1. Primary Chemical Bonding
Chemical bonding arises from electrostatic attractions between charged particles, driving atoms towards lower-energy configurations.
| Bonding Type | Definition & Nature of Force | Typical Examples | Macroscopic Properties |
|---|---|---|---|
| Ionic | Electrostatic attraction between oppositely charged ions in a giant 3D lattice. | NaCl,MgO,CaF₂ | High melting points, brittle, conduct electricity when molten or dissolved in water. |
| Covalent | Shared pair of electrons between non-metal atoms. | Cl₂,H₂O,CH₄ | Low melting points for simple molecules; very high for giant covalent macromolecular structures. |
| Coordinate (Dative) | Shared pair of electrons where both electrons are donated by the same atom. | NH₄⁺,H₃O⁺,Al₂Cl₆ | Identical in length, strength, and chemical behaviour to ordinary covalent bonds once formed. |
| Metallic | Electrostatic attraction between positive metal cations and a delocalised sea of electrons. | Na,Mg,Fe,Cu | High thermal and electrical conductivity, malleable and ductile due to layers of ions sliding. |
Ions and Formulas
You can predict the charge on a simple ion from the group of the element. Metals lose their outer electrons to form positive ions. Non-metals gain electrons to form negative ions.
| Group | 1 | 2 | 3 | 5 | 6 | 7 |
|---|---|---|---|---|---|---|
| Ion charge | +1 | +2 | +3 | -3 | -2 | -1 |
| Example | Na⁺ | Mg²⁺ | Al³⁺ | N³⁻ | O²⁻ | Cl⁻ |
Learn these five compound ions. Each one is a group of atoms with an overall charge.
| Compound ion | Formula |
|---|---|
| Sulfate | SO₄²⁻ |
| Hydroxide | OH⁻ |
| Nitrate | NO₃⁻ |
| Carbonate | CO₃²⁻ |
| Ammonium | NH₄⁺ |
The positive and negative charges must cancel. Put brackets round a compound ion when you need more than one of it.
- Aluminium sulfate: Al³⁺ and SO₄²⁻. Two Al³⁺ give +6 and three SO₄²⁻ give -6. Formula: Al₂(SO₄)₃.
- Calcium hydroxide: Ca²⁺ and OH⁻. One Ca²⁺ needs two OH⁻. Formula: Ca(OH)₂.
- Ammonium carbonate: NH₄⁺ and CO₃²⁻. Two NH₄⁺ balance one CO₃²⁻. Formula: (NH₄)₂CO₃.
- Magnesium nitrate: Mg²⁺ and NO₃⁻. One Mg²⁺ needs two NO₃⁻. Formula: Mg(NO₃)₂.
Single, Multiple and Co-ordinate Bonds
- A single covalent bond is one shared pair of electrons. It is drawn as one line, for example H-H or Cl-Cl.
- A double bond is two shared pairs, drawn as two lines (O=O, or each C=O in CO₂). A triple bond is three shared pairs (N≡N).
- In a co-ordinate (dative) bond, both electrons in the shared pair come from one atom. It is drawn as an arrow from the atom that donates the pair.
- NH₄⁺: NH₃ + H⁺ → NH₄⁺. The lone pair on N is donated to H⁺. Draw one arrow from N to H. Once formed, all four N-H bonds are identical.
- Al₂Cl₆: two AlCl₃ molecules join to form a dimer. A lone pair on a Cl atom of each AlCl₃ is donated to the Al atom of the other. Draw two arrows from Cl to Al.
Metallic Bonding
The attraction between a lattice of positive metal ions and a sea of delocalised electrons.
Why does magnesium (650 °C) melt at a higher temperature than sodium (98 °C)?
- Mg²⁺ has a greater charge than Na⁺.
- More electrons are delocalised: two per Mg atom, but only one per Na atom.
- The Mg²⁺ ion is smaller than the Na⁺ ion.
- So the attraction between the ions and the delocalised electrons is stronger in magnesium. More energy is needed to overcome it.
2. Electronegativity and Bond Polarity
The power of an atom to attract the electron density in a covalent bond towards itself (Pauling scale).
Electronegativity increases across a period (greater nuclear charge, smaller atomic radius) and decreases down a group (increased electron shielding and distance). Fluorine is the most electronegative element (4.0 on Pauling scale).
