Properties of elements and compounds Notes

How Bonding and Structure Affect Melting Point, Boiling Point and Conductivity

The key idea:

The strength and type of forces between particles determine melting point, boiling point and electrical conductivity.

We’ll look at four main structure types:

1. Giant ionic lattices
2. Simple molecular covalent substances
3. Giant covalent structures
4. Metallic structures

Giant Ionic Lattices (e.g. NaCl, MgO)

1. Bonding & structure
   1. Made of positive ions (cations) and negative ions (anions).
   2. Arranged in a regular 3D lattice.
   3. Held together by strong electrostatic attractions (ionic bonds) between oppositely charged ions.
2. Melting & boiling point
   1. Very high melting and boiling points.
   2. You must supply a lot of energy to overcome the strong ionic bonds throughout the lattice.
3. Electrical conductivity
   1. Solid:
      1. Does not conduct electricity → ions are fixed in place and cannot move.
   2. Molten (liquid) or dissolved in water (aqueous):
      1. Does conduct → ions are free to move and carry charge.

Simple Molecular Covalent Substances (e.g. H₂O, CO₂, CH₄)

1. Bonding & structure
   1. Made of small molecules (like H₂O, CO₂).
   2. Inside each molecule: strong covalent bonds between atoms.
   3. Between molecules: only weak intermolecular forces (e.g. London forces, dipole–dipole, hydrogen bonding).
2. Melting & boiling point
   1. Generally low (often gases or liquids at room temperature).
   2. When they melt/boil, you’re breaking intermolecular forces, not the covalent bonds inside each molecule.
3. Electrical conductivity
   1. Do not conduct electricity (no ions, no free electrons).
   2. Exception: some special cases like acids in water (they form ions), but the pure covalent molecules themselves don’t conduct.

Giant Covalent Structures (e.g. Diamond, Silicon Dioxide, Graphite)

1. Bonding & structure
   1. Atoms joined by strong covalent bonds in a continuous network (no separate molecules).
   2. Example:
      1. Diamond: each C bonded to 4 others in a 3D tetrahedral network.
      2. Silicon dioxide (SiO₂): each Si bonded to 4 O, each O to 2 Si in a 3D network.
      3. Graphite: carbon atoms in layers; strong covalent bonds in each layer, weak forces between layers.
2. Melting & boiling point
   1. Very high – you must break a huge number of strong covalent bonds to melt them.
3. Electrical conductivity
   1. Diamond, SiO₂: do not conduct (no free electrons or ions).
   2. Graphite: does conduct (delocalized electrons move along the layers).

Metallic Structures (e.g. Cu, Fe, Al)

1. Bonding & structure
   1. Positive metal ions in a regular lattice.
   2. Surrounded by a “sea” of delocalized electrons that can move throughout the metal.
   3. Strong electrostatic attraction between ions and electrons → metallic bonding.
2. Melting & boiling point
   1. Usually high – metallic bonds are strong.
3. Electrical conductivity
   1. Solid or liquid: conducts – delocalized electrons move and carry current.
4. This is why metals are used in wires and electrical components.

Why Do Ionic Compounds Conduct When Molten or in Solution, but Not When Solid?

This is a classic exam favorite.

In The Solid State

1. Ions are locked in a rigid lattice.
2. They can vibrate but cannot move from place to place.
3. No mobile charged particles → no electrical conductivity.

When Molten (Liquid)

1. The lattice breaks down.
2. Ions are free to move around in the liquid.
3. Positive and negative ions move in opposite directions when a voltage is applied → current flows.

When Dissolved In Water (Aqueous)

1. Water molecules surround and separate the ions (solvation/hydration).
2. Again, ions are free to move → solution conducts electricity.

Using Properties to Deduce the Type of Bonding

You can often work backwards from observed properties to figure out what kind of bonding and structure a substance has.

Step-By-Step Secision Process

1. Check melting and boiling point
   1. Very high (often > 1000 °C):
      → Could be giant ionic, giant covalent, or metallic.
   2. Low (often gas/liquid at room temperature):
      → Usually simple molecular covalent.
2. Test electrical conductivity
   1. Conducts when solid:
      → Likely metallic (delocalised electrons).
   2. Does not conduct when solid, but does when molten/aqueous:
      → Ionic (mobile ions only when not solid).
   3. Does not conduct in any state:
      → Usually simple covalent or giant covalent (like diamond, SiO₂).
      → Exception: graphite conducts in one direction.
3. Check solubility
   1. Dissolves in water, conducts when dissolved → ionic.
   2. Dissolves in non-polar solvents (like hexane), low mp/bp, non-conductive → simple molecular covalent.
   3. Insoluble, very high mp/bp, non-conductive (except graphite) → giant covalent.
   4. Insoluble, conducts as solid, shiny, malleable → metallic.