Introduction

• Atoms form bonds to get a full outer shell of electrons.
• There are three types of bonding:
  • Ionic bonding
  • Covalent bonding
  • Metallic bonding
• To conduct electricity there much be charged that can move (ions or electrons).

Ionic Bonding

• Ionic Bond - Electrostatic attraction between opposutley charged ions.
• In ionic bonding, metals transfer electrons to non-metals producing positive metal ions and negative non-metal ions.
• Dot-cross diagrams show outer elctrons only.

Example

Covalent Bonding

• Covalent bond - is a chemical link between two atoms in which electrons are shared between them.
• Ony non-metals can get a full shell by sharing electrons.

Example

Dative Covalent Bonds

• Dative bond - Is a covalent bond between two atoms where one of the atoms provides both electrons that form the bond.
• One atom provides both of the shaired pair of electrons.

Example

Metallic Bonding

• Metallic Bond - Electrostatic attraction between positive ions and delocalised electrons.
• Metals lose their outer shell electrons to produce a lattice of positive metal ions surrounded by delocalised electrons.

Example

Giant Ionic Lattices

• The lattice is made up of oppositely charged ions.
• It has a high melting and boiling point (strong forced of attraction between ions need to be broken).
• Do not conduct when solid (iosn not free)
• Conduct when molten or dissolved in water.
• An example, Sodium Chloride.

Simple Covalent Lattices

• Consists of molecules held together by weak intermolecular forces.
• Low melting and boiling points (weak forces of attraction between molecules are easily broken).
• Do not conduct (no mobile charge carriers).
• Example - Iodine and ice.

Giant Metallic Lattices

• Lattice of metal ions surrounded by delocalised electrons.
• High melting and boiling points (stron forces of attraction between metal ions and free electrons).
• Conduct when solid (free electrons)
• Example - Magnesium and Copper

Giant Covalent Lattices

• Lattice of non-metal atoms joined by strong covalent bonds.
• Very high melting and boiling point.
• Diamond does not conduct (no free electrons).
• Graphite is the only non-metal that conducts as a solid (contains delocalised electrons).

Electronegativity and Bond Polarity

• Electronegativity - A property of an atom which increases with its tendency to attract the electrons of a bond.
• If there is a big difference in electronegativity between the atoms the electrons will be pulled towards the more electronegative atom.
• The most electronegative elements are Fluorine, Chlorine, Oxygen and Nitrogen.
• Polar molecules have permanent dipoles that don't cancel out.
• Non-polar molecules either have no dipole or temporary dipoles that cancel out.

Example

• Chlorine is more electronegative than hydrogen.

  δ+   δ-

  H-Cl

  
  

Van der Waal's

• Van der Waal's - Are the weak forces which contribute to intermolecular bonds.
• Arise from temporary dipole (uneven distribution of electrons) in one molecule that iduces dipole in another molecule.
• The more electrons, the stronger the Van der Waal's forces.
• Occur in all simple covalent substances.

Dipole - dipole

• Dipole - separation of electrical charges.
• Attraction between molecules with permanent dipoles.
• δ+ ends attracted to δ- ends.

Hydrogen Bonds

• Hydrogen Bond - is a type of attractive interaction between electronegative atom and a hydrogen atom bonded to another electronegative atom.
• Need H attached to N/O/F.
• Exposed H nulceus is strongly attracted to lone pair on N/O/F.

Diagram

Anomalous properties of water

• Water has some unusual properties due to the presence of hydrogen bonding:
  • Ice is less dense than water because ice has an open structure.
  • Water has a high melting and boiling point than expected due to the strength of hydrogen bonds that have to be broken.

Shapes of Molecules

• Pairs of electrons reoek each other and get as far apart as possible.
• Lone pairs repel more than bonding pairs - they change the angle by 2.5°.

Table