Topic 3
Topic 3 on a page
Everything from this topic in one place, for the last read before an exam. If anything here is unfamiliar, go back to the full page.
Three types of strong bonding
| Ionic | Covalent | Metallic | |
|---|---|---|---|
| Particles | positive and negative ions | atoms sharing pairs of electrons | metal ions and delocalised electrons |
| Attraction between | oppositely charged ions | two nuclei and a shared pair | metal ions and delocalised electrons |
| Examples | NaCl, MgO | Cl₂, H₂O; diamond, graphite | Na, Mg, Cu |
| Conducts? | only when molten or dissolved (ions move) | usually no; graphite and graphene yes | yes, solid and liquid (electrons move) |
All three are held by strong bonds. Ionic and metallic solids and the giant covalent structures all have high melting temperatures. A simple molecular substance does not, because only weak forces between the molecules are overcome on melting (Topic 2).
Definitions to learn word for word
- Ionic bonding: the strong net electrostatic attraction between oppositely charged ions (Giant ionic lattices).
- Covalent bonding: the strong electrostatic attraction between two nuclei and the shared pair of electrons between them (Covalent bonds).
- Dative covalent bond: a shared pair of electrons in which both electrons come from the same atom (Dative bonds).
- Metallic bonding: the strong electrostatic attraction between metal ions and the delocalised electrons (Metallic bonding).
- Electronegativity: the ability of an atom to attract the bonding pair of electrons in a covalent bond (Electronegativity).
- Polarisation: the distortion of the electron cloud of an anion by a neighbouring cation (Polarisation).
- Isoelectronic ions: ions with the same number of electrons.
- Bond length: the distance between the nuclei of two bonded atoms.
- Bond angle: the angle between two bonds from the same atom (Electron-pair repulsion).
Evidence for ions
| Evidence | What is seen | What it shows |
|---|---|---|
| High melting temperature | NaCl melts at 801 °C | many strong attractions between ions |
| Conduction | none when solid; yes when molten or dissolved | charged particles, fixed in the solid and free in the liquid |
| Brittleness | shatters when hit | a shifted layer puts like charges together, so they repel |
| Electron density map | almost zero density between the ions; nearly circular contours | separate ions |
| Migration (KMnO₄ on damp paper) | purple moves towards the positive electrode | MnO₄⁻ is negative |
| Migration (CuCrO₄ in a U-tube) | blue to negative, yellow to positive, green in the middle | Cu²⁺ is positive and CrO₄²⁻ is negative |
See Evidence for ions and Migration of ions.
Ionic strength, radius and polarisation
| Idea | Rule |
|---|---|
| Stronger ionic bonding | higher charge and smaller radius (Ionic radius and bond strength) |
| Down a group | radius increases: more shells |
| Isoelectronic ions, N³⁻ to Al³⁺ | radius decreases: same electrons, more protons |
| Which argument? | count the electrons first. Different shells: shells. Same electrons: protons. |
| Polarising power (cation) | increases with smaller radius and higher charge: Al³⁺ > Mg²⁺ > Na⁺ |
| Polarisability (anion) | increases with larger radius and higher charge: I⁻ > F⁻; O²⁻ > F⁻ |
| Consequence | more polarisation, so more covalent character in an ionic bond |
Ten-electron radii in pm: N³⁻ 171, O²⁻ 140, F⁻ 133, Na⁺ 102, Mg²⁺ 72, Al³⁺ 53.
Dot-and-cross rules
- Show only outer electrons, with dots for one atom and crosses for the other.
- Ionic: show the transferred electron in the anion, put each ion in square brackets with its charge outside, and draw no shared pairs.
- Covalent: the shared pair sits in the overlap. Hydrogen has 2 electrons and the other atoms here have 8. Draw all lone pairs.
- A double bond is 4 electrons and a triple bond is 6 electrons.
- Dative: both electrons of the pair come from the donor. Draw the arrow from donor to acceptor. NH₄⁺ and Al₂Cl₆ are the two examples, and the ion goes in brackets with its charge (Dative bonds).