Bond Polarity vs Molecular Polarity
A bond between atoms with differing electronegativities is polar (Cδ+−Oδ−). However, whether the whole molecule is polar depends on molecular symmetry:
- Carbon dioxide (CO₂): Linear shape (O=C=O). The two equal bond dipoles act in opposite directions and cancel out completely. CO₂ is non-polar.
- Water (H₂O): Bent shape due to two lone pairs. The bond dipoles do not cancel, giving a permanent molecular dipole moment.
3. Intermolecular Forces
Intermolecular forces exist between independent covalent molecules, dictating physical properties such as boiling point and viscosity:
- London Dispersion Forces (Induced Dipole-Dipole): Present between all molecules. Constant electron movement creates an instantaneous temporary dipole, which induces an opposing dipole in an adjacent molecule. Strength increases with greater electron count and larger molecular surface area.
- Permanent Dipole-Dipole Forces: Electrostatic attractions between the permanent δ+ end of one polar molecule and the δ− end of another. Stronger than London dispersion forces for molecules of comparable electron count.
- Hydrogen Bonding: The strongest intermolecular force. Requires a hydrogen atom covalently bonded to a highly electronegative atom (F,O,N) with at least one lone pair of electrons.
- Anomalously high boiling points: H₂O, HF, and NH₃ boil at much higher temperatures than Group 16, 17, and 15 hydrides because significant energy is required to break extensive hydrogen bonds.
- Density of Ice: In ice, each water molecule forms four hydrogen bonds in an open tetrahedral lattice (hexagonal rings). Molecules are held further apart than in liquid water, making ice less dense than liquid water (ice floats).
4. VSEPR Molecular Geometry
Valence Shell Electron Pair Repulsion theory states that electron pairs around a central atom repel each other and adopt positions of maximum separation to minimise electrostatic repulsion.
Repulsion: lone pair-lone pair > lone pair-bonding pair > bonding pair-bonding pair.
| Total Pairs | Bonding Pairs | Lone Pairs | Shape Name | Bond Angle | Example |
|---|---|---|---|---|---|
| 2 | 2 | 0 | Linear | 180° | BeCl₂,CO₂ |
| 3 | 3 | 0 | Trigonal planar | 120° | BF₃,AlCl₃ |
| 3 | 2 | 1 | Bent (V-shaped) | 118° | SO₂ |
| 4 | 4 | 0 | Tetrahedral | 109.5° | CH₄,NH₄⁺ |
| 4 | 3 | 1 | Trigonal pyramidal | 107° | NH₃,H₃O⁺ |
| 4 | 2 | 2 | Bent | 104.5° | H₂O,H₂S |
| 5 | 5 | 0 | Trigonal bipyramidal | 90° & 120° | PCl₅ |
| 5 | 4 | 1 | Seesaw | 87° & 102° | SF₄ |
| 5 | 3 | 2 | T-shaped | 88° | ClF₃ |
| 6 | 6 | 0 | Octahedral | 90° | SF₆ |
| 6 | 4 | 2 | Square planar | 90° | XeF₄ |
5. Crystal Structures
There are four types of crystal structure: ionic, metallic, macromolecular (giant covalent) and molecular. The type of structure explains the melting point and the electrical conductivity.
| Type | Examples | Particles and forces | Melting point | Conducts? |
|---|---|---|---|---|
| Ionic | NaCl | Positive and negative ions; strong electrostatic attraction in all directions | High | Only when molten or dissolved |
| Metallic | Mg | Positive ions and delocalised electrons; metallic bonding | High | Yes, as a solid |
| Macromolecular | Diamond, graphite | Atoms joined by many strong covalent bonds | Very high | Diamond no; graphite yes |
| Molecular | Iodine, ice | Molecules held by weak forces: van der Waals (I₂) or hydrogen bonds (ice) | Low | No |
You may be asked to draw a structure with a stated number of particles, for example "one Na⁺ ion and the 6 Cl⁻ ions around it" or "one water molecule and the 4 molecules it hydrogen-bonds to". Label the particles and the forces.
Ionic: Sodium Chloride
- A giant lattice of Na⁺ and Cl⁻ ions. Each Na⁺ is surrounded by 6 Cl⁻, and each Cl⁻ by 6 Na⁺.
- High melting point, because there is strong electrostatic attraction between the ions in all directions.
- Conducts only when molten or dissolved, because the ions are then free to move. In the solid, the ions are fixed in place.
