Mission 1: Salt is made of charged particles. So why doesn't a salt crystal conduct electricity?

Which ions atoms form

• Metal atoms lose electrons to form positive ions: cations.
• Non-metal atoms gain electrons to form negative ions: anions.
• Group 1 and 2 metals, and aluminium, lose all their outer-shell electrons.
• Non-metals gain enough electrons to fill the outer p sublevel.
• Mg 1s² 2s² 2p⁶ 3s² loses two: Mg²⁺. N 1s² 2s² 2p³ gains three: N³⁻.
• Charge = protons − electrons. The number of protons never changes.

Transition elements: more than one ion

• Transition elements can form ions with different charges.
• They lose their 4s electrons before 3d.
• Cobalt, [Ar] 3d⁷ 4s²: Co²⁺ is [Ar] 3d⁷; Co³⁺ is [Ar] 3d⁶.
• A Roman numeral gives the charge: cobalt(II) chloride, CoCl₂.

Ionic bonds, formulae and names

• Ionic bond: electrostatic attraction between ions of opposite charge.
• It acts in all directions: each ion attracts every oppositely charged neighbour.
• Formula: combine ions so the total charge is zero. Mg²⁺ and N³⁻ → Mg₃N₂.
• Name the cation first, then the anion ending in -ide: magnesium nitride.
• Formula to name: MnO₂ has 2 × O²⁻ (−4), so Mn⁴⁺: manganese(IV) oxide.
• Number prefixes (di-, tri-) are for covalent compounds, never ionic ones.

Seven polyatomic ions to know

• Keep each ion whole. Brackets show more than one: Ca(NO₃)₂, (NH₄)₂SO₄.
• "-ate" does not fix the charge: nitrate 1−, sulfate 2−, phosphate 3−.

The ionic lattice

• Ionic compounds are giant 3D lattices of alternating cations and anions.
• There are no molecules, so the formula is empirical: the simplest ratio of ions.
• Al₂O₃ means two Al³⁺ for every three O²⁻, not a molecule of five atoms.

Properties from the lattice

• Low volatility, high melting point: every ion is held by many strong attractions.
• Conducts only when molten or aqueous: then the ions are free to move.
• The solid does not conduct: its ions are fixed, and there are no free electrons.
• Often soluble in water: polar water molecules attract the ions out of the lattice.
• Insoluble in non-polar solvents such as hexane. CaCO₃ and BaSO₄ don't dissolve in water either.

Lattice enthalpy: how strong the ionic bonding is

• Lattice enthalpy, ΔH⦵lattice: the enthalpy change to separate 1 mol of ionic solid into gaseous ions.
• MgO(s) → Mg²⁺(g) + O²⁻(g). Always positive: larger means stronger ionic bonding.
• Larger with higher ion charges, and with smaller ion radii (centres closer).
• MgO 3791, NaCl 790 kJ mol⁻¹ (data booklet). MgO's ions are doubly charged and smaller.
• Same charges? Radius decides. When charges differ, charge usually has the bigger effect.

Mission 2: Air is 78% nitrogen. Why does almost nothing react with it?

What a covalent bond is

• Covalent bond: the electrostatic attraction of both nuclei for a shared pair of electrons.
• The pair sits between the nuclei. Both pull on it, holding the atoms together.
• Octet rule: atoms tend to end up with eight valence electrons. A tendency, not a law.
• The octet is an outcome. The attraction is the bond.

Drawing a Lewis formula

• A Lewis formula shows every valence electron, bonding and lone pairs, as dots, crosses or lines.
• 1. Count valence electrons: add one per negative charge, remove one per positive charge.
• 2. Join the atoms with single bonds. 3. Complete the outer atoms' octets with lone pairs.
• 4. Spare pairs go on the central atom; if short, make multiple bonds (not B, Be).
• CO₂ has 16 valence electrons: O=C=O, two lone pairs on each O.
• An ion goes in square brackets, charge outside: [O–H]⁻, three lone pairs on O.

Fewer than eight

• Some central atoms have an incomplete octet.
• BF₃: boron's three electrons make three bonds. Boron has 6 electrons, no lone pair.
• BeCl₂: beryllium has only 4 electrons around it.
• Count both electrons of every shared pair for each atom.

Single, double, triple

• Single, double and triple bonds share one, two and three pairs of electrons.
• More shared pairs: the bond is shorter and stronger.

• More shared electrons pull both nuclei closer. (Data booklet, sections 11 and 12.)