Giant carbon structures
| Diamond | Graphite | Graphene | |
|---|---|---|---|
| Bonds per carbon | 4 (tetrahedral) | 3 (layers) | 3 (single layer) |
| Delocalised electrons | no | yes, along the layers | yes |
| Conducts | no | yes | yes |
| Key property | very hard | soft, slippery; high melting temperature | very strong, light, transparent |
| Use | cutting tools | electrodes, lubricant | electronics, composites |
See Giant covalent structures.
Electronegativity and polarity
Pauling values: H 2.1, C 2.5, N 3.0, O 3.5, F 4.0, Na 0.9, Mg 1.2, Al 1.5, Cl 3.0, Br 2.8, I 2.5.
- Equal values: non-polar bond. Different values: polar bond, δ+ on the less electronegative atom. A very large difference: ionic.
- A molecule is polar if it has polar bonds and the dipoles do not cancel. Symmetrical shapes (CO₂, BCl₃, CCl₄, SF₆) are non-polar. Lone pairs (H₂O, NH₃) usually make the shape unsymmetrical, so the molecule is polar (Polar molecules).
- Jet of liquid and charged rod: deflected means polar (water, ethanol), not deflected means non-polar (cyclohexane).
Shapes
| Species | Bonding pairs | Lone pairs | Shape | Bond angle |
|---|---|---|---|---|
| BeCl₂ | 2 | 0 | linear | 180° |
| CO₂ | 2 (double bonds) | 0 | linear | 180° |
| BCl₃ | 3 | 0 | trigonal planar | 120° |
| C₂H₄ | 3 regions at each C | 0 | planar | about 120° |
| CH₄, NH₄⁺ | 4 | 0 | tetrahedral | 109.5° |
| NH₃ | 3 | 1 | trigonal pyramidal | 107° |
| H₂O | 2 | 2 | bent | 104.5° |
| PCl₅ (gas) | 5 | 0 | trigonal bipyramidal | 120° and 90° |
| SF₆ | 6 | 0 | octahedral | 90° |
Counting method: outer electrons of the central atom, plus 1 for each single-bonded H or halogen, plus 1 per negative charge, minus 1 per positive charge, divided by 2. Lone pairs repel more than bonding pairs, and each lone pair takes about 2.5° off the angle (Electron-pair repulsion, Shapes with no lone pairs, Lone pairs).
Metallic properties
- Conductivity: delocalised electrons move through the lattice, in the solid and the liquid.
- High melting temperature: a lot of energy is needed to overcome the attraction between ions and delocalised electrons. Generally higher with more delocalised electrons per ion, a higher charge and a smaller ion: Na 98 °C, Mg 650 °C, Al 660 °C.
- Malleable: layers slide and the ions remain in the electron sea.
The ten traps
- Saying ionic compounds conduct with "free electrons" instead of mobile ions.
- Saying solid NaCl does not conduct because it "has no ions", instead of "ions fixed in position".
- Giving the observation for a migration practical without the reason (the ion's charge).
- Using "more shells" for isoelectronic ions, or "more protons" for a group.
- Mixing polarising power (cation) with polarisability (anion), or the direction of distortion.
- Drawing a double bond with two electrons, or leaving out lone pairs.
- Drawing a dative arrow the wrong way, or leaving off the brackets and charge on NH₄⁺.
- Defining electronegativity as the ability to gain electrons.
- Calling CO₂ non-polar because "the bonds are non-polar", instead of "the dipoles cancel because it is linear".
- Counting a double bond as two regions, forgetting the charge step, or giving one angle for PCl₅.
Mixed test
Try these with no notes, then check the answers.
- Explain why solid magnesium oxide does not conduct, but molten magnesium oxide does.
- Calculate the number of ions in 0.250 mol of Al₂O₃. Use .
- A purple crystal of potassium manganate(VII) is placed on damp filter paper with a d.c. supply across it. State what you see and explain it.