Metallic: Magnesium
- A lattice of Mg²⁺ ions in a sea of delocalised electrons.
- High melting point, because there is strong attraction between the positive ions and the delocalised electrons.
- Conducts as a solid, because the delocalised electrons can move through the structure.
Macromolecular: Diamond and Graphite
Diamond
Giant macromolecular lattice where each carbon atom forms four strong covalent bonds in a tetrahedral arrangement (109.5°). Extremely hard, very high melting point (>3800 K), electrical insulator.
Graphite
Each carbon forms three covalent bonds in planar hexagonal layers (120°). One delocalised electron per carbon travels freely within and along the planar layers, making graphite an electrical conductor along the sheets. Weak London forces between sheets allow layers to slide easily (solid lubricant).
- Diamond: each C is covalently bonded to 4 others, tetrahedrally. Very high melting point, because many strong covalent bonds must be broken. Does not conduct, because there are no delocalised electrons.
- Graphite: layers of hexagons, where each C bonds to 3 others. Delocalised electrons between the layers let it conduct. Weak forces between the layers let the layers slide. Very high melting point, because many strong covalent bonds must be broken.
Molecular: Iodine and Ice
- Iodine: a lattice of I₂ molecules held together by van der Waals forces. Low melting point, because only weak forces between molecules are overcome; the covalent I-I bond does not break. Iodine sublimes when warmed. It does not conduct.
- Ice: each H₂O molecule hydrogen-bonds to 4 others in an open tetrahedral lattice (hexagonal rings).
- The molecules are held further apart than in liquid water. So ice is less dense than water and floats.
Energy Changes in Changes of State
- Melting and boiling need energy. The energy overcomes the forces between particles. In giant structures, it breaks bonds.
- The temperature stays constant during a change of state. The energy separates the particles; it does not raise the temperature.
- Freezing and condensing release energy, as forces form between the particles.
- For molecular substances, the boiling point depends on the strength of the intermolecular forces. Stronger forces need more energy to overcome. For example, H₂O (hydrogen bonds) boils at 100 °C, but H₂S (no hydrogen bonds) boils at -60 °C.
Go deeper
These UK AQA pages cover the same ideas in more detail. They may include content that is not in 9620.
6. Practice Questions
Show Answer & Mark Scheme
Step 1: Count valence electron pairs around central Xenon:
- Xe is in Group 0 (8 valence electrons).
- 4 single bonds to Fluorine atoms contribute 4 electrons: 8 + 4 = 12 electrons = 6 electron pairs.
- 4 bonding pairs and 2 lone pairs.
Step 2: Apply VSEPR reasoning:
- 6 electron pairs arrange octahedrally to minimise repulsion.
- The 2 lone pairs occupy opposite axial positions (180° apart) to minimise intense lone pair-lone pair repulsion.
- The resulting molecular shape is square planar.
- Bond angle: exactly 90 degrees.
Show Answer & Mark Scheme
Mark Scheme:
- (a) Al₂(SO₄)₃: two Al³⁺ (+6) balance three SO₄²⁻ (-6). (1 mark)
- (b) (NH₄)₂CO₃: two NH₄⁺ balance one CO₃²⁻. (1 mark)
- (c) Mg(NO₃)₂: one Mg²⁺ balances two NO₃⁻. (1 mark)
Show Answer & Mark Scheme
Mark Scheme:
- Mg²⁺ has a greater charge than Na⁺ (and Mg²⁺ is smaller). (1 mark)
- Magnesium has more delocalised electrons: two per atom, compared with one for sodium. (1 mark)
- So the attraction between the positive ions and the delocalised electrons is stronger, and more energy is needed to overcome it. (1 mark)
Show Answer & Mark Scheme
Mark Scheme:
- Both are macromolecular (giant covalent) structures. (1 mark)
- Melting needs many strong covalent bonds to be broken, which needs a lot of energy. (1 mark)
- In graphite, each C bonds to 3 others, so there are delocalised electrons between the layers that can move. (1 mark)
- In diamond, each C bonds to 4 others, so there are no delocalised electrons. (1 mark)
Show Answer & Mark Scheme
Mark Scheme:
- In solid sodium chloride the ions are fixed in the lattice and cannot move. (1 mark)
- When molten, the ions are free to move and carry charge. (1 mark)
- Iodine is made of neutral I₂ molecules, with no ions or delocalised electrons to carry charge. (1 mark)