Coordination bonds

• Coordination bond: a covalent bond whose shared pair comes wholly from one atom.
• The donor has a lone pair. The acceptor has room for a pair.
• H₂O + H⁺ → H₃O⁺: an O lone pair goes to H⁺. Arrow: O→H.
• In CO, one of the three shared pairs comes from oxygen alone.
• Once formed, it is identical to the others: all three O–H bonds in H₃O⁺ match.

Mission 3: Water and carbon dioxide both have three atoms. Why is one bent and the other straight?

Electron domains repel

• An electron domain: a lone pair or a bond. A multiple bond is one domain.
• Domains round a central atom get as far apart as possible: the VSEPR model.
• 2 domains: linear, 180°. 3: trigonal planar, 120°. 4: tetrahedral, 109.5°.
• BeCl₂ is linear, BF₃ trigonal planar, CH₄ tetrahedral.

Electron-domain vs molecular geometry

• Electron-domain geometry: where all the domains point, lone pairs included.
• Molecular geometry: where the atoms are. Lone pairs are left out of the name.

What changes the angles

• Lone pairs repel more strongly than bonding pairs, so they squeeze bond angles.
• CH₄ 109.5° → NH₃ about 107° → H₂O about 104.5°.
• A multiple bond holds more electron density, so it repels more than a single bond.
• Ethene: the H–C–H angle is about 117°, a little under 120°.
• Lewis formulas are drawn flat. They show the pairs, not the angles.

Mission 4: Water and methane molecules have almost the same mass. Why does water boil over 260 °C higher?

Polar bonds

• Electronegativity: how strongly an atom pulls a shared pair. Values are in the data booklet.
• Unequal electronegativities: the pair is shared unequally. That is a polar bond.
• The more electronegative atom is δ−, the other δ+. Or draw a dipole arrow.
• H–Cl: H 2.2, Cl 3.2. Difference 1.0, δ− on Cl.
• Bigger difference, more polar. Only the difference counts, not the size of the values.

Polar molecules

• A molecule is polar if its bond dipoles leave a net dipole moment.
• Equal dipoles arranged symmetrically cancel: CO₂, BF₃ and CCl₄ are non-polar.
• In H₂O and CH₂Cl₂ they do not cancel: both molecules are polar.
• Always check both: are the bonds polar, and does the shape cancel them?

Forces between molecules

• London (dispersion) forces act between all molecules. More electrons, more contact: stronger.
• Dipole–dipole: between two polar molecules, such as two CH₃Cl molecules.
• Dipole–induced dipole: a polar molecule polarizes a non-polar one, such as H₂O and O₂.
• Van der Waals forces is the umbrella term for these three.
• Hydrogen bond: H on N/O/F, attracted to a lone pair on N/O/F of another molecule.
• Hydrogen bonds act between molecules. The O–H bond inside water is covalent.

Which force is strongest?

• Only at comparable molar mass: London < dipole–dipole < hydrogen bonding.
• O₂ (M_r 32), London only: boils at −183 °C.
• H₂S (34), polar, adds dipole–dipole: −60 °C.
• CH₃OH (32), adds hydrogen bonds: 65 °C.
• Masses very different? London forces can win: octane (M_r 114, non-polar) boils at 126 °C.

Properties of molecular substances

• Melting and boiling overcome forces between molecules. Covalent bonds stay intact.
• Weak intermolecular forces: low melting and boiling points, so volatile.
• No conduction: no ions, and every electron is held in a bond or lone pair.
• Solubility: it dissolves if new solute–solvent attractions replace those broken.
• Polar and hydrogen-bonding solutes dissolve in water; non-polar ones in non-polar solvents.

Chromatography and R_F

• Components separate by their relative attraction to two phases.
• Stationary phase stays put: in paper chromatography, the paper and the water it holds.
• Mobile phase: the moving solvent. More attraction to the stationary phase: shorter distance, lower R_F.
• R_F = distance moved by spot ÷ distance moved by solvent, both from the baseline.
• Example: spot 2.4 cm, solvent 8.0 cm. R_F = 2.4 ÷ 8.0 = 0.30, no units.
• Same paper, solvent and temperature: a substance always gives the same R_F.

Mission 5: Diamond and graphite are both pure carbon. Why does only graphite conduct electricity?

Covalent networks: diamond

• In a covalent network, covalent bonds run through the whole solid. There are no molecules.
• Carbon and silicon both form these giant networks.
• Diamond: each C bonded to four others, tetrahedrally.
• Very hard, very high melting point: melting must break strong covalent bonds. Insoluble.
• No conduction: all four outer electrons of each C are in localized bonds.

Graphite and graphene

• Graphite: flat layers of hexagons. Each C is bonded to three others.
• Each C's fourth electron is delocalized across its layer: graphite conducts along the layers.
• Only London forces hold the layers together. They slide, so graphite is soft.
• Graphene: a single layer of graphite. It conducts, and is very strong for its mass.