- Explain why the melting temperature of MgO (2852 °C) is much higher than that of NaCl (801 °C).
- P³⁻, S²⁻, Cl⁻, K⁺ and Ca²⁺ are isoelectronic. Put them in order of radius, largest first, and explain.
- Explain why lithium iodide has more covalent character than potassium iodide.
- Explain how a molecule of NH₃ bonds to a molecule of BF₃, and state what the arrow in the displayed formula shows.
- Explain why graphite conducts electricity but diamond does not.
- Calculate the electronegativity differences for C–H and O–H, and state which bond is more polar and which atom is δ+ in each.
- State the shape and bond angle of BF₃ and NF₃, and say whether each molecule is polar.
- Explain why the bond angle in H₂O is smaller than in NH₃.
- Explain why aluminium (660 °C) melts at a much higher temperature than sodium (98 °C).
- In the jet experiment, ethanol is deflected towards a charged rod and cyclohexane is not. State what each result shows.
Answers cover these until you have tried the questions
- In the solid, the Mg²⁺ and O²⁻ ions are held in fixed positions in the lattice, so there are no mobile charged particles. In the molten liquid the ions are free to move, so they carry charge.
- Each formula unit Al₂O₃ gives 2 Al³⁺ and 3 O²⁻, which is 5 ions. Amount of ions = 0.250 × 5 = 1.25 mol. Number of ions = 1.25 × 6.02 × 10²³ = 7.525 × 10²³, which is 7.53 × 10²³ ions.
- The purple colour spreads towards the positive electrode. The colour is due to MnO₄⁻ ions, which are negatively charged, so they are attracted to the positive electrode.
- Mg²⁺ and O²⁻ have higher charges than Na⁺ and Cl⁻, and are smaller. The electrostatic attraction between the ions is stronger, so more energy is needed to overcome it.
- Order: P³⁻ > S²⁻ > Cl⁻ > K⁺ > Ca²⁺. All have 18 electrons in the same shells. The number of protons increases from 15 to 16, 17, 19 and 20, so the same electrons are pulled closer to the nucleus and the radius decreases.
- Li⁺ and K⁺ both have a 1+ charge, but Li⁺ is much smaller, so it has a higher charge density and a greater polarising power. It distorts the electron cloud of the I⁻ ion more, so more electron density lies between the ions and the bond has more covalent character.
- N in NH₃ has a lone pair. B in BF₃ has only 6 electrons in its outer shell, so it has room for another pair. The lone pair on nitrogen is shared with boron, forming a dative covalent bond in which both electrons come from nitrogen. The arrow points from N to B, from the donor of the pair to the acceptor.
- In graphite each carbon uses 3 electrons in bonds and the fourth is delocalised, free to move along the layers. In diamond all 4 outer electrons are used in bonds, so there are no delocalised electrons.
- C–H: 2.5 − 2.1 = 0.4. O–H: 3.5 − 2.1 = 1.4. O–H is more polar. In C–H the hydrogen is δ+ (lower electronegativity than carbon), and in O–H the hydrogen is δ+ as well.
- BF₃: 3 + 3 = 6, 3 bonding pairs, no lone pairs: trigonal planar, 120°. The B–F bonds are polar but the dipoles cancel, so BF₃ is non-polar. NF₃: 5 + 3 = 8, 4 pairs (3 bonding and 1 lone): pyramidal, about 107°. The dipoles do not cancel, so NF₃ is polar.
- Both molecules have four electron pairs. H₂O has two lone pairs and NH₃ has one. Lone pairs repel more than bonding pairs, so the bonding pairs in H₂O are pushed closer together: 104.5° compared with 107°.
- Al³⁺ has a higher charge and a smaller radius than Na⁺, and aluminium has 3 delocalised electrons per ion compared with 1 for sodium. The attraction between the metal ions and the delocalised electrons is stronger, so more energy is needed to overcome it.
- The ethanol jet is deflected, which shows its molecules are polar. The cyclohexane jet is not deflected, which shows its molecules are non-polar.