy of S lower than that of P?
Quick checkWhich ion has the largest radius?
Na, Mg, Al: 1, 2, 3 delocalised electrons per atom and smaller, more charged ions. Melting point and conductivity rise.
Si: giant covalent, the highest melting point (1414 °C). A semiconductor.
P4, S8, Cl2: weak forces between molecules. S8 > P4 > Cl2: more electrons, stronger forces.
Your turn(a) Why does aluminium conduct electricity better than sodium? (b) Why is the melting point of sulfur higher than that of chlorine?
Quick checkWhich Period 3 element has the highest melting point?
Highest oxidation state = number of valence electrons: +1 in Na2O up to +6 in SO3. Electronegativity difference shrinks, so bonding goes from ionic to covalent.
| oxide | Na2O | MgO | Al2O3 | SiO2 | P4O10 | SO3 |
|---|---|---|---|---|---|---|
| structure | giant ionic | ionic, some covalent | giant covalent | simple molecular | ||
| nature | basic | amphoteric | acidic | |||
| pH in water | 13 | 9 | 7 | 7 | 2 to 3 | 1 |
Na2O + H2O → 2NaOH P4O10 + 6H2O → 4H3PO4 SO3 + H2O → H2SO4
Al2O3 + 6HCl → 2AlCl3 + 3H2O Al2O3 + 2NaOH + 3H2O → 2NaAl(OH)4
| chloride | NaCl | MgCl2 | AlCl3 | SiCl4 | PCl5 |
|---|---|---|---|---|---|
| pH | 7 | 6.5 | 3 | 2 | 2 |
Hydrolysis: small, highly charged Al3+ pulls on its water molecules until they release H+.
SiCl4 + 4H2O → SiO2·2H2O + 4HCl
PCl5 + 4H2O → H3PO4 + 5HCl
Your turn(a) Explain why Na2O gives an alkaline solution but SO3 an acidic one. (b) Element X is a solid of low conductivity. Its only chloride is a non-conducting liquid that reacts with water to give a white solid and a strongly acidic solution. Identify X.
Quick checkWhich oxide reacts with both hydrochloric acid and aqueous sodium hydroxide?
2Na + 2H2O → 2NaOH + H2
Li fast, Na darts and may explode, K very violent.
Reactivity increases down the group: the valence electron is further out and more shielded, so it is lost more easily.
Your turnWrite the equation for rubidium reacting with water, and name the products.
Only M2+: a third electron would come from an inner shell. More reactive down the group (Mg slow with cold water, Ba vigorous).
MCO3(s) → MO(s) + CO2(g)
Down the group, M2+ is larger, so its charge density falls. It distorts (polarises) the carbonate ion less, so the C–O bonds are weakened less: carbonates are more stable.
AnalogyA small, strong magnet held close bends a paper clip more than a large, weak one held further away. Mg2+ is the small, strong magnet.
Your turnWhich decomposes at a lower temperature, calcium carbonate or barium carbonate? Explain.
Quick checkWhich statement about the Group 2 elements, magnesium to barium, is correct?
| halogen | at room temperature | m.p. / °C |
|---|---|---|
| Cl2 | yellow-green gas | −101 |
| Br2 | red-brown liquid | −7 |
| I2 | purple-black solid | 114 |
Darker and higher melting down the group: more electrons, stronger forces between molecules, less volatile.
Halogens gain an electron: oxidising agents. Reactivity decreases down: the incoming electron is further out and more shielded.
Cl2 + 2KBr → 2KCl + Br2 (red-brown)
Br2 + 2KI → 2KBr + I2 (brown)
Br2 + KCl: no reaction. Halogen reduced, halide oxidised.
| bond | H–F | H–Cl | H–Br | H–I |
|---|---|---|---|---|
| kJ mol−1 | 562 | 431 | 366 | 299 |
Larger halogen, poorer overlap, weaker bond: thermal stability decreases. HI gives purple I2 fumes with a hot rod.
Your turn(a) Describe and explain what you see when bromine water is added to potassium iodide solution. (b) Why is HI less thermally stable than HCl?
Colourless, monatomic gases with low boiling points.
Full outer shells (He 2, the rest 8), so they rarely lose, gain or share electrons: unreactive.
Used for an inert atmosphere: argon in steel-making and bulbs, helium in balloons, neon in lamps.
Quick checkAstatine is below iodine in Group 17. Which is correct?
A more reactive metal loses electrons more readily.
Above H: reacts with dilute acid. Lead seems not to: PbCl2 coats it.
Above C: electrolysis. Below C: heat the oxide with carbon.
Al seems unreactive: its oxide layer protects it.
Mg + Cu2+ → Mg2+ + Cu 2Al + Fe2O3 → Al2O3 + 2Fe
More reactive metal oxidised, less reactive ion reduced.
2CuO + C → 2Cu + CO2 PbO + H2 → Pb + H2O
Na, K: stable. Ca to Cu: oxide + CO2. Ag: all the way to silver.
Worked exampleChromium does not react with ZnSO4(aq), but reacts with hot Fe2O3. Where is chromium?
Your turn(a) Excess zinc powder is added to blue copper(II) sulfate solution. Describe what you see and write the ionic equation. (b) Tin is between iron and lead. Suggest how tin is extracted from tin(II) oxide.
Quick checkWhich metal oxide is not reduced by heating with carbon?
4Fe + 3O2 + 2xH2O → 2Fe2O3·xH2O
Rusting needs oxygen and water. Salt and acidic gases speed it up.
Rust flakes off and exposes fresh iron, so it keeps going.
Paint, oil, plastic, tin or chrome. A scratch lets rust start.
Zinc loses electrons before iron. Galvanising, or Zn or Mg blocks.
AnalogyA bodyguard who takes the bullet: zinc is attacked first, even through a scratch.
Your turn(a) Why does a scratched galvanised bucket not rust, while a scratched tin can does? (b) How is rusting prevented on the parts of a steel pier under water?
1How does atomic radius change across a period?
2Name the amphoteric oxide of Period 3.
3Does chlorine react with potassium bromide solution?
4What two things does iron need to rust?