Fullerenes and allotropes

• Fullerene C₆₀: a closed cage of 60 C atoms. These are separate molecules.
• Only London forces act between C₆₀ molecules, so it vaporizes far more easily than diamond.
• Allotropes: forms of one element with different bonding and structure.
• Different structures give different physical and chemical properties.

Silicon and silicon dioxide

• Silicon: the diamond structure, each Si bonded to four Si. It melts at 1414 °C.
• Si–Si bonds are longer and weaker than C–C. Silicon is a semiconductor.
• Silicon dioxide, SiO₂: each Si bonded to four O, and each O to two Si.
• A network, so SiO₂ is an empirical formula: Si : O = 1 : 2.
• Melts near 1700 °C. Hard, insoluble, non-conducting.

Molecular or network? The properties

• Decide the structure first. Then explain the property from it.

Mission 6: Tin conducts electricity. So why doesn't liquid tin(IV) chloride conduct at all?

The metallic bond

• Metal atoms release their valence electrons into a shared, delocalized sea.
• Metallic bond: electrostatic attraction between the lattice of cations and the delocalized electrons.
• It is non-directional: each cation attracts the electrons all around it.
• The metal stays neutral: the electrons balance the cations' charge.

Properties and uses

• Electrical conductivity: delocalized electrons drift when a voltage is applied. Cations stay put.
• Thermal conductivity: delocalized electrons carry kinetic energy rapidly through the metal.
• Malleability: cation layers slide; non-directional bonding holds them in their new places.
• In an ionic lattice, a shift brings like charges together. They repel, and the crystal shatters.
• Uses follow: copper wiring, aluminium cooking pans, gold beaten into thin leaf.

What makes metallic bonding stronger

• Stronger with higher cation charge and smaller cation radius.
• Both raise the electron density: more delocalized electrons, closer to each cation's centre.
• Na⁺ → Mg²⁺: double the charge, a smaller ion. Magnesium melts far higher.
• Down group 1 the cations get larger: K 63 °C, Rb 39 °C.
• Mass does not decide: the heavier atom can melt lower.

Alloys

• An alloy is a mixture: a metal plus other metals or non-metals. No fixed formula.
• Brass (copper, zinc), bronze (copper, tin), stainless steel (iron, chromium, nickel, carbon).
• Metallic bonding is non-directional, so other atoms fit in. The alloy still conducts.
• Different-sized atoms disrupt the layers, so they slide less: alloys are usually harder.

Bonding is a continuum

• Ionic, covalent and metallic are models. Real bonding usually mixes them.
• The bonding triangle (data booklet, section 17) places any binary compound.
• x-axis: average electronegativity, (χ₁ + χ₂) ÷ 2. y-axis: difference, Δχ.
• Labels: metallic, ionic, polar covalent, covalent (the last two share one area).
• Higher up means more ionic character. The right-hand scale gives % covalent.

Placing a compound, predicting its properties

• Count each element once, whatever the formula. Use data-booklet electronegativities.

• The triangle gives bond type, not structure: SiO₂ is a network, CO₂ (3.0, 0.8) molecular.

Mission 7: When a plastic softens in heat, what is breaking?

Polymers

• A polymer is a macromolecule: many small monomers joined by covalent bonds.
• Natural: cellulose, starch, proteins, DNA, natural rubber. Synthetic: PVC, nylon, polyesters.
• Plastics don't conduct: every electron is held in a localized bond.
• Many plastics (thermoplastics) soften when heated: forces between chains are overcome. Chains stay intact.
• They last for decades: strong, non-polar C–C and C–H bonds react very slowly.

Addition polymers

• One bond of each alkene's C=C breaks, and the carbons link to neighbouring monomers.
• Nothing else is made: the polymer contains every atom of the monomers.
• Chloroethene, CH₂=CHCl → poly(chloroethene), PVC: [–CH₂–CHCl–]ₙ.
• The repeating unit keeps the two former C=C carbons in the chain. Other groups hang off.
• Back to the monomer: take two main-chain carbons and restore the C=C.

The big ideas

• Ionic: charges from configurations; lattice enthalpy rises with ion charge, falls with radius.
• Covalent: both nuclei attract a shared pair; more pairs, shorter and stronger. VSEPR gives shapes.
• Polarity needs polar bonds and a non-cancelling shape. Similar mass: London < dipole–dipole < hydrogen bonding.
• Networks melt very high; C₆₀ is molecular. R_F = spot ÷ solvent distance, from the baseline.
• Metals: cations in delocalized electrons; alloys harder. Triangle: average χ and Δχ. Addition polymers: C=C opens.