Bonding & Physical Properties (Oxford AQA International A Level Chemistry)

Ionic Crystal Structures

Table to show the particle arrangements in the Three States of Matter

Changes of state

Crystals

• Crystals are solids which are held together by forces of attraction such as:

  • Ionic bonds

  • Metallic bonds

  • Covalent bonds

  • Intermolecular forces

• The stronger the forces holding them together, the stronger the higher the enthalpy of fusion (the harder they are to melt)

• Ionic compounds are strong

  • The strong electrostatic forces in ionic compounds keep the ions held strongly together

Ionic crystal diagram

• They are brittle as ionic crystals can split apart

• Ionic compounds have high melting and boiling points

  • The strong electrostatic forces between the ions in the lattice act in all directions and keep them strongly together

  • Melting and boiling points increase with the charge density of the ions due to the greater electrostatic attraction of charges

  • Mg²⁺O^2– has a higher melting point than Na⁺Cl^–

• Ionic compounds are not volatile

  • Volatility refers to the vaporisation of a chemical

  • Large amounts of energy are required to overcome the strong electrostatic forces of attraction, which means that ionic compounds are not volatile

• Ionic compounds are generally soluble in water as they can form ion-dipole bonds

• Ionic compounds only conduct electricity when molten or in solution

  • When molten or in solution, the ions can freely move around and conduct electricity

  • As a solid, the ions are in a fixed position and unable to move around

Metallic Crystal Structures

Malleability

• Metallic compounds are malleable

• When a force is applied, the metal layers can slide

• The attractive forces between the metal ions and electrons act in all directions

• So when the layers slide, the metallic bonds are re-formed

• The lattice is not broken and has changed shape

How metals are malleable diagram

Strength

• Metallic compounds are strong and hard

  • Due to the strong attractive forces between the metal ions and delocalised electrons

Electrical conductivity

• Metals can conduct electricity when in the solid or liquid state

  • In the solid and liquid states, there are mobile electrons which can freely move around and conduct electricity

• When a potential difference is applied to a metallic lattice, the delocalised electrons repel away from the negative terminal and move towards the positive terminal

  • As the number of outer electrons increases across a period, the number of delocalised charges also increases:

    • Sodium = 1 outer electron

    • Magnesium = 2 outer electrons

    • Aluminium = 3 outer electrons

  • Therefore, the ability to conduct electricity also increases across a period

How metals conduct electricity diagram

• Since the bonding in metals is non-directional, it does not really matter how the cations are oriented relative to each other

Thermal conductivity

• Metals are good thermal conductors due to the behaviour of their cations and their delocalised electrons

  • When metals are heated, the cations in the metal lattice vibrate more vigorously as their thermal energy increases

    • These vibrating cations transfer their kinetic energy as they collide with neighbouring cations, effectively conducting heat

  • The delocalised electrons are not bound to any specific atom within the metal lattice and are free to move throughout the material

    • When the cations vibrate, they transfer kinetic energy to the electrons

    • The delocalised electrons then carry this increased kinetic energy and transfer it rapidly throughout the metal, contributing to its high thermal conductivity.

Melting and boiling point

• Metals have high melting and boiling points

  • This is due to the strong electrostatic forces of attraction between the cations and delocalised electrons in the metallic lattice

  • These require large amounts of energy to overcome 

  • As the number of mobile charges increases across a period, the melting and boiling points increase due to stronger electrostatic forces 

Macromolecular Crystal Structures

• Diamond and graphite are examples of macromolecular crystal structures

  • They are both formed from carbon and are known as polymorphs or allotropes

Diamond

• Each carbon is covalently bonded to four others with a bond angle of 109.5^o

  • The bonds point to the corners of a tetrahedron

• The result is a giant lattice with strong bonds in all directions

• Diamond is the hardest substance known

Diagram to show the three-dimensional structure of diamond

Graphite

• In graphite, each carbon atom is bonded to three others in a layered structure

• The layers are made of hexagons with a bond angle of 120^o

  • The shape is trigonal planar

• The spare electron is delocalised and occupies the space in between the layers

• All atoms in the same layer are held together by strong covalent bonds, and the different layers are held together by weak van der Waals forces

Diagram to show the layered structure of graphite

Molecular Crystal Structures

• Molecular crystals are held together by strong covalent bonds between the atoms, but exhibit intermolecular forces between the molecules

  • Intermolecular forces are weaker than covalent bonds

• Examples include iodine and ice

Iodine

• Iodine is a solid at room temperature

  • The van der Waals forces are just strong enough to maintain the solid crystal

  • Iodine is:

    • Brittle

    • Has a low melting point (114 C)

    • Does not conduct electricity as there are no available electrons or ions to carry a charge

The three-dimensional crystalline structure of iodine

Ice

• When water freezes, the water molecules remain fixed in position to form a three-dimensional crystalline structure

• This way of packing the molecules and the relatively long bond lengths of the hydrogen bonds means that the water molecules are slightly further apart than in the liquid form

The three-dimensional crystalline structure of ice